• 1.0 Introduction/Statement of Purpose
  Disclaimer
  Safety Precautions/Gear
  How to read/write a chemical reaction
  Units used throughout
  Discussion of Legality/Words of Encouragement
• 2.0 Reaction Vessels
  Glassware
  Pyrex/Borosilicate Glassware
  Soda-Lime
  Lead Glass, Vycor, Misc
  Plastics
  Ceramics
  Teflon
  Refactrory Compositions
  Metals
• 3.0 Lab Techniques (Basic)
  Refluxing
  Heating
  Distillation/Evaporation
  Filtering
  Selection of filter paper
  Vacuum Filtration
  Gravity Filtration
  Chromatography
  Electrolysis
  Titration
  Temperature control/Measuring
  Removing water from gasses or solids with drying agents
  Recyrstalization
  Measuring Weight and Volume
• 4.0 Lab Reagent Types
  Acid/Base Theory
  Acids
  Bases
  Oxidizing Agents
  Dehydrating Agents / Desiccants
  Poisonous Reagents
  Solvents
  Metals
  Halogens
  Alkali Metals
  Gasses
• 5.0 Simple Reactions

1.0 Introduction/Statement of Purpose 

There are lots of pages scattered across the internet that can serve the amateur chemist in their endeavors, however each one has its own focus.  Improvising a distillation apparatus, production of a specific chemical, some go further and tell how to stock at home labs.  Still though, it is the goal of this work to go one step further, to explain many of the basic concepts/materials/apparatuses used in chemistry.  It is further the goal of this work to depict this in a useful way directly relating to real life and real observations.  In doing so this should most closely represent anything you would run across in real life giving you the best idea as to what to expect.

        In addition to this there are two sections of experiments for the amateur and by the second half, experienced chemist to perform.  Each experiment is intended to develop skills that will be necessary to complete experiments found elsewhere in the future.  But not only does this text help to explain itself, but lays out strategies for projects that you devise yourself, sections on real research, and the gritty sections on out of control experiments and contingency plans.  Upon reading though all the material presented here you should be able to go out and perform chemistry with a degree of confidence that the at home chemist is not normally afforded.

1.1 Disclaimer

        Each person is responsible for his or her own fortunes.  It is not reasonable to hold the person writing a hunk of text responsible for your mistake should you choose to pursue their strategies.  Therefore by reading further you agree not to hold the authors of this document responsible for any injuries/fatalities that may occur from attempting to make any of the products that are outlined within.

Further, although many of these procedures have been attempted solely for the completion of this project, not every procedure was done or thoroughly researched. Chemistry inherently possesses some danger to it and there is always some chance somehow that things may go wrong. Please use common sense, if you are following a procedure to the letter and something occurs involving excessive heat generation or violent gas evolution and it is not mentioned in a procedure, do not just assume that it is normal, if it scares you take steps to rectify the situation, always have a contingency plan for any procedure you are doing for the first time.

People who are pregnant should not attempt many of the experiments described herein. Additionally chemistry should not be practiced around small children who may interfere with the proceedings, consume chemicals, or become more readily injured though any mishaps. You should not perform a procedure mentioned here until you are familiar with the procedure, the chemicals involved, and the possible extraneous reactions that could take place.

Please, for your own safety, consider all safety precautions, gas masks, gloves, aprons, gas scrubbing. Damage done with many chemicals can be forgiving, but over time it can be disastrous, and some never give you a second chance. You only get one life, take precautions now so you do not ruin it in the future.

1.2 Safety Precautions/Gear 

Before beginning on any chemical adventure there are certain steps that should be taken.  The first thing you should do before starting a new experiment is to find out the properties of any reagents that are being used, along with your intended products.  This can be accomplished in many ways, looking up in a chemistry dictionary for example.  However the most thorough way to determine the properties of a compound is to look up its Material Safety Data Sheet (MSDS).  At the bare minimum one should tie back loose hair when in the laboratory, not eat while doing an experiments and wear closed toe footwear, no sandals. Long pants are always a good thing; as are long sleeve shirts, not loose ones though.

       Regardless of the specific dangers of a chemical one should always wear gloves and goggles or a face shield.  Some experiments may even require you to wear a respirator when standing up wind is not enough.  There are very few chemicals that will go though a good pair of gloves.  Some solvents will eat them and cause them to become gummy on the outside but a good pair of gloves will last you.  One exception though is methyl mercury, which can penetrate many commonly available glove materials, but considering the toxicity of it, it is better to just not use this compound.

        When it comes to the clothing you wear during chemistry that is somewhat of a matter of debate.  The safest thing you could wear would be a full environmental suit, lacking that a disposable painting suit made of Tyvek® that covers the full body, even has a hood attached.  Another choice for a full body suit would be a Nomex® flight suit, widely available on eBay.  This is the choice for persons working at high heat or interested in pyrotechnics due to the high heat resistance of this fabric.  Common cloth jump suits like the one pictured also work to a lesser degree.  However there are other things to consider then fabric.

        A chemistry outfit does not necessarily have to be something altogether different from something you would normally wear.  The only thing that it has to be is something that you will not care if it gets ruined.  Secondly it should not be some absurdly flammable synthetic fabric.  Finally it should not be excessively loose or skin tight.  If the clothing is too loose it can knock over beakers or drag in reagents, if it is too tight and you get something on it, it might immediately soak though to the skin.  You should always be able to remove any effected clothing quickly.  In case of spill or fire.

        In the case of gloves there are several types.  Those that totally cover the digits, the palm, and a majority of the back of the hand are good for most situations.  Gloves that only cover the bottom part of the hand and cover the back with webbing are decent for handling solids but not liquids.  Disposable latex/nitrile gloves are good for most anything with nitrile being superior to latex, however there is a sweating problem.  Finally elbow length gloves are good in situations where you are handling large quantities of reagent or there is a reaction that is causing extraneous conditions, e.g. splashing though boiling so you can easily handle the reaction if need be without splatter hitting you to any extent.

        Face protection should prevent splashing into the eyes, goggles like those in the above picture are idea, they are covered from all sides and have enough air circulation to prevent them from being uncomfortable.  Face shields will work and afford some additional protection, however they do not prevent splatter from certain directions, therefore a good pair of goggles are preferred.  Just having some form of eye protection does help though, even safety glasses are better then nothing.  However if you wear contacts you should take them out before pursuing a chemical reaction, the vapors can cause them to fuse to your eye resulting in a painful removal surgery.  Please take out your contacts and work without them if you are comfortable with that, or put on regular corrective lenses.

        The other highly optional part of your protective arsenal is a gas mask.  They are available as a full mask, which covers the eyes as well, this would eliminate the need for additional protective eyewear.  And a half-mask, such as one pictured above which only covers the lower part of your face.  Masks are also available that utilize one cartridge or two, two providing easer breathing and longer cartridge life.  The protection afforded to you by your mask is directly related to the cartridge you use.  Some masks which are for military use and may be picked up surplus offer a wider variety of chemical agents to which they can offer some degree of protection, however they vary in their protection from mask to mask and therefore can only be compared on an individual basis.  Here are some gasses you may run across though and some of their filtration properties.

Gas	Filter Properties
Hydrogen Sulfide H2S	Filters designed specifically to protect against hydrogen sulfide can only do so for very limited periods of time and are designated rescue filters.  Being that they have a life of less then 10 minutes, are not designed for excessive concentrations, and are expensive, they are not economically feasible.  Hydrogen sulfide is extremely toxic and has killed many individuals working in amateur labs because it quickly deadens the sense of smell.  Common logic would dictate that you just not work with this foul smelling gas.
Organic Solvents	There are filters specifically designed to block out VOC (Volatile Organic Carbons) these work fairly well and are the most commonly available filters in my experience as they are used widely in painting.  They last for an extended period of time and are designed for constant use.  Good for working with solvents of all kinds but especially carcinogenic solvents such as halogenated hydrocarbons and benzene derivatives. (Methanol is notoriously difficult to filter out, however it does not present much of an inhalation hazard.)
Hydrogen Chloride Chlorine  Acid Gasses	There is a specific filter known as an acid gas filter.  It will block out hydrogen halides (except fluorides) and elemental halogens along with other acidic gasses such as SO2, although they cannot reliably block out nitrogen oxides.
Hydrogen Cyanide  HCN	Regular carbon filters can block out HCN, however one exposure can ruin the filter and therefore make you susceptible to anything else you attempt to protect yourself from.  As with H2S, it is better not to use this gas/liquid in your experimentation, as it is highly toxic.

        To put on a gas mask you place it over your face, straps behind your head, the exact technique is not as important as the tests you do before giving it the all clear.  First is the positive pressure test, find the exit hold for the air, usually right in the middle of the mask, cover it with your hand and breath out, if the air only 'farts' out from around the edges of the mask you are good.  Next is the negative pressure test, cover the intakes on your cylinders with your hands and inhale, no air should come in, the mask should attempt to deform inward to compensate for the uptake of air.  If you pass both of these tests your mask is correctly positioned on your face.  If either one of these tests prove negative, reposition your mask and try again.

        Special gear is necessary for working in other conditions then regular room temperature.  For working at high temperature a welding shop is a goodies store, offering welding aprons and gloves that can help up to 1000C.  On the other hand working at cryogenic temperatures might be a seasonal thing, attempting to buy thick winter gloves in the middle of summer might prove to be difficult.  Never the less do not let that deter you form the joy of experimenting under extraneous conditions.

        Aside from the personal gear that is worn it is also a good idea to have some extinguishing media around in case of fire.  A fire extinguisher will work for most situations except some metal fires.  In those cases sand (another good thing to have on hand) is used to smother the fire, however magnesium fires will not appreciate being smothered in sand and will continue to react.  In that case it is best to move away any hazardous chemicals you can get to and let the fire burn itself out.  Putting water on a magnesium fire or other highly reactive metal can easily lead to explosion.

        Finally it is good to have some neutralization chemicals on hand.  In case of an acid spill you should have a nice wide mouthed container of sodium bicarbonate (baking soda) laying around, tossing this on an acid spill will neutralize its corrosive properties and render it somewhat safer at least so you can clean it up.  Base spills are usually less of a problem but sodium bisulfate (used to adjust the pH of pools) can be conveniently located around your lab, or even boric acid just to get the situation contained.  Flushing most acids or bases with large quantities of water also helps the situation, however in the case of concentrated phosphoric or sulfuric acid very large quantities must be applied all at once to prevent flash boiling from the heat of hydration.

        Always remember that safety should come first.  It is not worth getting a severe burn because you are too cheap to but some long sleeve welding gloves or respiratory damage because you won't invest in a gas mask.  Always consider the possibility of long term damage and if you think there is something reasonably extra that you can do to prepare for an upcoming reaction spare no expense.

1.3 How to read/write a chemical reaction 

If you're just beginning to start a chemistry hobby there are a few skills that you should have.  The most useful of these is reading and writing chemical reactions.  There are many places on the web and in books and classrooms what will be able to more thoroughly explain this procedure then I will be able to, and it is not something easily explained, therefore this section is mostly to just show the standard way in which a chemical reaction will be written in this text. 

        Pictured above is your basic periodic table (See section 12.1 for a complete legible listing off all the elements in alphabetical order). The periodic table lists elements by increasing atomic number (which usually means increasing weight).  Also it has trends in it which can allow you to predict the properties of an element in its uncombined state.  The most common grouping that people fall back on is family.  That is the vertical groups of elements.  For example, on the far right, second to last column, beginning with the letter F, that family is the halogens.  In descending order they are Fluorine, chlorine, bromine, iodine and raidoactive astatine.  Fluorine is the most reactive of the group and the reactivity decreases as you go down.  

Reactions in this work will be in the form of  A + B ---> AB

        A specific example would be 2H2 + O2 ----> 2H2O Translating this from chemical speak to common tongue is simple.  Looking at a periodic table you find that the letter H represents hydrogen, a colorless odorless gas, and that O represents oxygen, and I would hope you are all familiar with the product H2O.  The prefix, that is the number before the letter signifies the number of whatever it is in front of that take part in the reaction.  The numbers behind the letters stand for the number of that element in a compound.  Hydrogen and oxygen are gasses and they exist not as individual atoms but as two atoms bound together, hence the two behind each of them.  So what it is saying is, four molecules of hydrogen, combine with two molecules of oxygen, to produce two molecules of water.  Chemical reactions of this type are balanced, with each molecule appearing in the same quantity on both sides, if that is not the case the equation is dubbed 'unbalanced' and steps must be taken to rectify it.

        Exceptions to this rule are nonstoichiometric reactions, reactions that do not have a specific reaction that takes place and a number of products can be formed under different conditions.  This includes a wide variety or organic reactions, a specific example being:  C6H6  --(HNO3/H2SO4)--> C6H5NO2  Notice how the by product, water is not included in the reaction, and that the amount of HNO3 (nitric acid) and H2SO4 (sulfuric acid) reacting with the C6H6 (benzene) to produce C6H5NO2 (nitrobenzene) is not a part of the equation, this can be read 'in the presence of' therefore allowing the reaction to be read, "Benzene, in the presence of nitric acid and sulfuric acid reacts to produce nitrobenzene".  Another time that information can appear in between the arrows could be a catalyst which only speeds up the reaction, additionally temperature information can appear there as can pressure.  Specific reaction conditions are not usually included in the condensed equation and it is not safe to assume that every reaction you see will run at STP.

        Now lets say you want to write a chemical reaction.  First you must know the chemicals/elements involved along with what you believe to be the products.  You write them out in the format previously described, then you attempt to balance the equation by adding to the products or reactants side.  If an element only shows up on one side something is automatically off and must be corrected.  Just remember that there is not always one correct chemical reaction.

In addition to the common elements there are also components that the average chemist will come across that almost behave as though they are elements. They are more evident in aqueous solutions in the form of ionic compounds. Collectively they are known as ions but more specifically positively charged species are known as cations and negatively charged species are known as anions. Most aqueous chemistry revolves extensively around cations and anions and it is quite useful to have a ready reference list of cations, anions, and their respective charges, and it just so happens that there are many lists avalible besides the one here:

Anions:

Acetate	C2H3O2^-	Hydrogen Carbonate	HCO3^-	Hydride	H^-
Arsenate	AsO4^3-	Carbonate	CO3^2-	Hydroxide	OH^-
Arsenite	AsO3^3-	Chloride	Cl^-	Nitrate	NO3^-
Azide	N3^-	Hypochlorite	ClO^-	Nitrite	NO2^-
Bismuthate	BiO3^-	Chlorite	ClO2^-	Nitride	N^3-
Bisulfate	HSO4^-	Chlorate	ClO3^-	Oxide	O^2-
Sulfate	SO4^2-	Perchlorate	ClO4^-	Peroxide	O2^2-
Hydrogen Sulfite	HSO3^-	Chromite	CrO3^2-	Phosphate	PO4^3-
Sulfite	SO3^2-	Chromate	CrO4^2-	Phosphite	PO3^3-
Thiosulfate	S2O3^2-	Dichromate	Cr2O7^2-	Metaphosphate	PO3^-
Hydrosulfite	S2O4^2-	Cyanide	CN^-	Phosphide	P^3-
Peroxy-disulfate	S2O8^2-	Thiocyanate	SCN^-	Permanganate	MnO4^-
Bisulfide	HS^-	Cyanate	OCN^-	Iodide	I^-
Sulfide	S^2-	Fluoride	F^-
Borate	BO3^3-	Formate	HCOO^-
Bromide	Br^-	Oxalate	C2O4^2-

Red = Usually insoluble in water Blue = Normally soluble in water Black = Follows no trend

Cations:

Hydronium Ion	H3O^+	Ammonium	NH4^+	Lithium	Li^+
Sodium	Na^+	Potassium	K^+	Magnesium	Mg^2+
Calcium	Ca^2+	Barium	Ba^2+	Chromous	Cr^2+
Chromic	Cr^3+	Manganous	Mn^2+	Manganic	Mn^3+
Ferrous	Fe^2+	Ferric	Fe^3+	Cobaltous	Co^2+
Cobaltic	Co^3+	Nickelous	Ni^2+	Nickelic	Ni^3+
Cuprous	Cu^+	Cupric	Cu^2+	Zinc	Zn^2+
Silver	Ag^+	Aluminum	Al^3+	Stannous	Sn^2+
Stannic	Sn^4+	Plumbous	Pb^2+	Plumbic	Pb^4+

Mind you these are just a fraction of the available cations and anions available. The charge of an unknown cation is usually more easily determined then that of an anion, especially if you are given a name. Charges of anions usually stay constant whereas metals can have differing charges, knowing the anion a metal is coupled with can give you an indication of what the oxidation state of the metal is. In addition some names are currently written out using the stock system. This greatly simplifies things, instead of a name like manganese dioxide you get manganese (IV) oxide, the Roman numeral four indicating that manganese is in the +4 state and therefore knowing that oxygen has a negative two charge you can determine the formula of this compound to be MnO2. The use of –ous and –ic at the end of some names to differentiate between the higher and lower oxidation states is an older phenomenon and is somewhat being phased out, however tin (stannous +2 and stannic +4) and lead (plumbous +2 and plumbic +4) are somewhat stuck in this system of naming. Regardless, there are many anions, many cations, existing in different situations, some not stable in water, some only found in water, and others only existing in the solid state. Just remember the overall charge of a molecule must remain neutral.

1.4 Units used throughout

The system used in this text will be the most accepted system in chemistry academia.  The all mighty metric system.  Units of weight will often be expressed in grams (g), of volume, in liters (l) and milliliters (ml) and time in seconds (s), hours (h), and days.  In addition temperatures will be measured in Celsius.

        When it comes to liquids though there are different units that come into play aside from milliliters.  The most useful unit is molarity.  This is defined as the number of mols of a substance (solute) dissolved in 1 liter of substance (solvent).  From here you can convert one molarity solution to another using the formula:

Molarity Initial (Mi) x Volume Initial (Vi) = Molarity Final (Mf) x Volume Final (Vf)

For Example:

Chemoleo wants to make a 1 M NaOH Solution in water.  So he weighs out one mole of NaOH, looking at the periodic table he finds the atomic mass of sodium to be 22.9, that of oxygen to be 15.9 and that of hydrogen to be 1.0, adding these together he gets the weight of one mole of NaOH to be roughly 40 g.  So after weighing out 40 g of sodium hydroxide prills he adds to them enough water to make the total volume 1 L thus making a 1 M solution.  This sits on his shelf for quite some time until one day he finds that he needs 100ml of a .5 M NaOH solution.  Having three components of the above equation he can solve for the initial volume of 1 M NaOH he needs to end up with a 100 ml amount of a .5 M solution.

1M x (Vi) = .5 x 100 ml

Vi = 50 ml

So Chemoleo must take 50 ml of his 1 M NaOH solution and add to it 50 ml H2O to bring the total volume to 100 ml of .5 M NaOH solution.  Remember to label any reagents you keep laying around Chemoleo.

The molarity unit is exceptionally good for one specific reason, it greatly simplifies calculation involving precise reactions and the amount of reagents your are dispensing.  Molarity is heavily used in stoichiometry and is the staple method of labeling many lab reagents.

Another method of measurement one will come across is the percent (%) solution.  There are different variations on this, the weight/volume method, the volume/volume method, and the weight/weight method. One common example of a % solution would be 6% NaOCl available OTC as bleach.  

BromicAcid just bought 3.8 L of a 6% NaOCl solution.  In order to keep his lab space organized he must retain the labeling method he has already begun for his other reagents, therefore he must determine the molarity of the 6% NaOCl solution.  So the solution is 6% NaOCl by weight, so 6.0 g / 100 g solution.  Now, density would come in handy here, however Bromic was unable to find the information on the web and is too lazy to do physical measurements, therefore he is assuming that the density of the solution is close to water so 1.0 g/ml  therefore 1000 ml or 1000 g would have approximately 60 g NaOCl.  NaOCl has an atomic weight of around 74 g/mol therefore Bromic has 60/74 = .80M solution of NaOCl to store on his shelf.

Often times, as seen above the density of the solution is necessary to determine a more precise molarity calculation from the percent solution.  Tables are available online and in the CRC and elsewhere that give molarity to percent to density conversions that will aid in this task.

Common Percent Solutions to Molarity

Substance Name	Percent Solution (in H2O)	Molarity
Sulfuric Acid H2SO4	100% 91% 40%	18.7 M 17.1 M 5.4 M
Nitric Acid HNO3	70% 90%	15.8 M 21 M
Hydrochloric Acid HCl	20% 28% 38%	6.0 M 8.7 M 12.4 M
Ammonia  NH4OH	4%	2.3 M
Acetic Acid  CH3COOH	5%	.8M
Sodium Hypochlorite NaOCl	6% 10% 12.5%	.85M 1.45 M 2 M

1.5 Discussion of Legality/Words of Encouragement 

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2.0 Reaction vessels

It all starts, even before the chemicals, with what will you be doing your reactions in? In the beginning it is common to improvise your glassware, just re-using old jars and bottles to store reagents that you procure or produce or run reactions.  However as time goes on you start to realize you might not be able to heat your bottles without them shattering, and those pop bottles that at one time seemed like a stroke of genius to store things in, are not melting like candles from the corrosive fumes.  Well, well all have to start somewhere, and even soda glass has its place, so take the time to read through these varied reaction vessels.

2.1 Glassware 

        Most laboratory work is safe to conduct some sort of glass apparatus. And that’s great news, glass is resistant to most chemical attack; notable exceptions being strong hot bases, and most definitely hydrofluoric acid/some fluorides. Glass also has a high melting point, technically glass is surprisingly defined as a liquid, but for all intents and purposes it acts as a solid. Glass will deform at high temperature but some types of glass will shatter along the way as you will see from the following descriptions. Another plus is that it is amorphous, and by lacking a crystal structure it is clear, allowing you to see reactions taking place inside the vessel and to allow measuring of liquids using graduation marks found on many pieces of glassware. Glassware is also convenient for storing regents for long periods of time; carrying out complex refluxing; distillation under high heat or pressure/vacuum, glass is the containment choice for nearly every chemist under most situations.

        There are many forms of common glassware including beakers, flasks, tubes, test tubes, funnels, pipits, graduated cylinders and watch glasses. There are also more exotic (and much more expensive) glassware products including separatory funnels, ground glass jointed distillation flasks and jacketed condensers. I am only going to explain the purpose of some of the more common glassware that a new home chemist would have.

Standard funnel	Beaker	Round Bottom (RB) Flask
Graduated Cylinder	Volumetric Flack	Sepretory Funnel
Erlenmeyer Flask	Claisen Adapter	90 Deg Vacuum Adapter
Still Head with Thermometer	Vigreux Column	2-Neck Flat Bottom (FB) Flask
	Liebig Condenser

Beakers: These are simple cylinders with a pour spout on the lib and a flat bottom. Many times beakers have graduations on the side but be warned that these are not as accurate as from a graduated cylinders. Beakers are used to mix or dissolve substances, as simple heating vessels and sometimes as heating or cooling bath containers.

Flasks: There are two types of flasks; Florence flasks (sometimes called boiling flasks) and Erlenmeyer flasks. Florence flasks have a round body with one or more necks going into them. Some have round bottoms and some have flat bottoms. Round bottomed flasks need stands to hold them up but are stronger so you can use a vacuum with them without fear of an explosion. Erlenmeyer flasks have a cone-like body and are used for simple heating and, with a side nipple, for vacuum filtration. Volumetric flasks are simply Florence flasks with a flat bottom and a very long neck with a mark at the 500ml or 1 liter line. They are used to prepare a solution of known molarity.

Tubes: Tubes are simply glass cylinders. Some are made of Pyrex but most are made of soda glass. By melting and blowing with the help of a burner the home chemist can make simple equipment to help with an experiment. For example he could wrap a cooking thermometer made of metal in glass to increase its chemical resistance or make a simple gas drying tube. Pyrex tubing must be melted with an oxygen rich flame. Tubes can be bent once heated to carry liquids or gas to different glassware.

Test Tubes: Test tubes are simply tubes with a rounded end and a lip made of Pyrex or soda glass. Small reactions can be run in them and small amounts of substances can be stored in them. For example a small bit of potassium metal could be stored in oil in a test tube.

Funnels: Funnels can be used to filter things when you use filter paper or to add liquids or powders into a small area like the mouth of a test tube.

Pipits: Pipits are glass tubes with a small hole in one end and a larger hole in the other. They have very accurate volume measurements on the side of the glass. Pipits are used to suck up liquids, measure small amounts of liquids and place small amounts of liquids somewhere. Never pipit by mouth; always use a rubber pipit bulb.

Graduated Cylinders: These are simply large tubes with a stand on the bottom and a pouring spout. They are used to measure volume and come in sizes from 10ml to 500ml. 100ml cylinders are the most common.

Watch Glasses: These are curved dome-like pieces of glass that can be used to hold powders, cover beakers or make "cold fingers" for sublimation purification of things like iodine crystals.

        These are only a few of the equipment made of glass that can be found in a well stocked chemist's lab. 

2.1a Pyrex/Borosilicate Glassware

        "Pyrex" is a brand of high quality borosilicate glass but is used to mean all sorts of heat resistant glass. Borosilicate glass is a heat resistant glass that most glassware is made of. Beware of cheap glass made by Bomex. This glass is more likely to crack under thermal or mechanical stress. Pyrex and Kimax are good mid priced glassware while Duran is top quality.

Never assume that your glassware is Pyrex or other another heat resistant type, back in my early days I was planning to boil down about 400ml of CuCO3 in water. At that time I did not have a nice hotplate like I do now so I tried to use the stove. I grabbed a cooking bowl made of glass and poured the greenish mess into it. I then placed it on the stove and started stirring it when after a few minuets it cracked into about four parts. The nasty stuff got all over the oven and dripped onto the floor and the range area. A thousand thoughts started to rush through my head. "CuCO3 + HCl in stomach --> Death?" I cleaned like a cheap animation on a laserdisk stuck in fast forward. Don't assume all glass is Pyrex without looking.

Before heating any borosilicate glassware it is good to check it carefully for defects.  Cracks and pits known as stars can lead to catastrophic failure at high temperature due to the expansion of the glass causing tension along the fractures created.  Although it is not something to worry about compulsively, if you are using your glassware over high heat containing any corrosives, oxidizing materials, or anything that presents a high hazard situation then beforehand do yourself a favor and give your glassware a quick check over for defects. 

2.1b Soda-Lime

2.1c Lead Glass, Vycor, Misc.

2.2  Plastics

2.3 Ceramic 

2.4 Teflon

2.5 Refractory Compositions

2.6 Metals

As with all these other reaction vessels, metals have their own notch were they work the best. The actual value of a metal vessel is of course directly related to what metal it is made out of:

Metal	Working Temperature	Chemical Resistance	Additional Properties	Obtained From
Nickel (Ni)	900 C	Very highly resistant to alkali conditions, resistant to non-oxidizing acids	Can be used to handle fluorine or other halogens.	Nickel can be bought in the form of crucibles from chemistry suppliers
Iron (Fe)	1200 C	Iron will dissolve in acids readily, however is it somewhat more resistant to alkalis, it oxidizes easily.	Iron oxide that forms on the surface of objects adheres loosely flaking off and leading to further oxidation.	Iron end caps for plumbing are cheap and readily available. The shiny end caps are galvanized and have a thin layer of zinc plated on them.
Stainless Steel	1000 C	More resistant to acids and bases then iron alone. Less easily oxidized in general.	Can cause hard to determine contamination to reactions due to varying compositions.	Mixing bowls, measuring cups, and other kitchen containers can often be found to be made of stainless steel.
Copper (Cu)	775 C	Somewhat resistant to acids, equally resistant to bases, better then iron, on par if not slightly better then stainless.	Forms soluble highly colored contaminates. Clean before every use due to oxidation by air. Can be used with fluorine or other halogens.	Copper end caps are available for plumbing; they are perfect for amateur experimenting.
Tin (Sn)	250 C	Weak against acids and bases.	Tin forms an oxide coating when exposed to concentrated oxidizing agents that can prevent it from reacting further.	Unknown, tin cans actually only have a insignificant tin coating, therefore they do not convey the properties of tin entirely.
Aluminum (Al)	550 C	Very weak against acids and bases.	Forms a tenacious oxide coating that prevents further oxidation in strong oxidizing conditions such as HNO3 >75%	Aluminum end caps and pipes are available in larger home improvement stores.
Silver (Ag)	700 C	Strong against acids and bases.	-NA-	Expensive, hard to find vessels made of silver, used for work with hydrazine.
Platinum (Pt)	1200 C	Very resistant to most anything	-NA-	Very, very, expensive, platinum vessels for chemistry are hard to come by as well due to these price constraints.

So what’s the consensus? If you cannot perform a reaction in glass for one reason or another then you need another choice. Examples of extreme circumstances being, high temperature, reaction mixture attacks glass, or thermal shock might be a problem, metals can be a cheap alternative.

3.0 Lab Techniques (Basic) 

3.1 Refluxing 

3.2 Distillation

        Distillation is one of the most essential procedures in all of at home chemistry in my honest opinion.  When it all comes down to it distillation is about removing a volatile product from a solution, usually for purification and by the same logic, extraction.  The most basic of distillation apparatuses is to the right.

        It consists of four essential parts.  The tube on the left is the distillation flask (this is also known as the reaction vessel), or in this case the test tube.  This leads to a condenser, in this case a glass tube.  The condenser serves the essential role of taking the gasses produced from heating the distillation flask and cooling them down, causing you to end up with a liquid.  In this case the condenser is the most basic design, air cooled, not very efficient, but made more efficient by a fan blowing on it.  Other better condensers utilize running water to cool them or mixtures that can obtain even cooler temperatures.  

        The condenser tube leads into a receiving flask.  This can be cooled as shown in the picture by ice, this is especially necessary if your product that you are distilling over is significantly volatile.  The final part of the setup, the one that really makes it work is your heat source.  A Corning hot plate is about top of the line for an at home lab but basically anything that gets hot and hopefully has a heating control will work.  Be sure to use borosilicate glass or another heat resistant material though for your reaction vessel, otherwise the heating will cause the glass to crack and break.  For a more precise heat control the use of a 'bath' is advised, this is just a beaker full of a fluid (oil in the picture above) that is heated directly, thereby heating your distillation flask more evenly and consistently and if things get out of control, it can act as a heat sink.  The bath temperature is usually about 10C higher then the reaction flask temperature as a rule of thumb.

Now, heating a liquid mixture to boiling and condensing the vapor may sound easy, but there are several small but important factors that need to be kept in mind.

First of all, the intent of the distillation is to end up with the most volatile compound of your mixture in the receiving flask and the less volatile compound(s) in the distillation flask. No problem you say, because I am condensing the most volatile component. This is, unfortunately, not entirely true. The other components in your distilling flask also have a vapor pressure, which rises with the temperature. As a result the vapor you are condensing mainly consists of the most volatile compound, but it will also contain a fraction of the less volatile compound(s). This fraction will be large if the boiling points of the compounds don't differ much (less than 50C) and small when the boiling points are quite different.

For example, distilling wine (usually 12,5 vol% of ethanol in water) will yield a solution which is greatly enriched in ethanol content, but it will still contain a considerable (about 50-40%) amount of water.

Secondly, as your distillation proceeds, the concentration of the most volatile compound decreases and the boiling point of your mixture rises (sic). As a result, your vapor will contain more of the lower boiling compound(s) as the distillation proceeds towards completion.

There are several ways to solve this problem, the most important being fractional and azeotropical distillation.

Fractional distillations is actually quite simple, the idea is to redistill your distillate until it's pure. Now, if we take our previous example of wine, about three distillations would be needed to attain reasonably pure ethanol. It goes without saying that this is both time and energy consuming. Being the inventive chaps they are, some chemists came up with a clever solution to this problem: The fractionating column. This simple but nifty device allows you to separate close boiling compounds in one run or in considerable less runs than a simple distillation would require.

The fractionating column is placed between the heated flask and your still head with the thermometer. Vapors which pass through it cool down as they rise and eventually condense, the compound(s) with the highest boiling points condense first against the walls. As a result there are countless condensation cycles taking place in your column. The condensation of the highest boiling compound(s) delivers condensation heat which evaporates the most volatile compound which is running down the wall. As a result the vapor coming out of the top will almost exclusively contain the most volatile component, unless you are distilling an azeotrope. One could say that inside the column several successive distillations are taking place simultaneously.

There are several types of fractionating columns and they all have their specific uses. The most commonly used column is the Vigreux column (picture?). It has a relatively small surface area but a high flow rate. The standard 30cm Vigreux column is ideal for separating compounds that have a difference in boiling point of 20C or more and it can be used for vacuum distillation. Vigreux columns can be made longer or stacked to improve separation, but above a length of one meter once should consider filled columns. Filled columns come in all different sorts and sizes, but they al work on the principle of maximal surface area and therefore they are usually filled with irregularly shaped objects. They retain a lot of liquid and they are also not very well suited for vacuum use. If you are planning to distill compounds with high boiling points or you are using large or long columns you should consider insulating the column to minimize heat loss.

A very important point when performing a fractional distillation is monitoring the temperature, certainly when separating more than two compounds. You’ll notice that the temperature suddenly skyrockets when the most volatile compound is depleted from your mixture because of the good separation of your column. Good measuring of temperature can only be achieved by correct positioning of the thermometer, which should be just below the bend towards the cooler, so that it’s being immersed in the vapor. (PICTURE!!!!!) Slight deviations of a few mm can cause temperature-reading failures of several degrees centigrade.

Azeotropical distillation could be explained as cheating. The trick here is that a third substance is added to your mixture. This then forms an azeotrope, preferably with the compound you do not need to isolate. An azeotrope is a mixture of two or more substances, which can’t be concentrated anymore by distillation because both components have the same vapor pressure at the azeotropic point.

In this case, the azeotrope needs to have a boiling point that differs substantially from the other component in your mixture. Then you can remove the otherwise hard to separate component azeotropically with the third substance. This process is often used during esterification. Toluene is added to the mixture, which forms an azeotrope with water and boils off. After cooling down the water and toluene separate back into two layers, but this is not always the case.

Azeotropes can also complicate a good separation attempt. A good example is nitric acid. Nitric acid forms an azeotrope with water, which contains 69,2% nitric acid. Simple distilling won’t work and neither does fractional distillation. So what now? The most commonly used method is to break the azeotrope by adding another substance. This is completely the opposite of azeotropical distillation where a substance is added to form an azeotrope, so beware of confusion.

In the case of nitric acid, the most applied method is to add sulfuric acid, as this has such a great affinity towards water it will easily “steal” the water from nitric acid. As a result the nitric acid behaves as if it were all by itself in your flask and thus obtains its’ normal boiling point, which is far lower than the hydrated sulfuric acid. One can also snoop the water off by adding a very hygroscopic salt like magnesium nitrate, the only condition being that it does not react with the acid.

The regulation distillation setups used in chemistry labs use ground glassware.  The joints are all tapered glass and fit together snugly, with or without the use of a sealant, which can be anything from silicone gel to concentrated sulfuric acid.  There are two extreme sizes for the professional distillation setup, the 14/20 sets which would be considered the smaller scale, and the 24/40 sets, considerably larger scale.  For example, the largest flask I have seen for a 14/20 set is 100 ml, the largest for a 24/40 is 10 L.  The smaller setup has the advantage of taking up less space and using less of the distillate to wet the vessel therefore resulting in a greater yield, it is great for distilling small amounts but the purification of 3.8 L of over the counter paint thinner may be a pain.  In contrast a 24/40 setup is perfect for this larger project.  But is not good for small amounts and takes up a larger space.  It is really up to the chemist and what scale they will be working on to decide.  However the 24/40 joints are more readily available online and elsewhere then 14/20.

Here is a quick checklist to follow before performing a distillation:

1. Check if your compounds form an azeotrope. If not, or if your starting material is below or above the azeotropic concentration, proceed with a normal distillation first. If your compound does form an azeotrope, proceed to step five.
2. Check the difference in boiling points.
3. Check the vapor pressure of the higher boiling component at the boiling point of the lower boiling compound. If this is significant and purity is required, use a column.
4. If you can’t achieve sufficient separation with a column or you don’t have a sufficient column at your disposal, check weather you can use an azeotrope to remove ONE of the components.
5. If your compounds form an azeotrope which you want to separate, check if there’s something that’ll form a stronger azeotrope with it or which has a greater affinity for the compound.
6. Last but not least, check if your components won’t decompose at the necessary temperatures. If yes, carefully read section 8.6 under the advanced techniques on how to work with and distill under vacuum .

Distilling HNO3 (Nitric Acid) from a mixture of a nitrate salt and sulfuric acid is a time tested way to isolate this useful oxidizing acid.  So an adventurous chemist combined an unweighed quantity of NaNO3 (sodium nitrate) into a flask with a large quantity of 94% H2SO4 (sulfuric acid) and attached a condenser though which water was run and in turn this ran to a receiving flask.  Some boiling stones were also added (pieces of obsidian) to help ease the boiling process and make it smoother.  However complications where run into shortly after heat was applied, the black boiling stones started to color the mixture black, and the gas running into the condenser was a dark red, not what would be expected for a clear to off yellow acid.  Regardless distillation was continued and in the end the chemist ended up with 20 ml of a dark red volatile distillate.  Upon addition to water it decolorized and left an acid solution.  Although not obviously apparent to the chemist at the time they had distilled NO2, a highly toxic gas that is one of the decomposition products of nitric acid.  Their ambitiousness and inexperience resulted in them heating the reaction mixture too high and the concentration of their acid only would have allowed 95%+ HNO3 to make it over, which would have called for a vacuum because high concentration HNO3 decomposes somewhat more readily then the more common 70% grade.  The chemist eventually realized their mistake by simply observing the physical properties of the product ran in contrast to those of the desired product they thought they had.  Luckily the chemist made it out unscathed.

Important Safety Concepts:

Heating:

When distilling one should use a safe method of heating, which prevents your glass from cracking and improves its life. It also protects you from a bursting apparatus which showers you in dangerous chemicals.

Flames:

Open flames are NOT a safe way of heating. Flames cause localized overheating and are especially a hazard when distilling flammable substances.

Electric Heating:

Electric heating mantles or hotplates can be used with certain exceptions. Under normal pressure one can safely use a hotplate (preferably with magnetic stirring) to heat an erlenmeyer. However, one should be careful with flammable substances. If the temperature of the hotplate surface is higher than the auto ignition point of your substance, it can not be used unless the whole apparatus is sealed to the entry of atmospheric oxygen and the output gasses are property taken care of. Same goes for electric heating mantles. Heating mantles should be used for flasks because a hotplate causes localized overheating with its flat surface! Note that heating mantles usually impair magnetic stirring, unless you buy a magnetic stirrer which fits the heating mantle, but these are usually within the >500$ range.

Baths:

Baths are ideal for heating flammable substances, certainly water baths. Water has the advantage of being not flammable and it’s high heat capacity can be beneficial when distilling. However, water evaporates rapidly when used above 80C. Oil baths don’t evaporate as quick as water, but they have other disadvantages. Most commonly available oils smell and are flammable. They also pose a severe explosion hazard when distilling oxidizing material. Ideal are silicone baths or other non flammable synthetic oils, but these are usually hard to get and/or expensive. For lower temperature applications or if you just don’t care about your hot plate a sand bath can be used by simply filling a pan with sand and putting on your plate, but at higher temperatures the insulating effect of the sand can burn out the heating element in a hot plate so be warned.

Sealing Joints:
Sealing your joints is very important. Improperly sealed joints will cause losses but they also pose a safety hazard, because they allow air to enter. This is mainly dangerous when distilling flammable substances or when distilling under vacuum. Joints are commonly sealed with grease. There are several types of grease and they all have their disadvantages. Vaseline is cheap to get and easy to use, but it’s chemical resistance is limited and it will contaminate organic solvents. Other options include silicon oil, but this is expensive and does not provide 100% chemical resistance either, and specialty high vacuum greases composed of higher fluorinated hydrocarbons but they can cost a bundle. Therefore the author recommends cheap, white, teflon tape available from any hardware store to seal pipe joints. This easily fits between the joints and has excellent chemical resistance.

• Evaporating to Dryness

        This is a common procedure for the inorganic chemist, less so for the organic chemist but none the less important.  Evaporating to dryness is a feasible method of recovery of a pure product providing:

1. Your intended product will not decompose at the temperatures necessary to volatize the solvent. [Note: Vacuum can decrease the necessary temperature to remove the solvent and prevent decomposition of your product]

2. Any other compounds in your solution besides your intended product are also volatile.

3. Your intended product will not volatize to any major degree at the temperatures used.

4. Your intended product will not explode at the temperatures used/is not pyrophoric at these temperatures.

5. An extremely pure product is not required/additional purification will take place.

        Examples of simple procedures would be:

· Making AgNO3 by dissolving silver in HNO3 and boiling off the HNO3/Water.

· Neutralizing BaCO3 with HCl then boiling off the water and HCl.

Examples of procedures that will not work are:

· Boiling off the water from commercial bleach (NaOCl decomposes, NaCl/NaOH impurities)

· Dissolving Na in MeOH then boiling off the alcohol (NaOMe decomposes)

· Dissolving Al in HCl then boiling off the HCl solution (AlCl3*xH2O decomposes to oxychlorides)

        Things to watch out for consist of azeotrope distillation when involving liquids and carry over of less volatile insolubles.  Additionally when a solution has nearly evaporated there may be a heavy precipitate on the bottom, this can cause 'bumping' in the flask which can bounce a flask off a hot plate or even crack it due to the pressure of the vapors rising though the precipitate.  To avoid the worst of this you can cool the solution when a precipitate starts to form then filter it then resume heating, or resort to magnetic stirring, or heat at a lower temperature then the boiling point of the liquid.  When creating a salt by reacting an acid with a metal/carbonate/hydroxide etc. be sure to use the stoichiometric quantity of acid if at all possible, excess acid will only have to boiled off resulting in noxious clouds that kill grass, other plants, your eyes, and lungs (this is assuming the acid is volatile, e.g. H2SO4 will not be easily volatile).

3.3 Filtering 

3.3a Selection of filter paper 

3.3b Gravity Filtration 

3.3c Vacuum Filtration

3.4 Chromatography

3.5 Electrolysis 

An amateur chemist can easily define electrolysis as any reaction that calls for the direct application of an electric current to a chemical, either on its own, or in solution, for purposes other then heating. Our point here is to understand how we can transform chemicals using electricity. The key word is “ions”. What is an ion? A quick and dirty definition would be: “It’s an electrically charged atom or molecule”. The positively charged are called “cations” and the negatively charged are called “anions”. Ions have different chemical and physical properties than the original atom or molecule. All anions have their own name, so Br^- is baptized Bromide and NO3^- is Nitrate. Cations with more than one possible charge also have names, so Cu+ has one electrons missing and it’s baptized Cuprous, Cu++ has lost two of them and becomes Cupric. [See table in section 1.3 for a list of common cations and anions]

        Many chemicals, including all salts, are made of opposite charged ions “glued” together by electrical forces. And in these cases electrolysis can help unglue them. There are basically three schools of electrolysis that you should familiarize yourself with.

Pure Compound / Molten Salt Electrolysis	This is the simplest form of electrolysis. Not in its practice, but in concept. A pure compound. ionic in nature, i.e., consisting of a cation and anion, is heated until it becomes liquid. Ionic liquids are good conductors of electricity and therefore are able to be electrolyzed directly. Once molten, electricity is applied and the compound breaks down into its constituent parts. For example, molten sodium chloride when subjected to DC electrolysis will break down into liquid sodium metal and chlorine gas.
Aqueous Electrolysis	The most common form of electrolysis for the at home chemist. An ionic salt is dissolved in water and a current applied. Depending on the reduction and oxidation potentials of your cations and anions you get different products. For example, electrolysis of water, with a little salt added to aid in conductivity will yield hydrogen and oxygen gas, excess salt and a higher current will yield hydrogen and chlorine gas.
Non-Aqueous Electrolysis	A compound is dissolved in any liquid other then water and current applied. Different products are possible under different conditions. Products are possible with non-aqueous electrolysis that are impossible in aqueous electrolysis or would require high temperatures for molten salt electrolysis. For example, it is possible to obtain sodium metal as a deposit in the electrolysis of sodium chloride in pyridine, whereas sodium metal would react instantly in water and molten salt electrolysis would require temperatures of several hundred degrees Celsius.

There are two different types of electricity available. There is alternating current (AC) and direct current (DC). Alternating current is the type of electricity that comes out of the wall, this is not good for electrolysis, alternating current changes which side of your cell is your cathode and which is your anode about 60 times a second, this means that almost nothing can be accomplished with it. If you have a cell full of water and something to make it conductive and put two electrodes into it, plugging it into the wall the only thing you will do is heat your solution to boiling with resistance heating, make a random explosive gas mixture above your water, and deform your electrodes. Therefore your only real choice for productive electrolysis is DC.

So where does DC come from? I will not detail the electronics here. There is plenty of information elsewhere on the internet. But you usually have two sources, converting the output of your wall adaptor to DC, or using batteries.

· Batteries: Not very useful, except for simple demonstrations. Those little square 9V batteries make me laugh! If you insist on using batteries, be a man, use 4 “C” size (big) batteries or a lantern battery or even a car battery, but if you go with the car battery then you’ll need a charger, and you could just use that directly anyways.

· Adapters: These are very common these days and most households have one or two spare ones from old equipments that broke. 6 to 9 volts is fine for most experiments and they usually deliver above 0.5A. Another common source of power that falls into this category packing a little more punch is the car battery charger. A good one can supply from 0 – 12 V and from 0 – 55 A, and can be procured cheaply from second-hand stores. Old ATX power supplies are marvellous for electrolysis because they yield high amperages (up to 20A) at low voltages (5 and 12V).

· Build your own simple power supply: If you have some skills on electronics, you can built your own power supply just using a transformer and a single 1N4001 diode.

Build your own variable power supply: An extra potentiometer (100Kohms) and a power transistor like TIP 31 (3A, 40W), TIP 41 (6A, 60W)or TIP3055 (15A, 90W; a.k.a. 2N3055), can give you control over the voltage supplied by the batteries, the adapters or your home built power supply. Don’t expect precision or stability though:

A power source capable of delivering at least 0.5 ampere would be nice. The product yield per hour depends on the current, so your current should be proportional to your hurry and your electrode surface area (too much current per square centimeter may cause unwanted results).

The electrolysis process is sensitive to the voltage applied, but if you want a “rule of thumb” number, 9V will perform most tricks. 

An exercise in calculating yields involving electrolysis.

Let’s say that you want to make bromine, and just to simplify things let’s say you have some lead (II) bromide laying around to perform molten salt electrolysis on.

First off you divide your reactions into half reactions, the half reactions, when added together cancel out but separately they give the number of electrons necessary and help in visualization.

Pb²⁺ + 2e^- Þ Pb_(l)

2Br^- Þ Br_2(g) + 2e^-

Notice how the number of electrons on one side of the equation match the number of electrons on the other side of the other equation. That is because while one thing is reduced (gains electrons) another thing must be oxidized (loses electrons) a good way to remember this is the mnemonic OILRIG [Oxidation Involves Loss, Oxidation Involves Gain (of electrons)]. Couple this with the mnemonic RedCat [Reduction occurs at the cathode] and you can figure out where your products will be produced at. When wires are color coated in their normal manner the cathode is the black wire, the negative (-) wire. By default then the red wire is the positive (+) wire, is the anode.

Now that you know your reaction, what kind of power source are you using? Maybe you’re using a simple wall adaptor, possibly from a phone charger and the phone broke. Looking on the plastic you might even find that it is 6.2 V and .5 A. That’s all the information you need off that. The next step is to calculate the number of coulombs ( C ) of electricity that pass though the molten mass. A coulomb is a unit of measurement specific to things involving electricity, it is equal to one amp times one second. So it actually measure the quantity of charge moving though the cell. So let’s say that we are going to run this for 30 minutes:

.5A ^. .5h ^. (3600s/1h) ^. (1C/1A^.s) = 900 C

To explain the above equation you can see the that .5A came from the power of the power supply, the .5 hours came from the time the cell was running, the other numbers are conversions to the number of seconds in an our and the number of amp-seconds in a coulomb. This is the generic representation and you can just plug in your numbers in the above equation to get your own unique answers. Now we see from the two equations way overhead that for every one mol of Br₂ generated two mols of electrons are simultaneously brought into existence. Now we need a new unit of measurement for electrolysis, a Faraday (F). A Faraday is the number of electrons necessary to reduce one mol of a single charge unit. A Faraday is equal to 96,500 C. From here we use a new equation:

g. Br₂ = 900 C ^. (1 F / 96,500 C) ^. (1 Mol Br₂ / 2F) ^. (159.8 g Br₂ / 1 Mol Br₂) = 1.5 g Br₂

In the above equation we took the number of coulombs that we got from the first equation and multiplied it by the conversion for Faradays to coulombs and that by the mols of Br₂ and the number of Faradays involved, e.g., in this case two mols of electrons are involved and therefore there are 2 F. This is all multiplied by the grams of Br₂ per mol to give 1.5 g of Br₂ produced overall. To get lead you could convert the grams of Br₂ produced and make it into mols, then you could multiply that by the grams per mol of lead or just substitute that information into the last part of the above equation.

Well, most of you are thinking, “Only 1.5 g…. what the heck, I wanted a liter!” Well, this is not the setup you would use for massive Br₂ production. But electrolysis is good for producing small quantities of hard to obtain chemicals, and you could increase your yields by either increasing the amps of your power supply or by running the setup longer.

Aside from a power supply the second most important consideration are the electrodes. Many of the metals are similarly conductive for all intents and purposes and therefore it is better to consider them in terms of their chemical resistance relating directly to whatever environment you are planning to perform electrolysis in and the ease of procurement of the electrode material.

· Copper can be easily obtained from common wires. Copper has a wide resistance to many environments.

· Zinc can be found as the outer metallic shell of common batteries (the cheap ones, called carbon-zinc). Not good for acidic environments or basic environments.

· Carbon or graphite: These are very useful electrodes, since they do not oxidize as anodes. Well, they don’t last forever; they are attacked by oxygen, originating CO2, or maybe just disaggregate in the solution, but are much, much, cheaper than platinum electrodes, so they are widely used. The most common are pencil’s graphite. These have very different compositions, and may or may not last long. In fact, they may even be very bad conductors. Another source of carbon electrodes is the carbon rod that every carbon-zinc battery has inside. They are better than pencil’s graphite. The best carbon electrode I found is a rod of graphite covered with a copper layer found in solder’s supply shop. When you strip the copper with ferric chloride (or electrolysis), a resistant graphite electrode is left. Another option is graphite from electric motors sliding contacts. They are small, but easy to find and quite resistant. I have a couple from large polisher that are 23x16x6 mm.

· Lead is used often as an inert electrode. You will probably find it in a sporting goods shop: fishing weights, gunshots in general, airgun bullets etc. It melts easy and you can cast your electrodes melting it with a blowtorch and a scoop.

· Nickel: Coins from some countries contain an appreciable amount of nickel and can be used as electrodes, additionally can be procured from scrap yards or specifically for electrolysis. Nickel is an excellent material for electrolysis of highly basic solutions.

· Iron: Iron is attacked readily under acidic conditions and somewhat slower under basic, but it is commonly available and may find some use in a pinch.

· Platinum: If you have the money, it’s the most resistant anode I know. Fine wire is ok, but for larger surfaces, I hear people use other metals plated with platinum.

· Silver/Gold: A cheaper alternative to platinum somewhat more reactive. Available from coin suppliers in the form of collectable currency either can be melted and cast into the appreciable shapes desired. Silver wire can be purchased from jewelers as can platinum.

· Misc. Electrodes: Mercury, lead dioxide, rare earth oxides plated on titanium, there are hundreds of possible electrodes that one may come across that are not covered here.

3.5a Molten Salt Electrolysis

Performing electrolysis on a molten binary salt usually yields predictable products providing you have a simple anion coupled to your metal ion. Lead bromide as provided in the example yields lead metal and bromine as a gas.

3.5b Aqueous Electrolysis

By Tacho.

1- Ions:

         Water and other solvents called “polar solvents” do something interesting: their molecule has two poles, a positive charge on one side and a negative charge on the other. When you put a salt in water, the salt’s negative ions are surrounded by water molecules pointing their positive side to it. Of course, the positive ions get surrounded by water molecules pointing their negative side to it. The result is that the salt dissolves in a liquid “soup” of ions that we will call “solution” or “electrolyte” from now on.

          There must be enough anions to neutralize all cations. The whole thing must be neutral. Yes, it can built up charges, and momentarily have an imbalance, but mother nature will find a way to put thing back the way she likes it.

         In the solution, you don’t have the individual salts anymore. Just ions. So, if you put two salts in water, say... sodium chloride and potassium bromide, you only have sodium cations, potassium cations, chloride anions and bromide anions. You can extract from this solution the original salts or sodium bromide and potassium chloride! If you could substitute both of those anions for the hydroxide anion, than you could extract sodium hydroxide and potassium hydroxide from the same solution.

         That’s what this work is all about: the substitution of ions in aqueous solutions using electrolysis. We do that by pushing electrons in and out of atoms or molecules using electricity.  

2- How?

        Suppose I want to get a copper salt, dissolve one of my electrodes. I mount the following setup, put distilled water in the flask and turn on the power:

        What do I get? Nothing! Why? Pure water practically does not conduct electricity! So we must add something that will make water conductive but won’t be part of the reaction. Let’s say we want copper sulfate, so let’s choose something with a sulfate ion attached to an “inert” ion. MgSO4 should work fine. Magnesium sulfate is Epson salt that every pharmacy should have. For now, take my word that Mg won’t be part of the reaction. Lets dissolve 3g of MgSO4 (the hydrate is okay.) in 30ml of distilled water and try again.

        Immediately you’ll see bubbles on the cathode and some blue tint by the anode. Cool! Copper ions should be blue! What is in those bubbles? The only thing in this soup that becomes a gas when it gains electrons (that’s what cathodes do, they give electrons) is hydrogen. Water molecules get broken and hydrogen bubbles away.

But wait... after a couple of minutes we notice that something is very wrong! The blue tint is solid and it’s precipitating! Copper sulfate doesn’t do that! What is happening?

        Here goes the explanation: When the hydrogen of water receives an electron, it joins a friend to become a gas molecule and bubbles away. It leaves behind the other half of the water molecule: OH- ion. This ion meets Mg+ ions or Cu++ ions and forms an insoluble hydroxide that precipitates (actually copper ions form more complex insoluble compounds, but let’s pretend it’s plain hydroxide). If you carefully put a piece of indicator paper close to the cathode while its bubbling, you will see that it’s alkaline.

        So how do we make our copper sulfate?

        The little snotty armchair chemist now smiles and says with his squeaky voice: “I know! I know! All you have to do is to use two half-cells connected by a salt bridge! It’s so simple! Like this:”

        In the real world, however, this setup is not efficient. If you use the 10g MgSO4 in 100ml water solution, under about 10 volts, you just can’t see any bubbles! That’s because ions are not agile movers, the salt bridge is a long way for them. Like any electric circuit, shorter paths increase current. So, a better design for the amateur would be:

        I tested this for the copper sulfate production and it works. For some reason, it works best using the lower part as the anode cell, where the sulfate forms. The solution by the anode is called anolyte and now is a mixture of Mg²⁺, Cu²⁺ and SO₄^2- ions. How to separate pure CuSO4 from it? If you look up the solubility information for the two salts you will find that 100 ml of cold water will dissolve 31 g of hydrated copper sulfate whereas it will dissolve 71 grams of magnesium sulfate therefore if you ran your electrolysis long enough then heated to reduce the volume of the solution, then cooled, the first thing to come out of solution should be the copper sulfate. Magnesium sulfate is soluble in glycerine, but copper sulfate is not, heating the solution and evaporating till it became a solid cake, then powdering and mixing that with glycerine and filtering off should make the copper sulfate purer, however there is always a new method to try. I’ll pick some pure copper sulfate I bought in a gardening shop to do the next experiment.

In this setup the cotton acts as a membrane to prevent the passage of certain ions. Another type of system accessible to the at home chemist consists of a flowerpot, unglazed and previously soaked in a strong acid placed into a larger container. The flowerpot and the larger container are filled so they contain the same level of electrolyte and an electrode is inserted in the flowerpot and another in your main solution outside. This is yet another way to separate your anolyte to prevent extraneous reactions.

        I substituted the two copper wires of the first setup (the one with the connector) by two pencil leads (graphite) and put them in a pure copper sulfate solution. When I turn on the power, quickly something starts building up in the negative electrode. That’s copper. It’s probably powdery and maybe too dark to look metallic, but it’s copper. That’s the principle behind electroplating. Don’t expect shiny metallic deposits though.

        What’s happening is that the Cu²⁺ cations in the solution gain a couple of electrons and become Cu, the metal.

3- Giving names to things happening.

What is happening in the anode is called oxidation. It doesn’t matter if that there is no oxygen involved! Loosing electrons by the anode is called oxidation and don’t you argue! Usually this oxidation results in metals becoming ions, ions gaining oxygen atoms, oxygen evolving in bubbles or halogen ions becoming the element.

        What is happening in the cathode is called reduction. It’s usually a metal ion becoming a metal, ions loosing oxygen atoms, elemental halogens becoming ions or hydrogen evolving in bubbles.

4- What happens when?

Question: What happens when I have many different metal pieces in the same electrolyte, connected and acting as anodes? Do they all go into solution at once? Do some of them go first? And what happens when I have many different metal cations by the cathode? Do they all get reduced together forming an alloy? 

Answer: There is a priority list. Theoretically, if you have two different metal pieces, like copper and zinc connected to the positive pole of a battery and immersed in an electrolyte, all zinc will go into solution before the first atom of copper gets oxidized.

Here is an incomplete list (potential table) that shows the tendency of an atom or molecule or ion to gain or loose electrons. If it shows a reduction potential of –2.71 for something, the oxidation potential for the same something will be +2.71:

Oxidized creature	Reduced creature	Reduction potential:
Li^+	Li	-3.04
K^+	K	-2.92
Ca^2+	Ca	-2.76
Na^+	Na	-2.71
Mg^2+	Mg	-2.38
Al^3+	Al	-1.66
2 H2O	H2(g) + 2OH-	-0.83
Zn^2+	Zn	-0.76
Cr^3+	Cr	-0.74
Fe^2+	Fe	-0.41
Cd^2+	Cd	-0.40
Ni^2+	Ni	-0.23
Sn^2+	Sn	-0.14
Pb^2+	Pb	-0.13
Fe^3+	Fe	-0.04
2H^+	H2(g)	0.00
Sn^4+	Sn^2+	0.15
Cu^2+	Cu^+	0.16
ClO4^- + H2O	ClO3^- + 2OH^-	0.17
AgCl(s)	Ag + Cl^-	0.22
Cu^2+	Cu	0.34
ClO3^- + H2O(l)	ClO2^- + 2OH^-	0.35
IO^- + H2O(l)	I^- + 2OH^-	0.49
Cu^+	Cu	0.52
I2(s)	2 I^-	0.54
ClO2^- + H2O	ClO^- + 2OH^-	0.59
Fe^3+	Fe^2+	0.77
Hg2^2+	2Hg(l)	0.80
Ag^+	Ag	0.80
Hg^2+	Hg(l)	0.85
ClO^- + H2O	Cl^- + 2OH^-	0.90
2Hg^2+	Hg2^2+(aq)	0.90
NO3^- + 4H^+	NO(g) + 2H2O	0.96
Br2(l)	2Br^-	1.07
O2(g) + 4H^+	2H2O	1.23
Cr2O7^2- + 14H^+	2Cr^3+ + 7H2O	1.33
Cl2(g)	2Cl^-	1.36
MnO4^- + 8H^+	Mn^2+ + 4H2O	1.49
H2O2(aq) + 2H^+	2H2O(l)	1.78
Co^3+	Co^2+	1.82
S2O8^2-	2SO4^2-	2.01
O3(g) + 2H^+	O2(g) + H2O	2.07
F2(g)	2F^-(aq)	2.87

        One should be able to get all the information one needs from this table, but lesser creatures like me always get confused trying to draw practical conclusions from it! So I organized the following two specific tables for things that can happen at your anode and things that can happen at your cathode. I believe they are foolproof:

Things that can happen at your anode (where oxidation takes place, positive pole in the electrolytic cell) in order of priority:

Priority a.k.a. Oxidation Potential in Volts:	Situation:	What I think should happen: What happened when I tried it:
higher than +0,76	If your anode has alkali metals, Al or Mg ...	...Alkali metals react with water, and can’t be electrodes in aqueous solutions. Aluminum behaves strangely as an anode (see “anodization” procedures). I just don’t know about Mg.
+0.76	If your anode has Zn (Zinc)...	...Zn^2+ ions go into solution. It may form insoluble salts (precipitate or form chunks in the anode) if it finds certain anions in the electrolyte like carbonate, hydroxide or hypochlorite. The Zn in the anode “dissolves” in the electrolyte. It works. Nice and easy. Using a two half- cell setup with MnSO4, the zinc anode dissolves into Zn^2+ ions, a colorless clear solution. Some black residue is left though, I don’t know what that is.
+0.74	If your anode has Cr (Chromium)...	...Cr^3+ (chromic) ions go into solution. It may form insoluble salts (precipitate or form chunks in the anode) if it finds certain ions in the electrolyte like borate or cyanide. The Cr in the anode “dissolves” in the electrolyte.
+0.41 or +0.04	If your anode is made of Fe (Iron)...	...Fe^2+(ferrous) ions go into solution at “priority” of +0,41V, or Fe+3 at “priority” of +0.04V. It may form insoluble salts (precipitate or form chunks in the anode) if it finds certain exotic ions in the electrolyte. The Fe in the anode “dissolves” in the electrolyte. See -0.77 priority. Funny things happen with an iron electrode. In my MgSO4 cell, at voltages lower than about 5 volts, a yellow/brown tint and a greenish precipitate appear due to iron ions(?), but if the voltage is above that, oxygen evolves from the anode and almost no color shows in the solution. If a piece of zinc from a battery is strapped to the anode, almost no brown color evolves .Makes sense.
+0.40	If your anode is Cd (Cadmium)...	...Cd^2+ ions go into solution. It may form insoluble salts (precipitate or form chunks in the anode) if it finds certain ions in the electrolyte like carbonate or borate. The Cd in the anode “dissolves” in the electrolyte. This is a nasty heavy metal, by the way. You shouldn’t be messing around with it!
+0.23	If your anode has Ni (Nickel)...	...Ni^2+ go into solution. It may form insoluble salts (precipitate or form chunks in the anode) if it finds certain ions in the electrolyte, like carbonate or borate. The Ni in the anode “dissolves” in the electrolyte.
+0.14	If your anode has Sn (Tin)...	...Sn^2+ (stannous) ions go into solution. It may form insoluble salts (precipitate or form chunks in the anode) if it finds certain ions in the electrolyte. The Sn in the anode “dissolves” in the electrolyte. See -0.15 priority.
+0.13	If your anode has Pb (Lead)...	...Pb^2+ (plumbous) ions go into solution. It may form insoluble salts (precipitate or form chunks in the anode) if it finds many common ions in the electrolyte, like sulfate, chloride or carbonate. The Pb in the anode “dissolves” in the electrolyte. Another heavy metal you should be careful about.
0.00	If you make the incredible “hydrogen bubbling on platinum” electrode...	...Literature is unanimous that hydrogen would become an H^+ ion.
-0.15	If your anolyte (the solution of ions around your anode) has Sn^2+(stannous) ions...	...The Sn^2+ become Sn^4+ (stannic) ions.
-0.16	If your anolyte has Cu^+(cuprous) ions...	...The Cu^+ become Cu^2+ (cupric) ions.
-0.17	If your anolyte has ClO3^- (chlorate) and OH^-(hydroxide = alkaline medium)...	...The ClO3^- ion becomes the ClO4^- (perchlorate) ion and some extra water is formed. Perchlorate is the sweet darling of pyrotechnics. I have not done it, but the whole electrolysis sequence Cl- -> ClO -> ClO2 -> ClO3 -> ClO4 is well described in the internet.
-0.22	If your anode has Ag (silver) and your anolyte has Cl^- ions...	...Insoluble AgCl (Silver Chloride) is formed. The AgCl forms an insulating white layer on the silver piece that reduces the current sharply. It resembles white (well, maybe cream) paint that darkens if exposed to sunlight for long.
-0.34 or -0.52	If your anode has Cu (Copper)...	...Cu^2+ go into solution. Or Cu^+ at –0.52. It may form insoluble salts (precipitate or form chunks in the anode) if it finds certain ions in the electrolyte, like carbonate or hydroxide. The Cu in the anode “dissolves” in the electrolyte. My results with a MgSO4 cell are described in the text.
-0.35	If your anolyte has ClO2^- (chlorite) ion and OH^-(hydroxide = alkaline medium)...	...The ClO2^- ion becomes the ClO3^- (chlorate) ion and some extra water is formed. See note for chlorate (oxidation potential –0.17).
-0.49	If your anolyte has I^- (iodide) and OH^-(hydroxide = alkaline medium)...	...The I^- ion becomes IO- (hypoiodite) ion and some extra water is formed.
-0.54	If your anolyte has I^-ion.	I2 is formed. It does. In a potassium iodide solution, using pencil’s graphite electrodes, brown iodine color rapidly develops around the anode while hydrogen bubbles form at the cathode. Interesting note: If you apply 60Hz AC directly from the transformer (6V) into the electrodes, a yellow color slowly develops, showing this reaction is not fully reversible. At least mine wasn’t. Also interesting: according to this list of priorities, a copper anode should not generate iodine, instead, the copper should dissolve into copper ions. That did not happen in my test. Brown color developed near the copper anode with no blue tint. Go figure!
-0.59	If your anolyte has ClO^- (hypochlorite) and OH^-(hydroxide = alkaline medium)...	...The ClO^- ion becomes the ClO2^- (chlorite) ion and some extra water is formed. See note for chlorate (oxidation potential –0.17).
-0.77	If your anolyte has Fe^2+ (ferrous) ion...	...The Fe^2+ becomes Fe^3+ (ferric).
-0.80	If your anode has Hg (Mercury)...	...Hg^2+ go into solution. It may form insoluble salts (precipitate or form chunks in the anode) if it finds the certain ions in the electrolyte. The Hg in the anode “dissolves” in the electrolyte. Dangerous heavy metal, by the way. Mercury, besides being toxic, have some chemical behaviors that I don’t understand.
-0.80	If your anode has Ag (Silver)...	...Ag^+ go into solution. It may form insoluble salts (precipitate or form chunks in the anode) if it finds certain anions in the electrolyte, like hydroxide, carbonate, iodide, bromide or chloride. Silver has many insoluble salts. The Ag in the anode “dissolves” in the electrolyte.
-0.90	If your anolyte has Cl^- (chloride) and OH^-(hydroxide)...	...ClO^- (hypochlorite) ion is formed. And some extra water. See note for chlorate (oxidation potential –0.17).
-1.07	If your anolyte has Br^- (bromide) ion...	...Bromine (liquid) is formed. It does. In a sodium bromide solution, using pencil graphite electrodes, brown bromine color rapidly develops around the anode while hydrogen bubbles form at the cathode. This one is fully reversible: If you apply 60Hz AC directly from the transformer (6V) into the electrodes nothing seems to happen, although bromine and hydrogen are being oxidized and reduced 60 times per second! Also: If you use a copper anode, no brown color develop, only a bluish precipitate (bromide? hydroxide?). This is expected, since copper is above the bromide ion in the priority list.
-1.23	whatever aqueous...	...At this point, in aqueous solutions, water molecules nearby the anode are split in O2 (gas), that evolves from the anode, and 4 H^+ ions, that remain in the solution. Theoretically, no oxidation should take place beyond this potential, because water should get oxidized first. Ha ha! So much for theory! Something called overpotential makes, among others, the next two oxidations possible:
-1.36	If your anolyte has Cl^- (chloride)...	...Chlorine gas evolves. This one sure works. The electrolysis of NaCl solution gives off hydrogen and chlorine. I believe (have not tested) that if you keep the voltage low enough, you can get a small amount of oxygen and no chlorine. But that may depend on the anode composition.
-1.49	If your anolyte has Mn^2+ (manganous) ion...	...Mn^2+ becomes MnO4^- (permanganate) ion. I have not tested this one personally, but I read in a book a detailed description of a laboratory procedure to obtain potassium permanganate by electrolysis. The book seems trustworthy.

Things that can happen at your cathode (where reduction takes place, negative pole in the electrolytic cell) in order of priority:

Priority a.k.a. Reduction Potential in Volts:	Situation:	What I think should happen: What was actually observed:
1.49	If your catholyte (the solution of ions around your cathode) has MnO4^- (permanganate ion) and some extra H+ (acidic solution)...	...MnO4^- turns to Mn^2+ (manganous) and some extra water are formed.
1.07	If your catholyte has Br2 (liquid bromine)...	...Br^- (bromide) ion is formed. As I said before, if you apply 60Hz AC directly from the transformer (6V) into graphite electrodes in a sodium bromide solution, nothing seems to happen, although bromine and hydrogen are being oxidized and reduced 60 times per second! So, it works!
0.96	If your catholyte has NO3^-nitrate ion and your solution is acidic…	...NO gas and some extra water are formed.
0.90	If your catholyte has ClO^- (hypochlorite) ion...	...Cl^- (chloride) and OH^-(hydroxide) are formed.
0.85	If your catholyte has Hg^2+ ions...	...Hg (metallic mercury) is formed.
0.80	If your catholyte has Ag^+ ions...	... Ag (metallic silver) is formed.
0.77	If your catholyte has Fe^3+ (ferric) ions...	...Fe^3+ becomes Fe^2+ (ferrous) ion
0.59	If your catholyte has ClO2^- (chlorite) ion...	...ClO2^- becomes ClO^-(hypochlorite) ion
0.54	If your catholyte has I2 (iodine) dissolved...	...Iodine becomes I^- (iodide) ion. As I said before, if you apply 60Hz AC directly from the transformer (6V) into graphite electrodes in a potassium iodide solution, a yellow color slowly develops, showing this reaction is not fully reversible. But the fact that is a slow development shows that most iodine was oxidized and reduced 60 times per second. So it works! However, a practical method of turning elemental iodine into I- is beyond my reach . I tested a solution of ethanol, iodine and MgSO4 and another with ethanol, iodine and NaCl. None seem to have any change under low voltage, At higher voltages, hydrogen evolved but no change in color. The anode was graphite.
0.52 or 0.34	If your catholyte has Cu^+ (cuprous) or Cu^2+ (cupric) ions...	...Both get reduced to metallic copper. Cu^+ at 0.52 priority and Cu^2+ at 0.34 priority. This one is easy. But to get a shiny smooth deposit is an art. As expected, when I did electrolysis using graphite electrodes in a solution of copper sulfate and zinc chloride, only copper deposited in the cathode (I tested it regularly dipping the electrodes in 20% HCl where copper does not react, but zinc bubbles). Only when almost (!) all blue tint from copper ions have gone from the solution, zinc starts depositing. This is how copper is purified industrially. Copper deposits before most common metals. Chlorine evolved from the anode, by the way, so I guess zinc sulfate was left in solution.
0.49	IO- (hypoiodite)...	...Becomes I- (iodide) and OH-.
0.35	ClO3^- (chlorate) ion...	... becomes ClO2^- (chlorite) ion
0.17	ClO4^- (perchlorate) ion...	...becomes ClO3^- (chlorate) ion.
0.15	Sn^4+ (stannic)...	...becomes Sn^2+ stannous ion.
0.00	If your catholyte is acidic and therefore has H^+ ...	H2(g) is generated. Therefore, none of the following metal ions can be reduced in acidic electrolyte. Which makes sense, since these metals are attacked by acids. In fact, they are attacked by acids exactly because they have a lower reducing potential then H^+, but that’s another story.
-0.04	Fe^3+ (ferric) ...	metallic Fe deposits at the cathode. No it doesn’t. At least not in my ferric chloride solution with graphite electrodes. I could get no deposit.
-0.13	Pb^2+ (plumbous)...	metallic Pb deposits at the cathode.
-0.14	Sn^2+ (stannous)...	metallic Sn deposits at the cathode.
-0.23	Ni^2+...	metallic Ni deposits at the cathode.
-0.40	Cd^2+...	metallic Cd deposits at the cathode.
-0.41	Fe^2+(ferrous)...	metallic Fe deposits at the cathode.
-0.74	Cr^3+ (chromic)...	metallic Cr deposits at the cathode.
-0.76	Zn^2+...	metallic Zn deposits at the cathode. It works. The deposit from a zinc chloride solution is powdery and dark, but it’s zinc allright.
-0.83	Whatever aquoeous...	Water by the cathode gest split in H2(g) that evolves from the cathode and + 2OH-, that remain in the solution. End of the line. No exceptions and no reductions from here!
	Aluminum ions, Magnesium ions and alkali metals ions...	Forget it! Not in aqueous electrolytes! Water gets reduced first!

5- Practical notes for amateur experiments:

c) What can be made using aqueous electrolysis?

Electrolysis is inefficient and slow. Consumes lots of electricity and takes a long time to produce little yields that, in most cases, must be submitted to other procedures to isolate the pure products.

If you are hoping to make a bottle of bromine or iodine in one sunny Thursday afternoon, you will be very disappointed.

On the other hand, it’s simple and within the reach of any amateur. It’ a good option when you need just a little bit of a chemical; instead of buying half a kilogram of an expensive, polluting and carcinogenic salt, make a little bit as needed. Also, it’s a way of obtaining chemicals that you just can’t buy.

· Salts in general can be obtained like sulfates, chlorides, nitrates, chlorates and perchlorates;

· Hydroxides and oxides;

· Metal powders and metal deposits on surfaces;

· Gases like oxygen, hydrogen and chlorine;

· Bromine and Iodine;

· Dilute acids;

· Organics like ethane, chloroform, etc.

3.5c Non-Aqueous Electrolysis

Additional notes for electrolysis:

1. Electroplating is the process of putting a thin layer of a metal onto a metal or sometimes, non-metal but rendered conductive substrate. The more advanced form of this is a process known as electroforming, whereby an object is coated with a thick mechanically sound layer of metal, non-metal objects can be formed from wax or other materials then coated with enough metal to make them useable for machining purposes. However electroforming in the home lab is very difficult.

3.6 Titration

3.7 Temperature control/Measuring

        Temperature control varies in its importance from reaction to reaction.  Two extreme examples would be a nitration reaction where you are controlling the temperature of the reaction to a range from 10C to 30C and at the total opposite might be trying to make phosphorus where your temperature might be 1200C and you are trying for as high a temperature as possible.  Both of these reactions pose their own difficulties for both measuring the temperature and controlling it.

        Your basic reactions are readily controlled with water in some form or another.  Water has a high specific heat and it can absorb a lot of heat before rising in temperature, or conversely, it can hold enough heat to warm another mass significantly before loosing its full heat.  So a cold water bath might be good for cooling, a warm/hot water bath for heating, and ice is always good to have around just in case.  However the other time tested method of cooling mixtures even cooler are euccentic mixtures.  KOH with ice can achieve extremely low temperatures, HNO3 and ice can too.  As can the dry ice/acetone slurry that is occasionally used.  Even cooler mixtures can be obtained with hydrocarbon baths with liquid nitrogen added periodically.  Liquid nitrogen could even be used directly, much lower then this is hard to accomplish in a home lab, but the lowest temperature coolant you will run across would be liquid helium, but I doubt you'll find this in any local super market.

        On the opposite side of the temperature scale you are shooting for heating.  Most any heating you will be doing will be the work of either electric heating, as in a hot plate, or chemical heating in the form of combustion.  Commonly electric heating concepts can get to 250C or so, lab grade hot plates can get even higher.  But to get really high you will have to use combustion, lacking a suitable furnace that is.  There are different kinds of torches, and different kinds of burners, Merc burners, bunsen burners, and more each have their own limitations.  Common butane torches can reach 800 - 900C but to get higher the use of MAPP gas can bring you there with an appropriate torch head.  Methane is also a good carrier of potential energy so hooking directly into your home gas line has its advantages.

        Measuring these temperatures though pose their own difficulties.  Common thermometers are fine for common temperatures, -30C - 340C can be found on many thermometers, but beyond this on either side of the temperature scale it becomes necessary to deviate from the norm.  Bimetal thermometers that hook into electrician tools can reach 1000C but you will have to invest in a good probe and the thermometer itself is sensitive to chemical attack.  Another type of thermometer exceptionally good for high temperature is the infrared thermometer.  Just point and click and you get an non-invasive temperature of up to 1000C for upper class models, but you will pay high for this too, although they are quite useful.

        The most useful tool for high temperature measurement available to the amateur chemist is the melting point of other compounds.  Molten metal baths can give an approximation of the temperature being used to keep the bath molten.  Different substances can be found to give a wide variation of temperature baths, just be wary of decomposition problems at these temperatures.

3.8  Removing water from gasses or solids with drying agents

3.9  Recrystalization

When the purity of a product is in question and you can spare a little bit in the quest for a more absolute product something can usually be worked out with recrystalization.  The basic principle is to pick out a solid that you want purified, say, ammonium nitrate.  The next step is to find a suitable solvent for it.  The solvent should posses the following properties:

1. Be able to dissolve a large quantity of the desired product when hot and have only a low solubility when cold.
2. Not affect the product be it by causing it to decompose or react with it.
3. Be useful under atmospheric conditions, not possess properties that are affected by heating and cooling.

        Finding such a solvent is usually difficult.  Many places do not list the solubility of a substance in anything but water, let alone finding hot and cold solubilities of a substance of different solvents on the same page.  As such some trial and error may be involved, or research can help, finding out what solvent a pioneer in the field used to recrysatalize one of your products may be a good start.

        Once your solvent is picked out, in the case of ammonium nitrate water can be used.  The first step is to heat the water to a high temperature but not quite boiling, then to saturate the solution with as much of the solute as can be dissolved, if there is still solid solute in the solution either it can be spooned out or more solvent can be added.  After the substance is all dissolved and while still hot a quick filtration can be used to remove insoluble materials such as glass/sand particles.  Careful though, a hot saturated solution can crystallize on a cooler solid surface quickly and can plug filter paper, have something to scrape the filter paper with handy.

        After your hot solution has been quickly filtered you let it cool.  Usually once a certain temperature is reached crystals will automatically start to come out of solution.  However on occasion a solution may become super-saturated, i.e. the solution should have crystals forming but there is nothing for them to form on, that is one way to look at it.  In this case one of two things can be done, a crystal of the original compound can be added, this is called a seed crystal, and new crystals will grow off it.  Or you can scrape the inside rim of your container right at the liquid air interface, this agitation can cause the growth of crystals.

        Allow the solution to keep cooling but don't mechanically cool it too low, if for example you cool a water solution to near 0C then most of the impurities may crystallize out with your intended product defeating the purpose of recrystalization.  But after you get a good crop of crystals then filtration is the logical next step.  Vacuum filtration is the best as you may be filtering off a large quantity of solid but gravity filtration may work depending on your circumstances.  Your freshly grown crystals should be washed while still in the filter with a few quantities of cold solvent to remove adhering particulates.

        Now that you have your crystals they may need to be dried, in a dessicator or under high vacuum are the two normal choices, the dessicator being the most readily available to the amateur chemist.  But this step may not be necessary, check in a chemistry book to see if the salt you produce is hydrated, if that is the case heating the salt will usually be necessary to create the anhydrous product, if your product is disquecent or hygroscopic a dessicator may be a good first choice, keeping it there for a few days may prove to be a good move, followed by immediate storage in an air tight container to prevent the re-entry of water.

        In addition to cooling a solution to cause a precipitate of crystals another solvent can be added to the solution.  In this case the solubility of a solid is less in both the solvents then it is in either one alone.

3.1k  Measuring Weight and Volume

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4.0 Lab Reagent Types (Intro, discuss overlap, and generalization) 

Rather then attempt to give you a chemical, name it, give you its properties, have you memorize those and move on to the next one I have organized this area to help you learn chemical properties more readily.  There is a bit of generalization here and some overlap, however rather then learning the chemical then the properties, the purpose of this is to tell you the properties then give you a list of the chemicals that posses these properties along with a bit of relevant data for each.  In doing this it is easier to go into deeper detail on exactly the designated title means not only under STP (Standard Temperature and Pressure) but also under extraneous conditions that you might be required to work with them at.

4.1 Acid / Base Theory (Aqueous Solution)

Discuss pH and pKa

4.2 Acids (Organic/Inorganic)

        The most loose definition of acids that most people are familiar with defines an acid as a chemical that is able to donate a hydrogen cation to water.  In doing this it generates the H3O+ cation which is the acidic component of water.  However this does not cover every acid under every circumstance by a long shot.  Never the less, water is a common solvent and in defining an acid it is easier to use definitions governed by water then add in the exemptions later for non aqueous systems.

Common Acids

Acetic Acid     CH3COOH	Commonly known as vinager, this acid forms no confirmed azeotrope with water.  It is somewhat strong in concentrated form, dissociating to an appreciable extent.  Acetate salts are usually soluble and are therefore a good source of metal ions in solutions, however solutions are slightly basic.
Hydrochloric Acid      HCl (Muratic Acid)	Sold as a solution in water of HCl gas, hydrochloric acid is a strong mineral acid.  The commonly avalible forms are 20% (The azeotrope), 38% (concentrated with a density of 1.19 g/cm^3)  It will attack anything in the reactivity series above hydrogen, most chlorides are at least slightly soluble.
Nitric Acid
Sulfuric Acid
Sulfamic Acid
Boric Acid
Hydrofluoric Acid
Cyanuric Acid
Phosphoric Acid

Metal Activity Series:

Time to introduce you to the metal activity series. Although important to other chemistry concepts it answers one question regarding acids that people ask most often. “What will an acid dissolve?” Below is a list of elements, towards the end of the list is hydrogen, and anything to the left of it will dissolve to some extent in acid. Those in the lighter color to the immediate left dissolve slowly-very slowly, going to the darker color even more to the left we find elements that will not only displace hydrogen from an acid but will react with steam. Finally those furthest to the left will readily react with water and their subsequent reaction with acid would only be described as intensely violent. This is a basic activity series, some series will have elements in slightly different relation to one another but this is the basic order.

Li K Ba Ca Na Mg Al Mn Zn Cr Fe Cd Co Ni Sn Pb (H₂) Cu Ag Hg Pt Au

So you’re looking at the list and you wonder, “What about those elements to the right of hydrogen?” A good question, those elements will not displace hydrogen from acid and as a consequence they could be considered inert in that respect. But that would be a mistake to assume they would remain inert in all respects. There is a way around this inertness, the addition of an oxidizing agent. The principle, let’s say for example you have a piece of copper that has some surface oxidation, now let’s say you put it into some hydrochloric acid, immediately the oxidized layer dissolves off tinting the acid a green/blue color. Pulling out the copper it looks fresh and clean, no oxidation. So, the oxidized layer dissolved, so if you were to leave it out the oxygen in the air would re-oxidize that top layer, you could dip it back into the hydrochloric acid, and dissolve yet more of the copper. In this case the atmospheric oxygen is the oxidizing agent, bubbling air though HCl while dissolving copper accomplishes this. But another way would be to add an oxidizing agent to your acid, an even better way would be to have an oxidizing acid. Perchloric acid (HClO4) and nitric acid (HNO3) are both oxidizing agents as well as acids [Hot concentrated H2SO4 is also an oxidizing agent], as a matter of fact nitric acid almost always functions as an oxidizing agent unless coupled with a very reactive metal, magnesium will actually liberate hydrogen for the first few seconds of reacting with nitric acid but after that it will be preferably oxidized first. Oxidizing acids will not dissolve some elements that have insoluble oxides, the formation of the oxide forms a protective layer pacifying the metal to further attack, a good example is aluminum in concentrated HNO3, also tin can be pacified in this way under some conditions.

Oxidizing Acids

By being an oxidizing agent the acid must simultaneously be reduced in the reaction. Therefore when copper is subjected to the action of nitric acid copper is oxidized and the nitrate anion is reduce to any of a number of nitric oxides depending on the conditions under which the oxidation took place. Here are some examples of the reactions of nitric acid:

2HNO_3(aq) + Mg_(s) Þ Mg(NO₃)_2(aq) + H_2(g)

Rarely occurs, only happens initially with magnesium or even more reactive metals [Na, K, Li, etc.], not important.

3Cu_(s) + 8HNO_3(aq) Þ 3Cu(NO₃)_2(aq) + 4H₂O_(l) + 2NO_(g)

This is an example of nitric acid acting as an oxidizing agent when dilute.

Cu_(s) + 4HNO_3(aq) Þ Cu(NO₃)_2(aq) + 2H₂O_(l) + 2NO_2(g)

This is an example of nitric acid acting as an oxidizing agent when concentrated. Notice the ratio of nitric acid molecules reacting with copper compared to the dilute reaction above.

P_4(s) + 20HNO_3(aq) Þ 4H₃PO_4(aq) + 20NO_(g) + 4 H₂O_(l)

Concentrated nitric acid can also oxidize elements such as phosphorus, silicon, sulfur, and occasionally carbon, especially when heated.

Fe_(s) + 6HNO_3(aq) Þ Fe(NO₃)_3(aq) + 3 H₂O_(l) + 3NO_2(g)

When metals capable of multiple oxidation states are dissolved in concentrated nitric acid they will usually take the highest normal oxidation state, in this case iron becomes +3 in preference to +2.

Similarly, when copper or mercury, some of the more reactive of the metals that follow hydrogen in the activity series, come into contact with hot concentrated sulfuric acid they can be oxidized and the sulfuric acid reduced.

Cu_(s) + 2H₂SO_4(l) Þ 2H₂O_(l) + SO_2(g) + CuSO_4(aq)

Perchloric acid is encountered to a considerably lessened extent in the laboratory, it has a nasty reputation for exploding for no reason, generating out of control reactions, creating fire hazards, and making unstable salts.

Another thing to consider when pondering weather a metal will dissolve in an acid is weather the salt formed would be soluble. One would not logically think that silver would dissolve in hydrochloric acid independent of its unreactivity simply based on the fact that the silver chloride thus formed is totally insoluble. Even a piece of barium metal tossed in H2SO4 may become pacified which is an amazing thing considering it would react very rapidly with water. Oxidizing ability aside there is another method to measure the strength of an acid, the pH scale and the pKa scale, which were discussed in the opening section.

4.3 Bases 

        As shown in the above picture bases can rapidly attack some metals just as acids can.  To the left in the above picture some aluminum turnings have been placed into a weak potassium hydroxide solution, to the right a weak acid solution is also attacking a similar amount of aluminum.  Hydroxides can attack a number of metals, especially when hot and concentrated, however the reactivity shown with aluminum, zinc, and magnesium can be considered special cases for the common metals.

Common Bases

Sodium Hydroxide  NaOH	Avalible over the counter as lye, sodium hydroxide serves the purpose of being the no-nonsense base, addition of sodium hydroxide to an aqueous solution automatically increases the hydroxide ion concentration and brings only the sodium cation along with it.
Sodium Carbonate  Na2CO3	Sodium carbonate is available as “Washing soda” it is usually the decahydrate (*10H2O) but that does not interfere with calculations as long as it is accounted for. Be wary of other impurities though. Sodium carbonate is a great base because the reaction with acidic components is driven foreword strongly by the loss of carbon dioxide from solution.
Sodium Bicarbonate NaHCO3	Less basic in solution then sodium carbonate but still able to neutralize acids well. It is safer on the skin and is therefore the choice base to have laying around in case of an acid spill.
Ammonia  NH4OH	Ammonia gas can simply be bubbled into solution to increase its pH. That is a great advantage to ammonia. Also it can be forced from solution after its purpose has been served, the gas itself will react with acids even if they are not aqueous either.
Trisodium Phosphate Na3PO4	Basic in water solution due to the equilibrium present between the phosphate anion and the hydrogen phosphate anion and the dihydrogen phosphate anion which take up hydrogen from the water and therefore leave hydroxide anions. This base is available as prills for a stripping agent in painting.

4.4 Oxidizing Agents (Organic/Inorganic) 

        The case to the right shows the effect of hydrobromic acid on hydrogen peroxide.  Whereas acidic peroxide solutions are one of the possible oxidizing agents that one can pick from, using hydrobromic acid/H2O2 solutions is not advisable.  The hydrobromic acid will act as a catalyst to decompose the H2O2 resulting in lessened yields, and in addition, the oxidation potential of the mix is enough to oxidize Br- anions to elemental bromine.  This is clearly shown, initially the H2O2 and the HBr solutions were clear, when mixed they immediately turned yellow, and upon standing for a minute or so the mix was a deep red with bromine vapors clearly stagnant above it.  Just goes to show you that you need to consider even the smaller things when attempting oxidation reactions.

Common oxidizing agents

Potassium Perchlorate   KClO4	Solid, white powder, non-hygroscopic, very slightly soluble in water.
Sodium Nitrate   NaNO3	White powder soluble in water, hygroscopic, slightly saline/bitter taste (don't taste it!).  Acid solutions will attack noble metals such as copper.  Widely available during the summer months as fertilizer.
Nitric Acid  HNO3	Clear - Yellow/Green liquid.  Available in various concentrations, >70% show remarkable oxidizing capabilities.
Hydrogen Peroxide   H2O2	Clear liquid, available in various concentrations from 2% to 99% solutions greater then 50% should be treated with care as combination with many things can cause them to explode.  Greatly attacks tissue.
Potassium Dichromate  K2Cr2O7	Bright orange solid, soluble in water.  Solutions of potassium dichromate with sulfuric acid were once one of the most routine things to clean lab glass with.  Potassium dichromate is considered carcinogenic.
Sodium Hypochlorite  NaClO	Clear-Yellow/Green liquid strong chlorine type smell.  Surprisingly good widely available oxidizing agent.  Considerably more powerful in concentrations greater then 12.5% and especially when hot.
Sodium Chlorate  NaClO3	White solid available as a weed killer in some areas.  Toxic and hygroscopic it has powerful oxidizing powers as a solid, when heated on its own it undergoes self oxidation-reduction  to perchlorate and chloride.
Potassium Permanganate KMnO4	Bright purple solid possessing great oxidizing ability as a solid and in either basic or acidic solution. Found as a treatment for water in areas where iron is a problem.

Aqueous Oxidations:

Hydrogen Peroxide Solutions

        Shown above is an attempt to dissolve nickel metal under various conditions.  Although not totally apparent the HCl solution and the H2SO4 solution showed little attack.  The HCl/H2O2 solution did show some attack.  However it was the H2SO4/H2O2 solution that showed incredible results.  As you can see the entire top of the nickel in the test tube to the far right has eroded to a point.  In addition the whole bottom of the test tube is full of nickel (II) sulfate crystals.  The mix of H2O2 with H2SO4 also looks entirely different from just H2SO4 acting alone, seen in the second to left picture, a cloudy mixture formed in that instance unlike the superb green mixture formed from H2SO4 reacting in tandem with H2O2.  The reason for this is H2O2 increases the process by oxidizing the noble metal, once the surface is oxidized the oxide dissolves in the acid, and once it dissolves in the acid the H2O2 can oxidize the surface again.

4.4a Molten Salt Oxidations / Solid State Oxidations:

4.5 Dehydrating Agents/ Desiccants

        In the case of the above picture nickel chloride is shown.  The kernel on the right being an anhydrous lump, and the green solution on the left being the same amount solvated in water.  This is just one example of a hygroscopic salt that changes color when hydrated.  Mind you, by being hygroscopic, a salt is not at the same time disquecent.  A disquencent salt will pull enough water from the air to put itself into a solution, an example being NaOH or CaCl2, a salt that is hygroscopic, but not disquecent, will form a stable solid hydrate that is more easily handled.  Another example of a color changing salt that forms a stable hydrate is copper sulfate, which is colorless when anhydrous but turns blue when its water removing capacity has been used up, it can be reactivated for use by heating for an extended period of time.

Chemical Name/Formula	Class	Details on Dehydrating Action
Sulfuric Acid  H2SO4	Dehydrating Acid Liquid	Heavy liquid, dehydrating action most apparent at high concentrations 90%+, concentrations higher are possible by dissolving the acid anhydride (SO3) in concentrated H2SO4, such solutions (called oleum) possess additional dehydrating strength, but remain liquids, will dehydrate sugar to carbon.
Phosphorus Pentoxide P2O5	Acid Anhydride Solid/Powder (becoming plastic like/liquid) Dehydrating Agent	Solid/Powder formed by burning phosphorus in air.  Disquecent, pulling water from air making a crust on the surface forming differing phosphor acids, phosphinic acid, phosphoric acid, etc.  Very strong dehydrating agent, forms N2O5 from concentrated HNO3.
Magnesium Sulfate MgSO4	Drying Agent Solid	White powder, commercially the heptahydate *7H2O is available, however this can be dehydrated in an oven maxed out for a few hours.  There is no color change upon hydration or dehydration.  MgSO4 is cheap and decent for drying some gasses and liquids however its action is not very strong.  The hydrated salt is a solid.
Calcium Chloride CaCl2	Drying Agent Solid (anhydrous) Liquid (hydrated)	CaCl2 is widely available for use as deicing or as a drying agent for use in basements.  It comes in the form of solid prills that will suck moisture from the air until they turn into a puddle.  The drying action of this solid is similar to anhydrous MgSO4 but the liquid hydrated state may run back into reactions.
Copper Sulfate CuSO4	Drying Agent Solid (Colorless) Green/Blue when hydrated	Solid that changes color when its drying action is used up.  Good for dying alcohols and such, action is stronger then CaCl2 or MgSO4 and it is regenerateable over high heat.

4.6 Poisonous Reagents 

4.7 Solvents 

Common Solvents:

Specific Solvents: (Red = Flammable; Blue = Non-Flammable; Green = Burns with difficulty)

WS =Water Soluble SWS = Slightly Water Soluble

Water H2O:

Diethyl Ether CH3CH2OCH2CH3 sws:

Methylene Chloride CH2Cl2 sws:

Acetone CH3COCH3 ws:

Methanol CH3OH ws:

Ethanol CH3CH2OH ws:

Tetrachloro ethylene CCl2CCl2:

Trichloro ethylene CCl2CClH:

Toluene C6H5CH3:

Xylene C6H4(CH3)2:

4.8 Metals 

4.9 Halogens 

Fluorine

Chlorine

Bromine

Iodine

4.10 Alkali Metals 

Not only are the alkali metals interesting for their reactivates and unusual properties, they exhibit definite trends as you move down the period which are easy to memorize.

· Lithium is the hardest alkali metal, but it can still be cut with a knife, as you move down the period from lithium to cesium the metals get softer and their melting points go down as well, cesium is a liquid only slightly above room temperature.

· Lithium is the least reactive of the metals, cesium the most, francium is only present on the earth in gram amounts at any time but it likely continues this trend.

· All alkali metals exhibit the +1 oxidation state as their main and only common oxidation state.

· Atomic radius increases as you move down the period, cesium, the largest stable element in the periodic table can stabilize the triodide anion I3- due to its size.

· The increasing reactivity of the alkali metals can be seen in their oxide formation, lithium forms predominately the normal oxide Li2O, sodium mainly the peroxide Na2O2, potassium mainly the superoxide KO2, and rubidium and cesium form almost entirely the superoxide.  Additionally they form somewhat stable ozides CsO3.

· All are soluble in liquid ammonia or hydrazine yielding brilliant blue solutions that contain 'solvated electrons'.  That are powerful reducing agents.

· React with water along the lines of   M(s) + H2O(l) ---> MOH(aq) + 1/2 H2(g)  Lithium is somewhat slow, sodium will ignite and can ignite clouds of hydrogen above it leading to explosion, by the time you get to cesium it will detonate.

· Lithium is the least dense metal of the period, it will float on most any oil you try to protect it under, cesium is the most dense.

        Distinguishing between the different alkali metals in solution can be incredibly difficult chemically as they all behave very similarly and most salts are very soluble.  The easiest test to distinguish between the alkali metal cations in solution is the simple flame test.  A circle of wire, preferably inert, e.g. platinum, is dipped into a concentrated solution of the salt.  It is then put into a high flame and the color of the flame observed.

Lithium = Red
Sodium = Orange/Yellow
Potassium = Purple
Rubidium = Red - Violet
Cesium = Blue  

        However the tests can be easily false, the colors can be over-run by other colors generated, and sodium, the most common contaminate of the other alkali metal salts due to most of them being produced from it, can easily over-shadow the other more sensitive colors of potassium and the like, leading to a false positive for sodium.

Lithium

Like other elements that start off their group lithium shows some properties that differ from the normal properties of the rest of the table.  When exposed to air lithium will form a black coating of nitride Li3N which reacts with water to produce ammonia.  Only magnesium metal appropriately heated similarly forms the nitride.  Nearly all lithium salts are hydrated and removing the waters of hydration are nearly impossible on some of them, particularly the chlorate and perchlorate which usually decompose before a majority of the water has been removed.  Lithium perchlorate actually contains more oxygen, gram for gram, then liquid oxygen.  The hydroxide of lithium forms a stable hydrate LiOH*8H2O and does not dissolve in atmospheric water if left exposed to the atmosphere, but it will form the carbonate like the other alkali metals, reacting with atmospheric CO2.  The carbonate of lithium is the least stable carbonate of the alkali metals and decomposes around 1000C.

        Lithium metal has a high electrode potential and therefore it is becoming popular in some batteries (this also means that it can be very reactive in the liquid state, it will destroy glassware when liquid).  Lithium is one of only a handful of metals that have actually become more expensive over the last 20 years.  Lithium is usually stored on top of oil under an argon atmosphere.

Sodium

        This is the most common alkali metal that we run across.  Sodium carbonate, sodium chloride, sodium vapor lights, sodium hydroxide, sodium bicarbonate, the list goes on.  Sodium is the most abundant of the alkali metals in the earths crust and it shows in our preference too it, that and the fact that we need sodium chloride in our daily dietary intake.  

Potassium

         Potassium follows the trend set up between sodium and lithium in that it is more reactive then either of them, being further down the group.  Therefore when it burns in air, shown left (Notice the purple flame that is also indicative of potassium ions in a flame test), it produces not only the oxide and peroxide, but predominately the superoxide.  KO2 is a very powerful oxidizing agent and it proves to be a nuisance when storing potassium.  When a block of potassium is stored under mineral oil or another inert substance unless it is in an air tight container and the liquid has been degassed the potassium can pick up oxygen and form the superoxide.  This leaves a yellow-orange coating on the surface of the pieces of potassium.  This in and of itself is usually no problem however upon cutting into a piece of potassium covered in this coating it can force the superoxide into the unreacted potassium and in worst cases this can cause an explosion, sending flaming bits of potassium everywhere and causing severe damage to the individual performing the manipulation.

Potassium superoxide and the superoxides of other higher alkali metals react with water along the following equation:

2KO2(s) + 2H2O(l) ---> 2KOH(aq) + H2O2(aq) + O2(g)

Potassium superoxide is also used in space capsules for the duel purpose of sequestering CO2 from the astronauts breath and to generate additional oxygen.  It is also used in some self contained breathing apparatuses:

4KO2(s) + 2CO2(g) ---> 2K2CO3(s) + 3O2(g)

Sequestering an additional water molecule and CO2 molecule by the following:

K2CO3(s) + CO2(g) + H2O(g) ---> 2KHCO3(s)

        Another interesting property of potassium is that it forms a carbonyl compound K(CO)6 however it is not the most stable of carbonyl compounds by a long shot.  It can explode for no reason at all at STP therefore any reaction that generates elemental potassium, or uses elemental potassium, in the presence of carbon monoxide  (i.e. reduction of the carbonate with charcoal) should be treated with caution as the in situ preparation of potassium carbonyl may cause explosions.

        Potassium-sodium alloys are liquids at STP and more reactive then either metal individually.  Industrially potassium is prepared by distilling it from a mixture of sodium metal and potassium chloride.  The replacement of potassium with sodium in the reaction is not immediately sensible due to potassium being the more reactive, however the reaction works due to the potassium formed having a significantly lower boiling point then the sodium therefore the reaction is pushed foreword as the potassium boils off.  The Castner cell also works with potassium hydroxide in place of sodium hydroxide and actually gives better yields and a lower melting solid.  However potassium is more flammable and reactive therefore this reaction is less favored.  The electrolysis of the chloride is also less favored due to higher working temperatures of the eutectic.  Thermite type reactions also work for the production of potassium, reducing potassium oxide or hydroxide with magnesium, aluminum forms aluminates that decrease yields significantly.  The most common potassium salt available to the amateur chemist is potassium chloride, it is widely available for use in water softeners and as a salt substitute in health food areas.

Rubidium / Cesium

4.11  Gasses

        A gas bubbler, as show on the left can be quite the useful tool for an at home lab.  However the construction of such a device is exceedingly simple and therefore purchasing a piece of equipment like the one shown can be avoided.  The basic principle is simple, gasses come in through the tube on the left, bubble though a solution contained in the body of it, going though a glass frit on the way to disperse them better thereby creating more surface area causing the gasses to be absorbed/washed better, and finally exit though the tube on the right which comes nowhere near the water thereby preventing the liquid in the container from being transferred to the next container.

        The most easily constructed substitute consists of a funnel inverted to dip just below the surface of the liquid into which the gasses will be dissolved, the gasses are forced into the funnel and they thereby come into contact with a large surface area of liquid.  This design also eliminates suck back but washing gasses like this requires more design work.

Gas generation and handling setups:

Above is a very simple gas generation apparatus. However since it lacks a trap (which would prevent fluid from coming back into the flask) it would only be good for low temperature generations to prevent a dangerous suckback.

This apparatus goes two steps further, one, it incorporates a trap to prevent liquid from flowing back into the possibly hot flask that may or may not react with water, and two, it has a final gas scrub which may possess a different liquid then the origional absorbant flask to make sure to destroy harmful vapors.

Even one step further, this setup has an initial trap, then a wash to take out impurities that may hurt the reaction, followed by another trap and finally the absorbance and the scrubbing step. Also there is a sepretory funnel for the addition of another liquid to the reaction flask to keep the reaction going.

Tips for working with gasses:

· When working with flammable gasses it is necessary to heat any part of your vessel with a non-flame, non-sparking heat source. The smallest spark can trigger an explosion and special care should be taken to remove sources of ignition from your work area.

· In the case of gasses that are highly unstable (arsine, diborane, phosphine, silane, hydrazine) and can decompose exothermically when heat is applied and if it is a case where that gas must be disposed of by incineration, it may be necessary to run the gas into a container containing damp sand, up though the bottom and though a tube at the top, this will negate the possibility of a sudden explosion of gasses flashing back and detonating your whole reaction setup.

· If working with a gas of very high toxicity (most of them on the list) and inhalation occurs, do not delay, get medical help immediately, gasses are absorbed fast and a delay of a few minutes could mean your life, call a poison control, or the local emergency services.

· Make sure your reaction apparatus is air tight beforehand, and keep some duct tape around for emergency fixes (real emergencies only, if some of these start to leak you might just evacuate before it gets to the duct tape)

Dealing with exit gasses:

Specific Gasses: (Green = Water Solution Basic Red = Water Solution Acidic) WS = Water Soluble

Acetylene HCCH:

Air: Air is a mixture of gasses of approximate composition by volume; (78%) Nitrogen, (21%) Oxygen, (1%) Argon, (<1%) Other gasses, CO2, Ne, He, etc. [Exact composition of the air around you varies with your elevation and your surroundings, however these numbers are relatively constant] Unless you take steps to the contrary, you will be working in air. Most reactions can be carried out with exposure to the atmosphere, there are many that cannot be though, be sure to take into consideration the properties of the atmosphere you make for your reaction, and how it will react with your reaction before beginning any involved chemical reaction.

Ammonia (NH₃) WS:

Arsine (AsH₃) WS: Arsine has a garlic-like odor similar to phosphine. Initial exposure to arsine produces few symptoms, headache, nausea, nothing to make a person worry too severely. However several hours or so after what some would call a mild exposure, a breath or two, vomiting and cramping set in and depending on the dose kidney failure, CNS depression, and death. Arsine is the most toxic way for the body to come into contact with arsenic. Because of its severe toxicity arsine should be avoided. To destroy arsine from exit gasses it is prudent to run the gasses though at least two washes of sodium hypochlorite, calcium hypochlorite, potassium permanganate, bromine water, or sodium hypobromite solutions. Be sure their volume is sufficient to provide excessive decomposition ability for more arsine then you can think might be produced. Do not run arsine into any incineration, fine As₂O₃ will be produced creating a terrible wide spread inhalation hazard. Arsine is really terrible, it will cause long-term reproductive damage, and cancer concerns, do not tinker with it. It will decompose to its elemental constituents at around 400°C providing there is not oxygen present to support its combustion. Electrolysis of solutions that contain arsenic cations under acid conditions can produce arsine, this is not meant to be a preparation however, it is a warning.

Butane CH₃(CH₂)₂CH₃: Butane is relatively uncreative to may chemical conditions. It makes a workable inert gas for some situations but the actual production of butane on a lab scale from other chemicals is needlessly complicated as it can be purchased for refilling lighters and some camping supplies easily. Butane containing exit gasses should be run into the intake of a torch to prevent butane vapors from accumulating around your work area and causing an explosion hazard. Although butane does not prevent as specific inhalation hazard it can act as an asphyxiant gas, especially if used in an enclosed area.

Carbon Dioxide CO₂ WS: Carbon dioxide has no detectable smell. You breath out carbon dioxide and plants take it in and as such it does not possess some terrible toxicity. But in concentrated form it is hazardous for your health. Therefore when working with dry ice or a carbon dioxide cylinder or another source that could provide a large amount of carbon dioxide quickly into a small space one should take care and have adequate ventilation. Aside from commercially available dry ice (solid carbon dioxide) and cylinders that are available for the soft drink industry, it is also marketed as a water solution seltzer. Production of carbon dioxide in the home lab is very simple. The addition of a markedly acid solution to a carbonate or bicarbonate, either solvated in water or simply as a powder will generate copious volumes of carbon dioxide, the production of which is controlled by the rate of acid addition:

CO₃^2-_(aq) + 2H⁺_(aq) Þ CO_2(g) + H₂O_(l)

Carbon dioxide produced in this way will contain significant quantities of water vapor that must be removed using a desiccant and additionally if your acid has a volatile component (e.g., HCl, HNO₃) a portion of the acid may carry over as well and need its own separate scrubbing. The use of carbon dioxide as an inert atmosphere is covered in section 8.4 an example of a preparation of carbon dioxide for lab purposes would entail a setup consisting of a sepretory funnel on the gas generation flask for the addition of acid at will. The exit gasses for this vessel would be scrubbed for volatile acid components such as HCl and HNO3 by passing though a carbonate solution which will in turn make more CO2 and passing those gasses though strong H2SO4 to remove any water that may be present. A minimum of one trap is required between the volatile acid scrubbing chamber and the sulfuric acid chamber as suck back of H2SO4 into saturated carbonate would not be favorable. Additional traps could be placed though out but the controlled addition of acid to the solution should ensure that gas continuously flows away from the reaction flask. Carbon dioxide is also the product of complete combustion but the carbonate method works better for a reasonable supply. Carbon dioxide is a gas slightly heavier then air, it possesses no oxidative properties except with strong reducing agents under heat and it does not act as a reducing agent, upon dissolution in water it makes a slightly acid solution of carbonic acid.

Carbon Monoxide CO: A colorless, odorless, poisonous gas. It’s toxic action is produced by its strong bonding with the oxygen carrying constituent of human blood, hemoglobin. Carbon monoxide inhalation is treated with oxygen, however since it has no odor, and the usual warning signs of poisoning are lethargy and headache, which can easily be overlooked cases of carbon monoxide poisoning usually go untreated resulting in chronic poisoning or fatal poisoning. The most common reason for carbon monoxide poisoning is from a faulty furnace in the home.

Carbon monoxide is the product of incomplete combustion of carbon containing molecules. It’s use in chemistry is actually quite extensive, however for the beginning chemist it is really not a gas of interest. It does however act as a decent reducing agent, and in environments where carbon is the reducing agent under extreme conditions carbon monoxide can be the main product, therefore taking precautions to remove the gas is an important measure. Other then generating insanely toxic carbonyls of the transition metals though it does not have the wide variety of uses that would render it highly important. It’s preparation is simple, the addition of concentrated sulfuric acid to formic acid results in its dehydration and subsequent generation of carbon monoxide gas.

HCOOH_(aq) + H₂SO_4(l) Þ H₂SO₄*xH₂O_(aq) + CO_(g)

Gas generated in this way is suitable directly for a number of applications provided the addition of the reagents is controlled. The exit gasses of apparatuses that use carbon monoxide or produce carbon monoxide should always be lead into the entry of a flame to burn the gas away entirely. Carbonyl compounds can also be pacified in this manner. However allow me to reiterate the main fact here, carbon monoxide is highly poisonous and gives no warning to its presence, apparatuses using it should be checked and rechecked to ensure they are air tight and exit gasses are properly incinerated, the author of this work recommends against intentionally using carbon monoxide in any preparations.

Chlorine Cl₂: (See Section 4.9)

Diborane B₂H₆: Diborane is the dimmer of borane (BH₃), which has never been isolated on its own. Diborane is toxic, inhalation of diborane results in headache, dizziness, unconsciousness, fluid in the lungs, and finally death. Diborane is spontaneously flammable in moist air, which can cause it to result in explosions. Diborane burns with a green flame producing powdery B₂O₃ that lays a fine dust on all surroundings. The main use of diborane is in the preparation of sodium borohydride, a moderately strong reducing agent that can be safely recrysatalized from water.

The preparation for diborane is similar to the preparation of silane or of hydrogen sulfide. Acid is slowly dripped over solid magnesium boride (MgB₂) resulting in the formation of this spontaneously flammable gas (Note, Diborane is the main borane produced, other higher boranes are also produced).

2MgB_2(s) + 4HCl_(aq) + 3H2O_(l) Þ B₂H_6(g) + 2MgCl_2(aq) + 2H_2(g) + B₂O_3(s)

Being that diborane is pyrophoric it is easy to assume it might be easily oxidized, that is in this case very true. Atmospheric oxygen and even water will oxidize diborane. Therefore all handling vessels must be free of moisture and especially oxygen (however it would be impossible to exclude moisture from the initial flask as it is necessary for the reaction and a part of most acids). Note that diborane was once considered for use as a rocket fuel but the produced boric oxide was too abrasive on the rocket cones, to say this another way, its oxidation is very exothermic and such an exothermic reaction anywhere near you would be disastrous. The industrial process for producing diborane involves the following reaction.

BF_3(g) + 6NaH_(s) Þ B₂H_6(g) + 6NaF_(s)

Additionally the reaction between sodium borohydride and elemental iodine produces diborane. Although simpiler both of these reactions are complicated by the use of two somewhat difficult to obtain chemicals, whereas the magnesium boride used in the first reaction can be obtained quite readily by a thermite reaction:

4Mg_(s) + B₂O_3(s) Þ 3MgO_(s) + MgB_2(s)

Boric oxide and magnesium shavings being more available then the boron triflouride and sodium hydride used in the industrial process.

Being that diborane is a very toxic gas any emissions from a system that may contain diborane should be appropriately dealt with as to reduce their danger. Although the substance is pyrophoric one cannot assume that diborane leaking from a reaction vessel will oxidize before it has a chance to do any damage to their person. Therefore one should do either or both of the following; 1) Diborane will react with water to produce boric acid and hydrogen gas readily, bubbling diborane though a a tall column of water, using an efficient bubbling mechanism (e.g., a glass frit) should greatly diminish the concentration of diborane. 2) Running the exit gasses of a system into the air intake of a burner. Much care should be taken with diborane, it is toxic, and explosive, and spontaneously flammable, this is yet another gas that poses such a danger as to make the author of this work suggest against working with it.

Ethylene CH2CH2: A colorless gas with a slightly sweet odor. Once used for anesthesia it is exceedingly flammable and explosive under the right conditions. It also has a unique property of acting to ripen fruit despite it being such a simple molecule. Although it can act as an asphyxiant is does not possess any excessive toxic effects on the human body. Unlike acetylene hazardous polymerization is less likely however this gas is quite reactive toward oxidizing agents. Here are a few examples:

CH₂CH_2(g) + Br_2(aq) Þ CH₂BrCH₂Br_(l)

CH₂CH_2(g) --(KMnO₄/NaOH_(aq))--> CH₂OHCH₂OH_(aq)

That double bond that ethylene has is quite prone to addition and it is likewise reactive as illustrated above. It has some use in synthesis for such reasons but no use as an inert atmosphere. The production of ethylene is one step further then the production of diethyl ether. In the case of ether two molecules of ethanol are condensed and a water molecule is lost, ethylene goes one step further, another water molecule pulled out of diethyl ether so to speak according to the following reaction:

2CH3CH2OH(l) --(High Heat/Excess H₂SO_4(l))--> 2CH2CH2(g)

High temperatures (>140C) and a significant excess of sulfuric acid (at least a 4x molar excess) to ethanol produce favorable conditions for ethylene formation. It is recommended that you add some fine sand to the reaction mixture to make a slurry that prevents some of the bumping associated with this reaction. As with working with other highly flammable gasses heating should be accomplished by an oil bath held on an electric/non-sparking heating source. The ethylene thus produced after being put though a weak basic wash is suitable for most any purpose.

To scrub ethylene from the exit gasses from a reaction either significant bubbling though an alkaline permanaganate solution or leading the exit gasses into the intake on a flame will work. Although highly flammable, ethylene is less dangerous then some of the gasses mentioned here and can be used with a measure of safety.

Fluorine F₂ WS (reacts): (See Section 4.9)

Hydrogen H₂: Hydrogen gas is highly flammable although otherwise not terribly reactive. With heating it will form hydrides with some metals and it is a good reducing agent at high temperatures but the kind of run-away reactions that can be expected from some reactions are simply in the lacking with hydrogen (except that whole highly flammable thing). Hydrogen is a very light gas that clears the reaction area very quickly owing to its ability to take flight. The simplicity of making hydrogen gas in the laboratory is only matched by the simplicity of making carbon dioxide. Section 4.2 on acids covers the hydrogen activity series, anything in the list above hydrogen will displace hydrogen from an acid and will therefore produce hydrogen gas as the metal is solvated. If the only acids you have are weak acids you need a somewhat reactive metal, aluminum or magnesium will displace hydrogen from acetic acid (it works with the other metals too but due to dilution the reaction is slow), but usually somewhat stronger acids are available. A great reaction to produce hydrogen gas is to drip hydrochloric acid onto magnesium scrap, the reaction is fast but it is therefore over fast and allows for a more precise control over the speed of hydrogen generation, other good choices might be common steel wool, aluminum foil, or zinc in any form (Do not use nails, you do not want to make a shrapnel bomb!).

Mg_(s) + 2HCl_(aq) Þ MgCl_2(aq) + H_2(g)

The gas thus produced would have to be scrubbed for two things, volatile acid components (HCL_(g)) and for water that vaporizes, it depends on how pure you want this gas as to how much scrubbing you are going to do on it. Running the gas though a sodium hydroxide mixture will eliminate acidic components and running it though a final sulfuric acid wash will remove gas and make the hydrogen suitable for most any application. Exit gasses containing hydrogen are best disposed of by incineration.

Hydrogen Cyanide HCN WS: This is actually a borderline gas/liquid, approx. Bp is 26C so slightly above room temperature, it has the odor of bitter almonds, however only 1/3 of the population can smell it due to a genetic defect. But that is not important, what is important is that this is one of the top four most toxic gasses listed in this countdown. Not only is it toxic, but it is prone to explosive polymerization if not properly stabilized and it is also flammable/explosive on its own. Hydrogen cyanide is more often then not accidentally made when a chemist dabbling in cyanides decides to acidify the solution, big mistake:

NaCN_(aq) + H₂O_(aq) Û NaOH_(aq) + HCN_(aq)

I say big mistake because that equilibrium lies too far to the right for my own comfort to begin with owing to the weakness of hydrocyanic acid and when acid is added separately it forces the equilibrium to the right by taking up the sodium hydroxide and therefore hydrogen cyanide is formed, although very soluble in water the acid continuously comes acidified solutions, and if significantly acidified a significant eruption of HCN can occur. Hydrogen cyanide has incredible knockdown power, one whiff can instantly put a person on the floor unconscious with no hope of recovery and death following within two or three minutes. Scrubbing hydrogen cyanide from exit gasses is a simple affair as it is so easily oxidized bubbling though a sodium hypochlorite solution or other mild to strong oxidizing agent will easily oxidize the cyanide anion to cyanate OCN^- a less toxic version. Hydrogen cyanide really is a chemical that I recommend others not even work with, as such the information here is more of a warning then anything amyl nitrate (inhalation) and sodium thiosulfate (ingestion) are somewhat of folklore medicines to help combat cyanide poisoning.

Hydrogen Fluoride HF WS: Another borderline gas/liquid, approximate Bp is 19C, slightly below room temperature. As with all soluble fluorides hydrogen fluoride would be considered toxic to start out with, however it goes one step beyond, if you get a splash of hydrogen fluoride on your arm the accepted scenario involves your flesh dissolving to the bone, then your bone dissolving, followed by several minutes of pain as you get terrible fluoride poisoning, finally ending with your heart stopping. But that’s liquid hydrogen fluoride, nasty stuff anhydrous.

In dilute solutions (<3%) it is fairly safe to work with, concentrations of this magnitude are available in over the counter products. They are still toxic but not to the same degree. As a gas anhydrous hydrogen fluoride is almost just as bad, eating away at lung tissue and poisoning your body. At least it’s not flammable. Hydrogen fluoride of all concentrations will attack glass show in the following equation:

SiO_2(s) + 6HF_(aq) Þ SiF₆H_2(aq) + 2H₂O_(l)

The ability of hydrofluoric acid to attack glass increases steadily with concentration to about 99.9%, supposedly totally anhydrous hydrogen fluoride will not attack glass, however notice that water is created in the attack on glass, therefore even a stray water molecule could catalyze the reaction. It will eat glass to a significant extent eliminating the ability to use glass vessels when handling hydrogen fluoride, either as a solution or as a gas.

CaF_2(s) + H₂SO_4(l) Þ CaSO_4(s) + 2HF_(g)

The formation of hydrogen fluoride by heating calcium fluoride with concentrated sulfuric acid is favored by the volatility of the hydrogen fluoride thus formed and continued heating to keep driving it off. Additionally hydrogen fluoride could be had by removing the water from over the counter solutions, either by distillation (be sure to check for azeotropes) or by dehydration via a strong desiccant.

As an aqueous solution hydrogen fluoride is a weak acid, this is to be expected if you look at the trend setup by the other hydrogen halides, but it is actually much weaker then could be predicted. This is because the bond between hydrogen and fluorine is incredibly strong and therefore ionization happens to a considerably lessened extent. It will dissolve many metals though and as it is doing so the heat of the reaction is probably vaporizing HF out of the solution and into your air.

Please, hydrogen fluoride is not only a terrible contact poison but cumulative poison, please consider this information a warning as to its dangers.

Hydrogen Halides except Fluoride HCl, HBr, HI WS: These are all noxious smelling gasses at room temperature and pressure. Their solution in water form the recognized acids hydrochloric acid, hydrobromic acid, and hydroiodic acid. Inhalation of these gasses can cause damage to the lungs and mucus membranes. Skin contact with concentrated vapors will result in discoloration and possibly necrosis. Hydrogen bromide presents a problem different from the other two in that bromine is not normally utilized by the body, hydrogen bromide inhalation, resulting in the increase of bromide in the body results in increase lethargy and in extreme instances, death. Although not exceedingly toxic these gasses all cause damage on the physical level, which if enough could cause death.

All of these gasses are soluble in water to significant extents, hydrogen chloride the least soluble and hydrogen iodide the most. Their acid solutions provide the most effective ways to generate the gasses. Dehydration of the solutions, expecially if they are initially concentrated will result in the formation of the free halogen halide.

Above is an apparatus ideal for the formation of a hydrogen halide gas from a hydrogen halic acid. In such an apparatus the acid to be dehydrated is put into 1 which is a funnel closed off from the rest of the system by the stopcock 2. This funnel extends downward with a stem that has a very small opening, almost like capillary tubing 3. This extends down to the bottom of the vessel and deep under the sulfuric acid 4. The gas upon formation bubbles though the sulfuric acid and out of the gas tube 5. The opening in the tubing is small enough that it draws down more acid until the stopcock is closed. This apparatus will generate large quantities of water-free hydrogen halide. The simplified version is just a container full of sulfuric acid into which your hydrohalic acid is dripped in with a sepretory funnel. The disadvantage of this version is the water spray and evaporation which calls for an additional wash of the exit gasses.

The hydrogen halides can also be generated from a chemical reaction. Hydrogen chloride can be generated (although not controllably) by the reaction between sodium chloride and concentrated sulfuric acid:

NaCl_(s) + H₂SO_4(l) Þ NaHSO_4(s) + HCl_(g)

Although for hydrogen bromide and hydrogen iodide this reaction is not as feasible. Some of the hydrogen bromide formed will be oxidized by the sulfuric acid to free bromine and for the reaction between sodium iodide and sulfuric acid most if not all of the hydrogen iodide produced is oxidized to iodine. However distillation of weaker solutions of sulfuric acid with these salts can result in aqueous azeotropes of the acids. The hydrogen halides of these salts can be generated by the reaction of a sodium salt with concentrated phosphoric acid, which lacks the oxidizing power to release the free elements.

HBr and HCl can also be generated as side products from organic halogenations. Usually ½ of the initial halogen is consumed in the reaction and the other half is released as the hydrogen halide, iodine however does not posses the power necessary to perform most organic halogenations and therefore this is not a good way to make hydrogen iodide.

Exit gasses containing these acidic gasses can easily be scrubbed by bubbling them though concentrated sodium hydroxide solution. Subjecting these gasses to high heat will dissociate them to some extent and oxidizes some of them part way but results in a multitude of products.

Hydrogen Sulfide H₂S WS: Yet another incredibly toxic gas, this one more so then hydrogen cyanide. It has it’s own unique smell, which everyone has probably smelled at one time or another, the smell of rotten eggs. However, although the smell of hydrogen sulfide becomes prominate at levels below lethal levels it has one major trick, it deadens the sense of smell quickly so the smell goes away and you think you’re okay, but actually you might just be about ready to die. Many beginning chemists have died from hydrogen sulfide thinking that because they could not smell it they were okay. Smell is only a way of detecting hydrogen sulfide if it is generated unexpectedly, if that happens leave the area. You do not want to mess around with this chemical.

The treatment of many sulfides with acid is the cause of hydrogen sulfide production usually. As with hydrogen cyanide, hydrogen sulfide dissolves in water forming a weak acid solution, and adding acid to a salt of hydrogen sulfide drives the equilibrium to the production of hydrogen sulfide gas. Sulfides can easily be made by the direct combination of an active metal with elemental sulfur followed by heat. They can also be the product of high temperature reductions of sulfates.

CaSO4_(s) + H2_(g) Þ CaS_(s) + H2O_(g)

S^2-_(aq) + 2H₂O_(l) Û HS^-_(aq) + OH^-_(aq) +H₂O_(l) Û H₂S_(aq) + 2OH^-_(aq)

In the first equation calcium sulfate is treated with hydrogen at a high temperature in the absence of oxygen, the products are gaseous water and calcium sulfide. The second equation shows the equilibrium that exists in a neutral solution of hydrogen a sulfide salt of of hydrogen sulfide. Addition of a base adds hydroxide, which appears on the far right, this drives the equation to the left, and as long as plenty of base is present sulfides are relatively safe. However if acid is added that will destroy the base and it will protonate the sulfide anion floating around it the solution, both of these will drive the equation to the right and produce hydrogen sulfide gas.

The allure of hydrogen sulfide is there though. It is useful in the laboratory, as an agent to detect certain metal compounds, as a reducing agent and also to generate some interesting acids.

Br_2(l) + H₂S_(g) Þ 2HBr_(aq) + S_(s)

If this reaction were to be carried out under a layer of water and stirring were applied it would not stop there, the sulfur formed would react with the bromine formatting S₂Br₂ which would react with the water resulting in the oxidation of sulfur converting it to SO₂ and making two more molecules of hydrogen bromide, most of which would dissolve in the upper layer of water making a concentrated hydrobromic acid solution. Solutions of formic, acetic, and other acids can be made in this manner. Exit gasses containing H₂S can be burned or scrubbed with two washes of concentrated basic solutions or alkali oxidizing agents such as KMnO₄. But the danger of carrying out such operations usually render this unfeasible, hydrogen sulfide kills, the author of this work recommends against not working with this toxic chemical.

Methane CH₄: A common gas with which most people are familiar, methane is a simple asphyxiant gas with the side pitfall of being an explosion hazard. It has no uncommon reactivities and behaves fairly inertly. Natural gas piped to homes is mostly methane with other agents added for smell and a small percentage of other hydrocarbons. Methane in home synthesis is somewhat of an extreme measure do to its lack of reactivity. Chlorinating methane should yield carbon tetrachloride, chloroform, methylene chloride, and methyl chloride. However the careful control of temperature and necessary supply of chlorine gas are usually outside of the normal working scope of the home lab. It is interesting to note the simple procedure by which methane can be generated in the home lab:

CH₃COONa_(s) + NaOH_(s) Þ CH_4(g) + Na₂CO_3(s)

By simply heating anhydrous sodium acetate with sodium hydroxide the reaction commences generating methane gas and sodium carbonate. Methane can be formed as a product of the electrolysis of some mixtures containing organic components or by the action of water or acid on aluminum or beryllium carbide. If a reaction is run in which methane provides an ‘inert’ atmosphere or a reaction that involves the production or use of methane is run, the exit gasses should be lead into the intake of a burner and incinerated to prevent and explosion hazard.

Nitric Oxide NO₂ WS (reacts): Also known as nitrogen dioxide this compound reacts with water to form nitric acid and nitrogen monoxide gas. Due to this fact it is incredibly toxic, think about it, it goes into your lungs and makes nitric acid which in turn causes your lungs to secrete fluid which causes pulmonary edema which means you’re going to die. Nitric oxide has a biting odor caused by its hydrolysis upon contact with fluids within your nose, also due to its hydrolysis you can taste it. It has a high boiling point and can be easily condensed at home, as a liquid/solid it is colorless in theory due to the formation of the dimmer N2O4, however it is actually more often then not brown-red as is the gas.

4HNO_3(aq) + Cu_(s) Þ 2NO_2(g) + Cu(NO₃)_2(aq) + 2H₂O_(l)

The easiest production of nitric oxide involves the action of concentrated nitric acid upon copper metal. However if the reason you seek this toxic gas is for the production of nitric acid then that really does not seem like a feasible method of production. Alternative methods of production include the moderate temperature decomposition of heavy metal nitrates (although complicated by the formation of oxygen), or by the action of a strong electrical discharge on a mixture of oxygen and nitrogen to produce a mixture of nitrogen oxides followed by a means of separation such as condensation.

6NO_2(g) + 3H₂O_(l) Û 3HNO_3(aq) + 3HNO_2(aq) Û 4HNO_3(aq) + 2NO_(g) + H₂O_(l)

As you can see nitrogen dioxide disproportionates in water to form nitrous acid and nitric acid, however the nitrous acid thus formed is unstable and decomposes resulting in the formation of yet another molecule of nitric acid and a two molecules of nitrogen monoxide. Giving the overall equation showing that six molecules of nitrogen dioxide will react with two molecules of water to form four molecules of nitric acid. It should be noted that nitrogen monoxide is easily oxidized to nitric oxide by the action of atmospheric oxygen with no other stimulus. Solutions of high concentration nitric acid with excess NO₂ dissolved within are known as red fuming nitric acid.

Exit gasses containing nitrogen dioxide should be appropriately scrubbed using a strong sodium hydroxide solution. The formed sodium nitrate and sodium nitrite may be recovered for future uses by evaporation of the scrubbing solution afterwards. Nitrogen dioxide is a very poisonous gas, if it is to be used safety measures should be planned out in advance and it should be used entirely in a closed reaction system, the operator of such a system should wear a respirator of some sort and exit gasses should be double washed in concentrated sodium hydroxide solutions, this is not a gas to take lightly.

Nitrogen N₂: Approximately 78 % of the air you breath is, by volume, nitrogen gas. Nitrogen is a diatomic molecule that has a very strong triple bond connecting the two nitrogen atoms. It is inert to a wide variety of common applications except electrostatic discharges. It is purified from air by liquefaction of air followed by fractional distillation. It is a widely available gas in the chemistry industry and can be found for welding applications. Laboratory preparation of nitrogen is very complicated on average, one method being by the action of a strong oxidant on an aqueous ammonia solution, or by the careful heating of ammonium nitrite, or by the careful decomposition of an azide. Other methods also exist but the best is going to be simply removing other components from air. Running normal air into a metal tube full of copper wool and heated externally very hot with a torch will remove the oxygen present if the tube is significantly long enough. You are left with nearly 98% nitrogen content with 1% or so of argon. Nitrogen forms a good inert atmosphere for many reactions. It is an asphyxiant gas like any other, however there is no need to take extra precautions with exit gasses that may contain nitrogen only worry about other things that may be there.

Nitrogen Monoxide NO:

Nitrous Oxide N₂O:

Nobel Gasses; Helium He/Neon Ne/Argon Ar/Krypton Kr:

Oxygen O₂:

Ozone O₃:

Phosphine PH₃:

Propane CH₃CH₂CH₃:

Radon Rn:

Silane SiH₄:

Sulfur Dioxide SO₂ WS:

Xenon Xe:

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5.0 Simple Reactions 

5.1 Shifting the equilibrium, a basic reaction technique, and an essential one : Le Chatelier's Principle

A very large percentage, if not all the reactions that you will run across involve some sort of equilibrium. Since I stated that nearly every reaction involves an equilibrium many of the reactions that are considered productive have the equilibrium lie to one side of the arrow though some mechanism. Here are a few examples, which will help to explain this concept:

1. NaCl_(s) + H₂O_(l) Û NaOH_(aq) + HCl_(aq)

2. NH₄OH_(aq) Û NH_3(g) + H₂O_(l)

3. Ca(NO₃)_2(aq) + H₂SO_4(l) Û CaSO_4(s) + 2HNO_3(aq)

4. Na_(s) + KCl_(s) Û NaCl_(s) + K_(g)

The examples above illustrate different reactions, which can occur and are shifted though some process, some of which can be shifted back afterwards. Let me explain the reasoning behind each reaction now.

1. Sodium chloride, common table salt when dissolved in water undergoes an equilibrium reaction, although a vast majority of the sodium chloride dissolves forming sodium cations and chloride anions, a small percentage reacts with the water forming sodium hydroxide and hydrochloric acid. The reason for this actually relates to the individual components NaOH and HCl.

NaOH(s) –H2Oà Na+(aq) + OH-(aq)

HCl(g) –H2Oà H+(aq) + Cl-(aq)

5.2 Basic Chemistry Experiments [With Step By Step Explanations]

5.2a  Nickel salt à Nickel Oxylate à Nickel Powder

            In this very first reaction we are going to utilize a number of skills that will prove to be useful to you for quite some time and also give a good idea as to how certain reactions run.  Here is the basic outline:

· Calculate the quantity of salt to react and prepare the solutions.

· Mix the solutions and allow precipitate to 'age' or alternatively filter.

· Dry the precipitate.

· Pyrolyze the precipitate at moderate temperature without access to oxygen.

        Now let's get started!

Warning!

Nickel salts are toxic and possibly carcinogenic, please wear appropriate protective gear including gloves and goggles especially.  A respirator should not be required for this experiment but during the pyrolysis of the formed oxylate carbon monoxide gas and nickel carbonyl may be produced, it is therefore advised to do this part of the experiment with ambient ventilation.

        The first reaction in this procedure will be the preparation of the solutions.  What is necessary in this case is a soluble nickel salt, in this case nickel (II) chloride, and the somewhat insoluble acid, oxylic acid.  The reaction we will be utilizing is:

NiCl_{2 (aq)} + HOCOOCOH _(aq) ----------> NiOCOOCO _(s) + 2HCl _(aq)
Nickel Chloride + Oxylic Acid ----> Nickel Oxylate + Hydrochloric Acid

        In the picture above the beaker on the left is the oxylic acid solution, which had to be heated to get a good concentration and on the right is the nickel chloride solution, the empty beaker in the middle is where the two will be mixed.  Our aim for this reaction will be 15 g of nickel oxylate and our calculations will be based on that.  Below are some steps we can take to get our solution concentrations.  

        This is an example of a process used in chemistry quite a bit, stoichiometry.  In this process you take your known values and calculate the concentration of other chemicals necessary to make them.  It reads like a math problem, the units from one area usually cancel in the next till you get to the end, which are the desired units.  How this would read would be 1) The amount of oxylate salt you want to end up with, in this case I stated it as being 15 g.  2) The grams of oxylate salt that comprise 1 mol.  In this case your salt Ni(OCO)2 weighs 146.7 grams for every mol, remember to look to your periodic table for this information, the sum of nickels atomic weight, the four oxygens, and two carbons add up to about 146.7, check for yourself.  3) Because the bottom part of this equation was the number of grams in 1 mol, 1 mol appears at the top, therefore this one section of the equation counts for the grams of Ni(OCO)2 in one mol.  4 & 5)  This section accounts for the mols of oxylate salt that are produced from each mol of NiCl2, the oxylate salt appears at the bottom to cancel the 1 mol of oxylate salt in step 3.  6)  This step carries over the mol NiCl2 idea from 5 and allows it to be canceled, this is balanced by the weight of 1 mol of NiCl2 step 7 above it, this leaves the grams of NiCl2 necessary for this reaction as the final uncanceled answer.

        Here is the same equation with numbers.  The calculations are preformed in normal mathematics fashion.  Multiply each number in the numerator.  Multiply each number in the denominator, and divide the numerator by the denominator.  You end up with 13.26 g of NiCl2 necessary for every 15 g of Ni(OCO)2 produced.  

        Notice how each of the units cancel out except the desired units?  That is the way this type of equation should work out.  However this is not the end of this calculation.  Normally nickel chloride forms a hydrated salt.  NiCl2*2H2O, therefore the weight of one mol is different.  To correct for this take the weight of NiCl2 necessary for this reaction, 13.3 g and divide it by the atomic weight of one mol NiCl2, this cancels the grams and leaves you with the mols, now you can multiply the 0.1025 Mol by the atomic weight of the hydrated salt, simply add the weight of 2 H2O molecules to the atomic weight of NiCl2.  The weight of the hydrated salt is 163.7 g/mol and multiplied by the mols of NiCl2*2H2O necessary you get 16.8 grams of NiCl2*2H2O are necessary to produce 15 g of Ni(OCO)2.

After calculation your first solution should be made up of approximately 17 g of NiCl2*2H2O (green salt) dissolved in a minimum amount of distilled water.  Alternatively 13.5 g of the anhydrous NiCl2 (yellow-gold salt) could also be used.

         Now it is necessary to calculate the quantity of HOCOOCOH (oxylic acid) necessary to complete the reaction.  You can calculate this off either figure, either the 15 g of nickel oxylate you wish to produce or the 13.5 g of anhydrous nickel chloride that you plan to react.  Although the oxylate measurement is better as it does not suffer from 'calculation fatigue'.  Therefore to easily do this divide your 15 g by the grams of nickel oxylate per mol and multiply by the atomic weight of oxylic acid in grams.

15 g Nickel Oxylate / 146.7 g Nickel Oxylate Per Mol = 0.1022 Mol Nickel Oxylate
0.1022 Mol x 90 g / Mol HOCOOCOH = 9.198 g HOCOOCOH

        However just as before oxylic acid comes usually as the dihydrate HOCOOCOH*2H2O therefore the calculation must be altered as before.  Taking the mols of oxylate 0.1022 and multiplying by the adjusted atomic weight of oxylic acid, 126 g / mol now gives a new value of 12.8 rounded to 13 grams of oxylic acid are necessary for complete conversion of nickel chloride to the formate.

The second solution is made by dissolving 13 g of the dihydrate of oxylic acid in a minimum quantity of warm H2O.  The warm solution will help to better solvate the somewhat weakly soluble salt.

         With your two solutions now made they are simply mixed together.  The actual mixing process should agitate them enough for good combination, however if you should feel the need to stir the solution feel free to do so.  Immediately after combining the solutions there is no noticeable change, however after a few short minutes the solution becomes cloudy.

        And after even more time a precipitate, seen easily on the left picture, will settle out.  The best advice here is to let this mixture sit some time.  To 'age' the precipitate, about 20 minutes will suffice if you're in a hurry though.  Now though you have several methods to get your precipitate.  Filtration is usually an option, although as mentioned before some precipitates can pass though filters, if this is the case the precipitate must age further.  Evaporating this solution to dryness is an option as well as the other byproduct from the reaction is HCl, which will evaporate from solution, however in light of its danger to the chemist this route should be avoided.  Another method that will work after some aging is to simply decant off the top layre.  A method used in this procedure.

        Notice in the left-hand picture that the precipitate appears blue now that the supernate has been discarded.  All but about 8 ml of visible water was able to be poured off before the precipitate started to come over as well, therefore the remaining solution and precipitate was heated on a hot plate to dryness with occasional stirring to spread out the lumpy pieces.  If the solution were filtered instead the filter cake could be washed with cold ethanol and heated gently on a hot plate to drive off water.  Notice in this variation though that in the flask at left there is nearly 25 ml of precipitate, in the middle picture there appears to be much less as it evaporates, and in the final picture, although it is not clear, the little brown bottle, that only holds 10 ml is about half full, thus the precipitate was considerably unsettled and contained a massive amount of water.

5.3 Basic Chemistry reactions [Simplified Explanations]

5.3a Preparation of Sodium Acetate 

CH₃COOH_(aq) + NaOH_(aq) Þ CH₃COONa_(aq) + H₂O_(l)

One hundred milliliters, over the counter 5% acetic acid (.8 M vinegar) is placed into a beaker. After this 3.2 g of solid sodium hydroxide in prill form is weighed out and put directly into the beaker with the acetic acid [Note: 6.7 g of sodium bicarbonate may be substituted for the sodium hydroxide] , the mixture is stirred occasionally until all the sodium hydroxide has dissolved and then the mixture is placed on a hotplate and heating is commenced on a medium setting. The water is boiled off at a medium pace until the solution is less then 10 ml in volume and the heat turned down even further, the solution is observed carefully until it appears there is no further change in volume, it is removed from heat and allowed to cool. Shortly after removal of heat the solution should begin to solidify, the solid is allowed to stand until it comes to room temperature and then it is removed with the aid of scraping and placed into a storage container, yield is roughly 9 g or 100% of the trihydrate CH₃COONa*3H₂O and is somewhat impure, it can be purified by recrystalization from a minimum amount of hot ethanol or methanol.

5.3b Preparation of Zinc Hyposulfite

SO_2(g) + Zn_(s) --(H2O)--> ZnS₂O_4(aq)

Zinc hyposulfite is an interesting and powerful reducing agent who’s preparation is somewhat simple. Begin by putting together one of the setups listed in section 4.11 for the collection of a soluble gas and fill the gas collection area with 100 ml of distilled water to dissolve the gas. As seen above the gas to be solvated is SO2, please look to the section on the preparation of SO₂ to setup this part of the apparatus. Please pressurize the apparatus initially with CO2 by the addtion of acid to bicarbonate or a similar procedure to make sure the apparatus is air tight except the exit lines. Sulfur dioxide although not lethaly toxic has a strong physical destructive potential to the eyes, nose, and lungs. Please be careful in the generation of gas and the scrubbing of exit vapors to protect yourself. Finally once you are ready to begin and have your 100 ml of water ready to collect your gas slowly start SO2 evolution. The gas dissolves in the water somewhat rapidly, keeping your solution cool externally helps and at 0C your water should dissolve nearly 23 g of SO2 when it is saturated making a solution of sulfurous acid H2SO3.

Upon saturation allow your vessel to stand before taking it apart. To avoid the pain of SO2 gas you can have a large bucket of water prepared, and providing your glassware is cool your whole apparatus, minus your SO2 collection water can be slowly lowered into the water and taken apart under water. Finally after you are able to pay your sulfurous acid solution the attention it deserves put it into an ice bath to keep it cool for the next step. Take roughly 25 ml of your cold solution and add it to a small flask to which you have a lid. Add 8 grams of zinc in the form of powder and shake the solution stoppered until the solution obtains a uniform clear color again. Quickly filter this mixture and place the filtrate in a darkened bottle to prevent decomposition, repeat this in 25 ml portions for the remaining volumes. You are now left with a solution of usable zinc hyposulfite, use within a few weeks for the best results.

5.4 Dissolution of Copper Metal, e.g. oxidizing acids 

5.4a Copper metal placed in HCl/H2O2 mixture 

5.4b Heating to get CuCl2 

5.4c Dissolving in water again and precipitating with NaOH 

5.5 Amorphic behavior of Aluminum (dissolve in HCl, dissolve in NaOH)

5.6 Useful Techniques (e.g. dissolving noble metals, different techniques.)

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6.0 Practical Concerns for running an amateur lab

Although it is not a nice issue to bring up, it remains true that in most places around the world, especially industrially advanced, chemistry at home is frowned upon. For example, in a place like the United States chemistry was a widely practiced hobby until the 70’s where environmental concerns and safety considerations made chemistry seem forbidden and dangerous. Saying that you are performing a chemistry experiment may shock people in your area and might force them to call the authorities.

The legality of the procedures that you perform usually will fall into questionable territory. Outside of using common over the counter reagents for their intended over the counter purpose you are walking a thin line. The use of sulfuric acid in the form of drain cleaner to act as a catalyst might not seem like a breach of the law but some chemicals will flatly state on the back of them that it is illegal to use that chemical for anything other then the instructions listed on the opposite side. Therefore it may be prudent to move all obtained chemicals to new containers, and if possible between the move, purify them.

Aside form these legality issues of chemical possession and use, you come to the disposal issue. You will not be able to recover every bit of product from every reaction, you won’t be able to continuously run any series of reaction without generating waste along the line somewhere. You will have to dispose of this waste you create, and although the disposal is completely up to you, you should always dispose of chemicals in the most environmentally way possible, disposing of certain chemicals, by dumping on the ground, flushing down a toilet, or throwing in the trash is a major crime that can bring about jail time, and or extraneous fines.

Also the illegal dumping of chemicals can cause immediate destruction to your local ecosystem. Your grass and trees may die, accidental releases are also a problem. The unintended release of large quantities of noxious gasses can also kill grass and make your neighbors life miserable. Increasing their likelihood on calling the authorities. Although you shouldn’t have to sneak around in the middle of the night, which would make others more suspicious you shouldn’t perform your reactions in the public eye. Doing so just increases that chance that someone will see you and object to what you are doing. And if that happens disastrous consequences can result, people with small children are often the most objective over chemicals, but, it is a reasonable response on their part.

Considering all of these aspects it is usually a good thing to not flash around the fact that you have a chemistry hobby. Your chemicals should be kept under lock and key just in case anyone wants to get to them just for the fun of mixing something up. This whole situation is a sad one in some respects but it is the way that you should follow your hobby, carefully disposing of waste, not willingly divulging more information then you have to, and working out of the public eye.

6.1 Legality

6.2 Storage of chemicals, Separation of reagents

6.3 Disposal of waste materials generated

6.4 Considering your neighbors/neighborhood in every reaction

6.5 Privacy & Security

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7.0 Choosing your own experiments 

        A good way to figure out what experiments you want to do on your own is to keep track of what experiments you see.  Whenever you're reading a book or surfing the internet and you come across an interesting reaction, one where you have all the reagents or they are readily available, an experiment where you could make a product that could be useful in creating another product down the road, or maybe just something that changes color in just the right way to catch your fancy, copy it down into a book.  Be sure to write down whatever important reaction details it gives (temperature, pressure, stirring, etc.) and a complete chemical reaction if it is available.  Also be sure to include some bibliographical information as to where you found this source so if you ever have to cite it or go back to it you won't have any difficulties.        

7.1 Researching 

Devising a way to reach B-Trinitroborazine

          I don't know how I came to the conclusion that I could make B-trinitroborazine but once it got in my mind it became one of my goals.  Partly this was because I could find no papers suggesting that this compound had been made, although I did find papers on B-trizaoborazine.  A starting point was necessary so off to the library I went.  I figured the parent compound borazine would make a good starting point so I searched for the most feasible way to make that in my at home lab.  The consensus after quite a long while searching was that B-trichloroborazine was the easiest of the borazines to make, from there the compound could be reduced to borazine with sodium borohydride.  So now I went back though the literature again to find a method of preparation of B-trichloroborazine.

          The highest yielding procedure involved the reaction between BCl3 (boron trichloride) and NH4Cl (ammonium chloride).  Tracing the footnotes I found I ended up at an old CAS copy of the original article.  Original articles tend to detail the techniques used better and take little for granted so this was a gold mine of information.  But still this wasn't a good enough starting point, NH4Cl could be easily procured but BCl3 could not.  So another step backwards was necessary.  How to make boron trichloride?  The answer was simple, react boron and chlorine.  But where to get the boron, another step backwards.  After some searching for techniques of making elemental boron on the internet I ran across a thermite type reaction involving B2O3 (boric oxide) and magnesium metal.  The resulting mass boiled in HCl (aq) to dissolve magnesium compounds and the boron thus filtered off.

          So now I had a starting point from B2O3.  Still another step backwards was necessary, to the commonly available B(OH)3 (Boric acid) which is sold for ant control and such.  This can be easily dehydrated at heat to yield boric oxide.  Now I had a possible pathway to my first starting point.

2B(OH)3(s) ---[Heat]-->  B2O3(s) + 3H2O(g)
B2O3(s) + Mg(turn) ---[Heat (Violent Reaction)]---> B(s) + MgO(s)
B/MgO(s) ---[HCl(aq)]---> MgCl2(aq) + H2O(l) + B(s)
2B(s) + 3Cl2(g) --[Heat]--> 2BCl3(g)
3BCl3(g) + 3NH4Cl(s) --[Heat]--> (-BClNH-)3(s) + 9HCl(g)

          Going though my notes I went back to my sources and got as much information specific to each reaction as possible so that I could feasibly do each reaction.  For example, the second reaction also produces MgB (magnesium boride) which when reacted with HCl produces B2H6 (diborane) which is a spontaneously flammable gas at STP.  Therefore extra precautions were devised for that step.  Step 4 called for a simultaneous distillation as Cl2 gas was passed over the reactant and the last step took place in a reaction tube with the NH4Cl coated onto glass wool and the hot BCl3 gas passed though it, the B-trichloroborazine condensed in the cooler part of the tube.

          All that just to get to the one of the starting materials for a larger reaction.  Remember the goal was to devise a way to B-trinitroborazine, something not accomplished as of yet.  The procedures written out for the production of B-triazoborazine shed some light on how the trinitro derivative might be produced.  One reaction called for reacting sodium azide with boron trichloride in an inert solvent (borazine compound hydrolyze redily in water so it is not an option).  The slightly soluble sodium chloride precipitated out and the triazo compound stayed in solution.  So treating with sodium nitrate might prove feasible if I could find a proper solvent.  Another reaction called for the reaction of borontrichloride with N3(CH3)3Si (trimethylsilyl azide), a surprisingly stable azide molecule.  The process being driven by heating the mixture, highly volatile Cl(CH3)3Si (Chlorotrimethylsilane) being distilled off leaving the product.  A similar procedure might be possible with the nitro derivative so that is where my research centered next.

          Preparation of trimethylsilyl nitrate was my next objective.  The parent compound trimethylsilyl chloride was incredibly difficult to make, involving passing hot CH3Cl thought a heated tube consisting of a silicon/copper alloy.  Therefore I opted to buy the hazardous chemical.  From there I found a preparation to make the nitro derivative from the organic chemistry periodicals but upon reading it my thoughts on a reaction pathway changed and I thought up another possible reaction.  The preparation of trimethylsilyl nitrate involved adding a solution of trimethylsilyl chloride in CH3CN (acetonitrile) to a mixture of AgNO3 (silver nitrate) in CH3CN, the insoluble silver chloride precipitating out.  Acetonitrile did not seem like a solvent that would react with B-trichloroborazine and judging from its polarity it should be soluble in it.  Already I know that silver nitrate is soluble in it, and if I mix solutions of the two the insoluble silver chloride should precipitate out leaving me with a solution of B-trinitroborazine in CH3CN.  From there the volatile acetonitrile could be evaporated under exceedingly low heat and my product recovered.

          Finally I had a pathway to my molecule of choice.  Now I just need the funding.

7.1a Internet 

7.1b Library 

7.1b1 Following up on footnotes 

7.1c CAS 

7.2  Scaling up and Scaling Down

7.3 Being through 

7.4 The Importance of Keeping a Log

7.5 Trouble Shooting

7.6 Words of encouragement, stories of “Try Try Again” 

7.7 Lengthy Story about procedure with details 

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8.0 Advanced Techniques 

8.1 Working at high temperatures 

8.1a Advanced Heating techniques 

8.1b Fuel 

8.1c Refactories 

8.1d Furnaces 

8.2 Fractional Distillation 

Distilling hydrobromic acid (HBr) seemed to be simple enough so the amateur chemist thought that they would give it a shot.  After mixing together the reagents in the prescribed manner the mix was put into a 1 L flat bottom flask connected to a still head, then to a lebig condenser and in turn to a vacuum adaptor and then a 500 ml round bottom receiving flask.  A quick look in a book revealed that the Bp of the azeotrope would be 122.5 C.  Heating was begun with magnetic stirring to keep the mix agitated.  Some time later the mix began to boil.  The first of the vapor to touch the thermometer bulb at the top of the still head only brought the temperature up in the mid 90's.  The liquid condensed and dripped into the receiving flask and over the course of the next hour the temperature continued to climb, when it reached 118 C the chemist removed their current receiving  flask and replaced it with a new once, they then set aside the forerun for future analysis and watched as the temperature continued to climb to 122 C where it held steady for over an hour, during this time over 100 ml of distillate came over.  Shortly thereafter though the temperature began to drop as vapors no longer reached the bulb, signaling that the HBr azeotrope had finally finished distilling over and distillation could be discontinued.  

8.3 Catalyst Tubes (e.g. H2SO4 production) 

8.4 Inert Atmospheres

There are some reactions out there where you just don't want the air coming into contact with your product or your reactants.  The solution for this is the removal of the components you don't want mixing with your solution.  That is where inert gasses come into play.  By being inert that does not necessarily mean that the gas itself it inert under nearly all conditions.  Just that it is not going to precipitate in the reaction going on.  Two extreme examples of this would be argon, which is for all intents and purposes, totally inert, and propane, which is not considered inert to most due to its high flammability, however hydrocarbons are fairly inert, just not with respect to oxygen and an ignition source, propane therefore provides a cheap and readily available inert gas, however it does have its problems.

        When working with an inert gas the standard procedure is to run it through a closed system for a length of time to flush out any previous gasses, then introduce your reagents quickly, preferably under the inert gas to prevent recontamination of the environment.  Then sealing the system back off except for an exit and slowly letting the gas continue running though it.  There is the concern of what to do with the exit gasses, which will depend on your reagents and your 'inert' gas but that is a specific problem.  Although less useful inert gasses can also be used outside of a closed environment.  All inert gases are by nature asphyxiants and therefore the gasses should not be vented into an enclosed area. They can aid evaporation if slowly run over a hot solution and they can also be used just being sprayed into a beaker to provide some protection, argon which is heavier then air is especially suited for this.

· Argon  :  Argon is a good blanketing gas approximately 38% heavier then air and will sink into the nooks and crannies of a distillation apparatus.  It is not known to combine with anything at STP, although it does form HArF when irradiated with hydrogen fluoride at -255C but it decomposes about -245C so you shouldn't have to worry about extraneous reactions.  Couple that with the fact that it is widely available for welding applications (although there is a bit of a startup fee including a cylinder, which you can rent, and a regulator) argon is well suited for most applications.  As an exit gas you only have to worry about any contaminates in it, argon itself can be vented to the environment with no ill effects.

· Butane / Propane  :  Butane, commonly available as a refill gas for lighters, and propane, commonly available for just about everything are both gasses at STP, both are easily liquefied under pressure and vapors from either of these gasses are heavier then the surrounding air.  As such they collect in low spots, unlike argon both of these gasses are highly flammable in the presence of oxygen, as such systems using them should be thoroughly flushed first and should not contain oxidizing agents or compounds that can yield free oxygen/halogens.  In addition care must be taken with exit gasses containing these flammable products, they should be lead directly into a burner of some sort where they can be burned without hassle.

· Carbon Dioxide  :  Available in cylinders for the carbonation of beverages or as 'dry ice' relatively pure carbon dioxide has its sources outside of the lab setting.  This gas is also denser then air, however it is more reactive then the others.  It will ruin Grignard reactions, react with hydroxides and strong bases of all sorts and more, but it does have its uses, just be extra wary of reactivity.  CO2, although not incredibly toxi,c can cause damage if inhaled in a concentrated form, but as long as it is used in vented conditions it is relatively safe.

· Freon  :  Freon is a tempting source for an inert atmosphere.  First because it is widely available and second because the regulating equipment is sold right along side it.  But that is where the benefits end, freon is flammable, and somewhat more reactive then these other inert gasses, it has oils in it and as for purity, well, it’s a grab bag. Use it at your own risk.

· Helium  :  Helium is the most widely available of the noble gasses and can be picked up from many places for the purpose of filing balloons, it even comes with its own cylinder.  However the purity of such forms is questionable, unless it is in the large industrial containers it is usually mixed with a certain percentage of air.  Beyond this helium is an extremely light gas that will not blanket a majority of your vessel, however extended flushing can help to overcome this, the price of commercial helium does not make it economically feasible in light of its light nature.

· Nitrogen  :  Has a limited availability as a welding gas.  Also available in the liquid form, but be wary of frost bite.  Nitrogen is the staple inert gas for the organic chemist, usually on tap in fume hoods.  It is fairly non-reactive, cheap and easy to transport.  Nitrogen is roughly equal to air in terms of density due to it making up about 80% of our atmosphere.  There is no need to worry about venting nitrogen gas.

· Neon / Krypton / Xenon  :  Availability of these nobel gasses is considerably more limited.  They posses properties nearly identical to argon for reference.  The exception to this being Xenon, which is the most reactive of this group, it will react with fluorine and some other high oxidizing molecules, but not to any extent that the amateur chemist should have to worry about.

· Sulfur Hexafluoride  :  Fairly inert gas used in blanketing the setup used for the electrochemical production of magnesium metal.  Not widely available, dense gas, not exceedingly toxic, TLV 1,000 PPM.

        Taking care of exit gasses depends on what exactly is contaminating them.  Disposing of radioactive chemicals by incineration of exit gasses is not advisable but the decomposition of organic material by leading exit gasses through a flame works in most cases.  If there is a specific component in the gas that you expect to survive the flame then take the extra step to bubble the exit gasses through an appropriate solution to neutralize the offending contaminate.  Occasionally a reaction will be called for to run under a specific gas that does not seem to be inert, it may well not be.  If something is run under a chlorine atmosphere you can be that it is probably necessary for the reaction to commence.  

        When discontinuing use of the inert atmosphere the vessel must become slowly accustomed to regular atmosphere again, unless the apparatus is to be disassembled under an inert atmosphere in a fume hood and cleaned there.  Unless the reaction products are explosive (boranes, phosphines, etc.), or exceedingly pyrophoric, simply removing the tube for the exit gas and turning off the flow of inert gas will suffice to bring the vessel back into atmospheric conditions as the gasses inside diffuse out and the gasses from the outside slowly work their way in.  Explosive mixtures between the air and flammable gasses may form so it is necessary to allow the apparatus to cool to room temperature before hand.

8.5 Solvent/Solvent extraction systems 

8.6 Vacuum Pumps and working under Vacuum

Vacuum distillation is necessary when your compound decomposes when being boiled at STP. It simply is distilling under reduced pressure and can be applied to all types of previously mentioned distillation. There are however a few important observations. Boiling points of different components don’t necessarily change in a linear way when the pressure is altered. Simply said, it could mean that whilst the difference in vapor pressure is 200mbar at STP, it could be 50mbar or 400mbar under reduced pressure.

Another problem is glassware and pumps. Your lab glass needs to be borosilicate glass and it should be in prime condition. Cracks or even scratches severely compromise the strength of your glass under vacuum. Only round-bottomed flasks can be used during vacuum distillation, no flat bottom flasks or erlenmeyers! What could possibly happen you think? Well, imagine a flask filled with boiling ethanol imploding. It suddenly comes into contact with a surge of fresh air and whooosh...flaming inferno all over and you being sprayed with glass pieces. That’s why. Not to mention what would happen with acids, poisonous or otherwise hazardous substances.

So your glassware is in perfect shape. On to the other problems, boiling stones don’t work under vacuum and boiling under vacuum can be very aggressive, so aggressive that the bumping can crack your glass, after which it implodes. Again, there are two solutions to this problem. A Claisen adapter (picture?) with a capillary tube that is immersed in the liquid and then provides bubbles is a way to go, or magnetic stirring, the authors personal preference.

It goes without saying that your joints should be sealed well. This can be accomplished by grease or commercially available teflon tape or specially designed teflon joint fitters, which are expensive. Don’t neglect this aspect, because a joint which fails after a while is another possible doom for you and your set up.

Special care should also be taken in the way you fix your apparatus. You must avoid any stress or strain caused by hanging flasks, flask that are being pushed up because they float in the water bath, etc. Carefully balance your apparatus before applying your vacuum.

Heating is not an easy task when using vacuum. Flames or any other forms of localized heating are a NO GO. A water or oil bath is preferable. You must also realize that your set ups heat conductance is much lower because of the insulating factor of the vacuum. That’s why the author also recommends magnetic stirring of your bath, because this dramatically increases heat transfer. Otherwise you risk a bath, which is at 90C while the inside of your flask is at 30C, which causes a huge amount of stress and is generally energy consuming.

Finally, wrapping your glass with a wire mesh or laminating it with heat resistant plastic protects you from flying glass should something go wrong.

When first starting a procedure involving distillation using vacuum there is a sequence of events that one should follow:

Applying vacuum should be done FIRST, before heating starts. If you start heating first and then apply vacuum, there is a severe hazard that your liquid will flash boil when the pressure drops. Flash boiling usually comes with a pressure spike, insane bumping and a lot of frothing. Short & sweet, it means death to your setup and possibly to you. You also should start up your stirring or bubbling at the same time your heating starts, because if you do it too late, there is again a risk of flash boiling.

When you're done distilling, allow the apparatus to cool down first, then allow air to enter. If done otherwise, hot solvent vapor can come into contact with fresh air and reach the explosive range...which again means flaming inferno all over and flying glass.

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9.0 When things go wrong 

9.1 Contingency Plans 

        In any reaction that has an abnormal risk it is always good to have a contingency plan.  How can you determine if a reaction has an abnormal risk?  In my opinion if you have to worry about a reaction then it has an abnormal risk, however you shouldn't have to plan for every possible screw up.  Here is an example, you add a piece of sodium metal to some ethanol with the intent to make sodium ethoxide.  However you find that your ethanol must be contaminated with a large water percentage, too bad you found this out when you added your Na to your ethanol.  It's boiling and bubbling and H2 is coming off it like there's no tomorrow.  If that H2 builds up the the reaction keeps up the Na will ignite it, possibly detonating your reaction vessel and thus spraying flammable liquid everywhere, most likely on fire itself, what are you going to do?  Your normal plan to dump the offending reaction on the ground and spray it with water seems to be bad but it's always been your backup before.  So you toss your reaction solution on the ground and spray it with your hose, the H2O hits the sodium and **Boom** not only does the sodium explode and spray the surrounding area with little chunks, but the ethanol was scattered too and is now burning merrily all over your grass and house and, wait, your arm's on fire.....

        Don't tell me you didn't see this coming, Na can be quite nasty when removed from its anhydrous environment that it stays so comfortable in.  You didn't have a back up plan and from here things could get even nastier.  So, the obvious solution is to come up with a backup plan.  So what could you have done in retrospect.... Sand is good for metal fires, that could have been a decent idea, maybe instead of water you could have tossed it in ice, that might have been slightly better but not a lot, how about having some real anhydrous ethanol on hand to dissolve excess sodium in, you know, like they do in professional labs.  Regardless, it's always good to figure out what might go wrong before hand then to figure out why things went wrong after the fact.  MSDS sheets can give a good indication of what to do incase a reagent gets out of control.  In addition just knowing a chemicals properties can help.

9.2 Don’t Mix .... or.... 

9.2a Explosive mixtures involving oxidizing agents 

9.2b Unstable Peroxides 

9.3 Flammability Concerns 

May of the reagents, particularly the solvents used in chemistry are quite flammable.  However some posses specific flammability concerns.  Highly flammable solvents include ether, which is known to creep along the ground for an ignition source, and carbon disulfide which can ignite from boiling water.  It is good to know if you keep these in an enclosed environment without oxygen they will not spontaneously ignite, however they do get out.  So as a common practice flammable liquids should not be heated with an open flame unless you are quite familiar with what might happen.

        Flammability is always augmented if the word pyrophoric is involved.  Some metal powders are pyrophoric (zirconium powder being a good example), some other solids are too (white phosphorus), as are some gasses (diphosphine), and liquids.  After just a short contact with the atmosphere, either from reacting with ambient moisture or oxygen these may ignite on their own.

9.4 When to abandon Ship 

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10.0 Finding things at home 

To some extent people have a chemical stockpile in their houses even if they are not practicing chemists.  For example many house holds have acetic acid, sodium chloride, sodium bicarbonate, sodium hypochlorite solutions, and more.  But that is not the extent of the chemicals available at your local hardware store or super market, or if you want to go even further specialty stores like hydroponics stores can be a uranium mine for the amateur chemist.  When ever you visit these places just keep an eye on the shelves and if something catches your eye look over the label for information relating to the compounds contained within.  This can give you the best idea of what you have available in your area.

        The second best alternative is the internet.  Not only are there sites dedicated to chemicals found at home, you can search compounds on google or other places and attempt to find them in some household use.  Some are pure some are not and some are easily separated.

10.1 Pure compounds 

10.2 Making Vs. Buying 

10.3 Extracting compounds 

10.3a Basic Principles (Comparing Properties) 

10.4 Mail Order 

10.5 Notes about purity 

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11.0 Advanced Experiments (Name Reactions) 

11.1 Canizzaro Rxn on Benzaldehyde (separation of products) 

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12.0 Index (Links throughout will be highlighted and click able to bring you to the specific index entry, e.g. H2SO4 will be highlighted and clicking on it will bring you to a page listing its properties, High temp oxidizing agent, dehydrating agent, different concentrations available.)

12.1 The Elements (See Section 1.3 for a depiction of the periodic table)

Actinium Atomic Symbol: Ac Atomic Number: 89 Atomic Weight: 227.0 g/mol Known oxidation state(s): +3

Hazard information: Highly radioactive, most stable isotope has a half-life of 22 years.

Aluminum Atomic Symbol: Al Atomic Number: 13 Atomic Weight: 27.0 g/mol Known oxidation state(s): +3

Hazard information: The presence of aluminum cations in soft drinks is a the suspect to some cases of Alzheimer’s. Aluminum dust poses two hazards, it can provide an environment that could possibly lead to an explosive mixture with the air and secondly it can cause irritation to the respiratory system and disorientation. Always wear gloves and a dust mask when working with aluminum in the powder form. Bulk aluminum is safe.

Additional information on Aluminum: Aluminum as a bulk metal is widely used in the building industry. It is easily spotted in a scrap yard for a few reasons, it is relatively light, and forms an oxide coating which is easily scraped off with a knife to reveal the clean metal underneath. Carry a small bottle of vinegar with you if you are hunting for aluminum in a scrap yard to test samples, scrape the surface of the aluminum clean and apply a little of the acid, it will react with aluminum forming bubbles if it is the real deal. Aluminum turnings are also available at some scrap yards. Aluminum powder is available from pyrotechnic suppliers. There are also guides online for turning bulk aluminum to powder. Aluminum powder cannot be made by the decomposition of aluminum formate or oxylate as the finely divided aluminum can react readily with the carbon dioxide produced to form aluminum oxide as the majority product.

Industrially aluminum is produced by the Hall process, electrolysis of aluminum oxide held in a molten cryolite [Na₃AlF₆] bath. On a home scale such a process would be demanding at best. On a side interesting note one of the first uses of sodium was as a reductant for producing aluminum from the oxide. This process has since been replaced by the Hall process noted above.

Aluminum is a highly reactive metal, it reacts readily with atmospheric oxygen and would simply rust to a pile if the oxide coating thus produced did not adhere so well. If for example a small amount of mercury is placed on a block of aluminum it continuously alloys with the aluminum rendering the oxide coating ineffective and will allow the oxygen in the air to rapidly oxidize large amounts of aluminum. Aluminum will react with nearly any acid and many bases readily. Many aluminum salts are soluble and therefore are a good source of choice anions in solution.

Americium Atomic Symbol: Am Atomic Number: 95 Atomic Weight: 241.1 g/mol Known oxidation state(s): +3

Hazard information: Radioactive element, treat with care.

Additional information on Americium: Americium oxide is the source of ionization energy in the vast majority of smoke detectors. It is a very small piece of this radioactive element.

Antimony Atomic Symbol: Sb Atomic Number: 51 Atomic Weight: 121.76 g/mol Known oxidation state(s): +3, +4, +5 (least common)

Hazard information: Excessive handling of antimony metal should be avoided as many of the salts formed even those on contact with air could be hazardous. Antimony and its salts have been linked to reproductive damage and cancer.

Additional information on Antimony: Used in alloying, with lead in solder and in other applications, a hardening agent. Antimony is toxic and forms some interesting salts, the pentafluoride is a component of superacids but obtaining this metal in an over the counter way is difficult. Antimony sulfide is used in pyrotechnics.

Argon Atomic Symbol: Ar Atomic Number: 18 Atomic Weight: 40.0 g/mol Known oxidation state(s): No common oxidation states

Hazard information: Argon is an asphyxiant gas, use with ventilation. Argon directly exiting from cylinders may be cold enough to induce frost bite.

Additional information on Argon: (See section on inert atmospheres 8.4)

Arsenic Atomic Symbol: As Atomic Number: 33 Atomic Weight: 74.9 g/mol Known oxidation state(s): +2, +3, +5 (least common)

Hazard information: Excessive handling of arsenic metal should be avoided as many of the salts formed even those on contact with air could be hazardous. Arsenic and its salts have been linked to reproductive damage and cancer. Arsenic can show progressive physical and neurological damage, the progressive signs of arsenic poisoning are well covered. Arsenic trioxide was once known as “Inheritance powder”.

Additional information on Arsenic: The only widely available compound containing arsenic is arsenic trioxide, I have seen it marketed for the purpose of killing a variety of insects, in ant traps and less commonly to kill mice. It’s use has been phased out since the beginning of the 20^th century though. It is also found in some specialty solders and in semiconductors. From its trioxide it could be reduced with an active metal such as magnesium to form the metal. Another available form of arsenic comes in the form of some herbicides and pesticides which contain arsenic organic molecules. Arsenic is a chemically reactive metal with interesting properties especially evident in the covalency of its high oxidation state compounds.

Astatine Atomic Symbol: At Atomic Number: 85 Atomic Weight: 210.0 g/mol Known oxidation state(s): NA

Hazard information: Highly radioactive, most stable isotope has a half-life of 8 hours.

Barium Atomic Symbol: Ba Atomic Number: 56 Atomic Weight: 137.3 g/mol Known oxidation state(s): +2

Hazard information: Barium salts are highly toxic, a small amount of a soluble barium salt that makes its way into your body will make you have a very bad day, diarrhea, blood in stool, headache, stomach pains, etc. The metal itself is highly reactive towards water along the lines of sodium and can cause minor explosions and presents a flammability hazard on its own. The free metal will burn the skin if it comes into contact with it. Should be stored under oil, most reactive of the common alkali earth metals.

Additional information on Barium: When exposed to air barium will from an appreciable percentage of the peroxide. Very few barium salts are available to the general public, the few that I know of are barium sulfate which is obtainable from medical clearances (it is used to make the intestines more visible with though xray, it is one of the very few safe barium salts), and barium ferrate, which is present in the coating on VHS tapes. In theory a large quantity of VHS tape could be ashed (heated till it turned to ash) then reduced with an active metal (aluminum or magnesium) then dissolved in water, the barium oxide thus formed would react with the water and convert to the somewhat soluble barium hydroxide which could be extracted by evaporation and crystallization.

Furthermore barium is available in both the hobby of pyrotechnics (carbonate, nitrate, perchlorate, sulfate) and pottery (carbonate) for colorization. These can be scrounged up from local sources or from online sources. Barium metal could be produced by aluminothermic reduction of the oxide or carbonate or hydroxide and subsequent distillation under high vacuum. Reaction of barium oxide and aluminum metal at high heat furnishes an alloy of high barium percentage >50% on cooling. Barium can also be procured though electrolysis of an eutectic mixture of barium salts in the molten state.

Berkelium Atomic Symbol: Bk Atomic Number: 97 Atomic Weight: 249 g/mol Known oxidation state(s): +3, +4

Hazard information: Highly radioactive, half life sufficiently short to render amateur experimentation futile.

Beryllium Atomic Symbol: Be Atomic Number: 4 Atomic Weight: 9.0 g/mol Known oxidation state(s): +2

Hazard information: Beryllium salts and beryllium metal dust are highly toxic and carcinogenic.

Additional information on Beryllium: Some aircraft parts, specifically gyroscopes are occasionally made of almost entirely beryllium, easily differentiated by their unearthly lightness. Machining beryllium is dangerous as shavings and powder can cause ‘metal fume fever’ and terrible pain. Beryllium is a reactive metal that forms an oxide coating that prevents further atmospheric attack. It is hard to find on the civilan market though exept as the afformentioned use and in a very few copper alloys. Because of the beryllium ions small size and high charge density it forms unique cations when dissolved in water involving several water molecules.

Bismuth Atomic Symbol: Bi Atomic Number: 83 Atomic Weight: 209.0 g/mol Known oxidation state(s): +3, +5 (rare)

Hazard information: Bismuth is fairly benign and safe to handle, the toxicity of bismuth salts is almost entirely dependent upon the anion to which it is coupled.

Additional information on Bismuth: Bismuth is available as environmentally friendly buck shot for re-loading guns in areas where guns are permitted, but by this route it is fairly expensive. Also it can be found in some areas that sell minerals and collectable rocks, bismuth forms beautiful crystals when solidified from a properly formed melt and are sold as a pure material, again, the price can be exhorbant. The internet is always another choice for bismuth metal if all else fails.

Bismuth trioxide has found use in pyrotechnics and this could be reduced with an appropriate aluminothermic reduction. Also bismuth subsilicate is available as an over the counter stomach soothing remedy, it may be possible, although economically disastrous to extract this small quantity of bismuth. Some bismuth salts, especially those where bismuth is in the +3 state and attached to three different molecules are prone to decomposition in water due to the formation of the stable oxy compound. For example, a solution of bismuth trichloride left to stand may decompose in the following manner:

BiCl_3(aq) + H₂O_(l) Þ OBiCl_(s) + 2HCl_(aq)

Many compounds will do the same hydrolysis reaction if left in solution too long, bismuth nitrate may form bismuth subnitrate, bismuth chloride may precipitate as bismuth oxychloride and there are many more. Dissolving bismuth is a difficult chore although it comes ahead of hydrogen in the activity series and should theoretically dissolve in acid it does so sluggishly at best, it is necessary to add an oxidizing agent to get a decent rate of solvation of the native metal. And as I just mentioned it is necessary to recover your bismuth salt quickly lest it hydrolyze, the hydroxide is a good choice as it will allow conversion to other appropriate salts at a later date. The bismuthate anion BiO₃^- in which bismuth has a +5 charge is an excellent oxidizing agent prepared by the reaction of dry bismuth trioxide with sodium peroxide or by the action of molten NaNO₃/NaOH on bismuth trioxide, it will oxidize manganate to permanganate.

Boron Atomic Symbol: B Atomic Number: 5 Atomic Weight: 10.9 g/mol Known oxidation state(s): +3

Hazard information: Elemental boron is toxic, dust should be avoided, boron compounds differ widely in their toxicity, for example, the chloride is a strong irritant/corrosive liquid, whereas the acid is the only acid that is actually good for the eyes.

Additional information on Boron: Boron has two widely available salts, borates/metaborates are available to some extant as borax in the cleaning industry. Borax as found in cleaning products usually has the formula Na₂B₄O₇*5H₂O solutions of borax can be treated with a strong acid such as HCl to precipitate out boric acid. Boric acid can also be bought as a somewhat pure substance from pharmacies and also from grocery stores for the purpose of pest control. From boric acid heat can be applied to dehydrate it to boric oxide. And from the oxide elemental boron can be had.

Na₂B₄O_7(aq) + 2HCl_(aq) + 5H₂O_(l) Þ 2NaCl_(aq) + 4B(OH)_3(s)

2B(OH)_3(s) --(Heat)--> B₂O_3(s) + 3H₂O_(g)

B₂O_3(s) + 3Mg_(s) --(Heat)--> 2B_(s) + MgO_(s) + x[MgB₂]_(s)

B(s)/MgO(s)/MgB₂(s) --(HCl_(aq))--> B_(s) + MgCl_2(aq) + B₂H_6(g)

In the above reactions we start from the commonly available borax with a precipitation reaction to get to our boric acid. Of if you have boric acid start from step 2. From here the acid is dehydrated and easily goes to boric oxide. The oxide is then pulverized with a hammer or other suitable object and mixed with either magnesium powder or turnings in a stoichiometric amount. The mix is ignited and a thermite reaction ensues, this generates lots of heat but the reaction must be covered loosely immediately to prevent the oxidation of the boron thus formed, at the same time a small amount of magnesium boride is formed as a side reaction. Finally after the reaction cake has cooled and is powdered, it is digested in hydrochloric acid, the magnesium oxide being a basic oxide is readily dissolved in the HCl and the magnesium boride reacts with the HCl to produce diborane. The diborane is a spontaneously flammable gas and therefore small explosions may result, it is therefore advisable to cover the cake first with water then add acid in small amounts to prevent excessive sudden gas evolution. The magnesium chloride stays in solution, the boride is decomposed and what you are left with is boron as a precipitate at the bottom of the reaction vessel.

The first question that comes to many peoples mind when they see this thermite type reaction is weather they can substitute aluminum for the magnesium as aluminum powder is more readily made/acquired. Yes, it could be substituted in theory, but there is one drawback, see step 3 where the magnesium boride is formed as a side reaction. It could be assumed, and in this case correctly that aluminum boride [AlB₁₂] would be formed analogously in this reaction. But the problem comes in reaction 4, aluminum boride is very inert, it will not react with the HCl and therefore you end up with very impure boron as you are unable to separate the aluminum boride (in addition the aluminum oxide is very hard to dissolve out). So what you are left with is a neigh insoluble mass of boron, aluminum oxide, and aluminum boride from which the boron is very difficult to remove. One possible removal method would be to run chlorine gas over the heated mass to produce boron trichlroide and run that over heated zinc powder to facilitate the reaction along the lines of :

B_(s) + Cl_2(g) Þ BCl_3(g)

2BCl_3(g) + 3Zn_(s) Þ 2B_(s) + 3ZnCl_2(s)

Although this method facilitates boron powder it makes the reaction considerably more difficult in the manipulation of chlorine gas and boron trichloride. However one could make boron trichloride directly from boric oxide and sodium chloride and run that over the zinc therefore skipping the active metal reduction with magnesium and replacing it with this zinc step.

Boron will form an additional bond at its lone pair making it a negative cation, an excellent example of this is sodium borohydride [NaBH₄] in which the boron atom has a negative one charge due to the extra bond to hydrogen.

Boron is an elemental color emitter, its combustion produces a beautiful green color and its esters produce the same.

Bromine Atomic Symbol: Br Atomic Number: 35 Atomic Weight: 79.9 g/mol Known oxidation state(s): -1, +3, +5, +7 (rare)

Hazard information: Bromine is a highly corrosive red liquid. It will attack rubber, your lungs (causing pulmonary edema), your eyes (causing blindness), and your skin (causing painful ulcerations). Skin exposure should be treated with a reducing agent such as sodium thiosulfate which will help to destroy the bromine before it destroys any more of you. Although it is not highly toxic it does have sedative effects that can result in death due to depression of the central nervous system.

Additional information on Bromine: (See section 4.9 for further information) Free bromine is found as the diatomic molecule Br₂ that is whenever bromine is free it is always coupled with another bromine molecule. Your best bet to finding commercially available bromine sources is going to be from pool/spa suppliers. Bromination sources include sodium bromide but more often you may find a complex organic compound that actually acts as the brominating agent. If possible the sodium bromide provides the much easier compound from which to extract bromine although the organic compound could yield a combination of bromine and bromine chloride (although this decomposes above 10C, the chlorine gas that makes its way though and comes into contact with your condensed bromine in a receiving flask could react with the bromine there). That is if it is sufficiently gassed with chlorine in powder form at a temperature sufficient to distill off the bromine [>59C].

As for bromine production from sodium bromide. 1) Running chlorine gas though a solution of warm sodium bromide will cause the chlorine to replace the bromine in the compound resulting in free bromine. This reaction really is complicated by working with chlorine gas. 2) Reacting aqueous sodium bromide with an oxidizing agent under acidic conditions can result in the formation of free bromine which can be distilled off:

2NaBr_(aq) + H₂SO_4(l) + H₂O_2(aq) Þ Na₂SO_4(aq) + 2H₂O_(l) + Br_2(l)

In the above reaction it is the hydrogen peroxide that acts as the oxidizing agent, other oxidizing substances; potassium permanagnate, potassium bromate; etc. could be used in its place. Additionally different acids could be substituted, hydrochloric acid could be substituted but there is the possibility that it could be oxidized resulting in free chlorine contaminating the reaction. An additional benefit to the addition of concentrated H₂SO₄ is the heat of hydration which allows the mixture to obtain a temperature to distill off the Br₂ formed without the need for significant, if any, additional heating.

Cadmium Atomic Symbol: Cd Atomic Number: 48 Atomic Weight: 112.4 g/mol Known oxidation state(s): +2

Hazard information: Highly toxic, carcinogenic, poisoning from cadmium compounds is rare though due to their ability to induce vomiting rapidly.

Additional information on Cadmium: Cadmium serves very few purposes in the life of the general populous. One of the only sources of any form of cadmium, aside from meager alloys and coatings, is found inside of household rechargeable batteries. This is in the form of a cadmium oxide electrode. Another source of cadmium is in the form of pigments, cadmium sulfide (yellow-brown) and selenide (red) being the main ones. Cadmium sulfide could be dissolved in dilute HCl and the mixture heated to reflux, hydrogen sulfide would be evolved though which is highly toxic. The resulting CdCl₂ could be re-dissolved in neutral water and the solution electrolyzed to yield the metal. Cadmium is resistant to alkalis but readily attacked by acids.

Calcium Atomic Symbol: Ca Atomic Number: 20 Atomic Weight: 40.1 g/mol Known oxidation state(s): +2

Hazard information: Flammable as a bulk solid, spontaneously flammable in powder/fine turnings. Calcium is non-toxic but it can cause skin damage if handled without gloves from the basicity of the hydrolyzed metal and the dehydrating action on the skin. Reacts readily with water forming hydrogen gas, which can ignite and explode.

Additional information on Calcium: Calcium is produced most often by the electrolysis of straight molten CaCl2, in this process the cathode must either be barely touching the surface of the melt and slowly raised up or, constantly rotated to provide a cohesive non-porous mass of calcium metal. The addition of up to 15% KCl can depress the melting point of the mixture without noticeable potassium formation at the cathode but at percentages beyond this potassium formation becomes evident. Additionally mixing calcium chloride with chlorides of other alkali earth metals can form eutectics which may prove useful, but despite finding patents on such mixtures, they have found no use in industry. During electrolysis of the molten chloride there is a very small range over which electrolysis can progress successfully, between 780 and 800 C, during this small frame calcium produced will be a solid and the melt will be molten, lower then this and the melt solidifies, higher and the already highly reactive calcium will be molten and almost guaranteed to catch fire. Remember, chlorine gas would be produced at the anode to complicate matters even further.

Chemical reduction of calcium oxide is another route to calcium metal production. When calcium iodide and sodium metal are heated together in a metal vessel at high heat and the mixture allowed to cool, calcium metal crystallizes out. Aluminothermic reduction of calcium oxide with aluminum metal over high heat under high vacuum has been used to isolate calcium metal, however it does not work as well as similar reductions of other heavier alkali earths.

Calcium itself is a great reducing agent due to the low volatility of its oxide and chloride. Heating cesium/rubidium/potassium chlorides with calcium metal under high vacuum will distill over the free metals. Calcium carbide, a somewhat available chemical can also act as a potent reducing agent.

Californium Atomic Symbol: Cf Atomic Number: 94 Atomic Weight: 251.1 g/mol Known oxidation state(s): +3, +4

Hazard information: Highly radioactive element. However the half-life is long enough to work with the element in macroscopic quantities. Cf²⁵² is the most widely available isotope and is for sale in milligram quantities. The isotope with the longest half-life is Cf²⁵¹ with a half-life of nearly 900 years.

Carbon Atomic Symbol: C Atomic Number: 6 Atomic Weight: 12.01 g/mol Known oxidation state(s): -4, +4 (Carbon can form hybrid orbitals resulting in unique states)

Cerium Atomic Symbol: Ce Atomic Number: 58 Atomic Weight: 140.1 g/mol Known oxidation state(s): +3, +4

Hazard information:

Additional information on Cerium: The decomposition of cerium oxalate by heat results in the formation of cerium dioxide instead of the expected free metal.

Cesium Atomic Symbol: Cs Atomic Number: 55 Atomic Weight: 132.9 g/mol Known oxidation state(s): +1

Chlorine Atomic Symbol: Cl Atomic Number: 17 Atomic Weight: 35.5 g/mol Known oxidation state(s): -1, +1, +3 (rare), +4, +5, +7

Chromium Atomic Symbol: Cr Atomic Number: 24 Atomic Weight: 52.0 g/mol Known oxidation state(s): +2, +3, +6

Cobalt Atomic Symbol: Co Atomic Number: 27 Atomic Weight: 58.9 g/mol Known oxidation state(s): +2, +3

Copper Atomic Symbol: Cu Atomic Number: 29 Atomic Weight: 63.5 g/mol Known oxidation state(s): +1, +2

Curium Atomic Symbol: Cm Atomic Number: 96 Atomic Weight: 247.1 g/mol Known oxidation state(s): +3, +4

Hazard information: Curium is a radioactive bone-seeking element. It is available in gram quantities but is quite expensive and outside the price range of the average at home chemist.

Dysprosium Atomic Symbol: Dy Atomic Number: 66 Atomic Weight: 162.5 g/mol Known oxidation state(s): +3

Hazard information: Spontaneously flammable in powder form, reacts slowly with water and halogens.

Additional information on Dysprosium: Formed by the reduction of its fluoride with calcium Dysprosium is a rare element with which you have little probability to run across.

Einsteinium Atomic Symbol: Es Atomic Number: 99 Atomic Weight: 253 g/mol Known oxidation state(s): +2

Hazard information:

Erbium Atomic Symbol: Er Atomic Number: 68 Atomic Weight: 167.3 g/mol Known oxidation state(s): +3

Hazard information:

Europium Atomic Symbol: Eu Atomic Number: 63 Atomic Weight: 152.0 g/mol Known oxidation state(s): +2, +3

Hazard information: Highly reactive, spontaneously flammable and reactive with water.

Fermium Atomic Symbol: Fm Atomic Number: 100 Atomic Weight: 254 g/mol Known oxidation state(s): +3

Hazard information: Highly radioactive, most stable isotope only has a half-life of 3 hours.

Fluorine Atomic Symbol: Fl Atomic Number: 9 Atomic Weight: 19.0 g/mol Known oxidation state(s): -1

Francium Atomic Symbol: Fr Atomic Number: 87 Atomic Weight: 223.0 g/mol Known oxidation state(s): +1

Hazard information: Highly radioactive, most stable isotope only has a half-life of 3 hours. Less then 25 g on Earth at the same time.

Gadolinium Atomic Symbol: Gd Atomic Number: 64 Atomic Weight: 157.3 g/mol Known oxidation state(s): +3

Gallium Atomic Symbol: Ga Atomic Number: 31 Atomic Weight: 70.0 g/mol Known oxidation state(s): +2, +3

Germanium Atomic Symbol: Ge Atomic Number: 32 Atomic Weight: 72.9 g/mol Known oxidation state(s): +2, +4

Gold Atomic Symbol: Au Atomic Number: 79 Atomic Weight: 197.0 g/mol Known oxidation state(s): +1, +3

Hafnium Atomic Symbol: Hf Atomic Number: 72 Atomic Weight: 178.5 g/mol Known oxidation state(s): +2, +3, +4, +6

Helium Atomic Symbol: He Atomic Number: 2 Atomic Weight: 4.0 g/mol Known oxidation state(s): No common oxidation states

Holmium Atomic Symbol: Ho Atomic Number: 67 Atomic Weight: 164.3 g/mol Known oxidation state(s): +3

Hydrogen Atomic Symbol: H Atomic Number: 1 Atomic Weight: 1.0 g/mol Known oxidation state(s): -1, +1

Indium Atomic Symbol: In Atomic Number: 49 Atomic Weight: 114.8 g/mol Known oxidation state(s): +1, +3

Iodine Atomic Symbol: I Atomic Number: 53 Atomic Weight: 126.9 g/mol Known oxidation state(s): -1, +1, +3, +5, +7

Iridium Atomic Symbol: Ir Atomic Number: 77 Atomic Weight: 192.2 g/mol Known oxidation state(s): +1, +2, +3, +4, +6

Iron Atomic Symbol: Fe Atomic Number: 26 Atomic Weight: 55.9 g/mol Known oxidation state(s): +2, +3, +4 (rare), +5 (unstable), +6 (rare), +7 (rare)

Krypton Atomic Symbol: Kr Atomic Number: 36 Atomic Weight: 83.8 g/mol Known oxidation state(s): +2 (rare)

Lanthanum Atomic Symbol: La Atomic Number: 57 Atomic Weight: 138.9 g/mol Known oxidation state(s): +3

Lawrencium Atomic Symbol: Lr Atomic Number: 103 Atomic Weight: 262.1 g/mol Known oxidation state(s): NA

Hazard information: Would be highly radioactive as the most abundant isotope only has a half-life of 8 seconds.

Lead Atomic Symbol: Pb Atomic Number: 82 Atomic Weight: 207.2 g/mol Known oxidation state(s): +2, +4

Lithium Atomic Symbol: Li Atomic Number: 3 Atomic Weight: 6.94 g/mol Known oxidation state(s): +1

Lutetium Atomic Symbol: Lu Atomic Number: 71 Atomic Weight: 175.0 g/mol Known oxidation state(s): +3

Magnesium Atomic Symbol: Mg Atomic Number: 12 Atomic Weight: 24.3 g/mol Known oxidation state(s): +2

Manganese Atomic Symbol: Mn Atomic Number: 25 Atomic Weight: 54.9 g/mol Known oxidation state(s): +2, +3, +4, +6, +7

Mendelevium Atomic Symbol: Md Atomic Number: 101 Atomic Weight: 258.1 g/mol Known oxidation state(s): +2, +3

Hazard information: Longest half-life of a mendelevium isotope is just shy of two months. This highly radioactive element is not something you will likely run across.

Mercury Atomic Symbol: Hg Atomic Number: 80 Atomic Weight: 200.6 g/mol Known oxidation state(s): +1 (diatomic), +2

Molybdenum Atomic Symbol: Mo Atomic Number: 42 Atomic Weight: 95.5 g/mol Known oxidation state(s): +2, +3

Neodymium Atomic Symbol: Nd Atomic Number: 60 Atomic Weight: 144.2 g/mol Known oxidation state(s): +3

Neon Atomic Symbol: Ne Atomic Number: 10 Atomic Weight: 20.2 g/mol Known oxidation state(s): No Common Oxidation States

Neptunium Atomic Symbol: Np Atomic Number: 93 Atomic Weight: 237.1 g/mol Known oxidation state(s): +5

Nickel Atomic Symbol: Ni Atomic Number: 28 Atomic Weight: 58.7 g/mol Known oxidation state(s): +2, +3

Niobium Atomic Symbol: Nb Atomic Number: 41 Atomic Weight: 92.9 g/mol Known oxidation state(s): +2, +3, +4, +5

Nitrogen Atomic Symbol: N Atomic Number: 7 Atomic Weight: 14.0 g/mol Known oxidation state(s): -3, +5

Nobelium Atomic Symbol: No Atomic Number: 102 Atomic Weight: 259.1 g/mol Known oxidation state(s): NA

Hazard information: Although nobelium has nine known isotopes, none of them have a long enough existence to determine any of the physical or chemical properties of this element.

Osmium Atomic Symbol: Os Atomic Number: 76 Atomic Weight: 190.2 g/mol Known oxidation state(s): +2, +3, +4, +6, +8

Oxygen Atomic Symbol: O Atomic Number: 8 Atomic Weight: 16.0 g/mol Known oxidation state(s): -2, +2 (rare)

Palladium Atomic Symbol: Pd Atomic Number: 46 Atomic Weight: 106.4 g/mol Known oxidation state(s): +2, +4

Phosphorus Atomic Symbol: P Atomic Number: 15 Atomic Weight: 31.0 g/mol Known oxidation state(s): -3, +3, +5

Platinum Atomic Symbol: Pt Atomic Number: 78 Atomic Weight: 195.1 g/mol Known oxidation state(s): +2, +4

Plutonium Atomic Symbol: Pu Atomic Number: 94 Atomic Weight: 239.1 g/mol Known oxidation state(s): +3, +4, +5, +6

Polonium Atomic Symbol: Po Atomic Number: 84 Atomic Weight: 210.0 g/mol Known oxidation state(s): +2, +4

Potassium Atomic Symbol: K Atomic Number: 19 Atomic Weight: 39.1 g/mol Known oxidation state(s): +1

Praseodymium Atomic Symbol: Pr Atomic Number: 59 Atomic Weight: 141.0 g/mol Known oxidation state(s): +3

Promethium Atomic Symbol: Pm Atomic Number: 61 Atomic Weight: 146.9 g/mol Known oxidation state(s): +3

Protactinium Atomic Symbol: Pa Atomic Number: 91 Atomic Weight: 231.0 g/mol Known oxidation state(s): +5

Radium Atomic Symbol: Ra Atomic Number: 88 Atomic Weight: 226.0 g/mol Known oxidation state(s): +2

Radon Atomic Symbol: Rn Atomic Number: 86 Atomic Weight: 222.2 g/mol Common oxidation statse: +2, +4, +6 (rare)

Rhenium Atomic Symbol: Re Atomic Number: 75 Atomic Weight: 186.2 g/mol Known oxidation state(s): +1, +2, +3, +4 (stable), +5, +6 (stable), +7 (stable)

Rhodium Atomic Symbol: Rh Atomic Number: 45 Atomic Weight: 102.9 g/mol Known oxidation state(s): +3

Rubidium Atomic Symbol: Rb Atomic Number: 37 Atomic Weight: 85.5 g/mol Known oxidation state(s): +1

Ruthenium Atomic Symbol: Ru Atomic Number: 44 Atomic Weight: 101.1 g/mol Known oxidation state(s): +3, +4, +5, +6, +8

Samarium Atomic Symbol: Sm Atomic Number: 62 Atomic Weight: 150.4 g/mol Known oxidation state(s): +3

Scandium Atomic Symbol: Sc Atomic Number: 21 Atomic Weight: 45.0 g/mol Known oxidation state(s): +3

Selenium Atomic Symbol: Se Atomic Number: 34 Atomic Weight: 79.0 g/mol Known oxidation state(s): -2, +2, +4, +6

Silicon Atomic Symbol: Si Atomic Number: 14 Atomic Weight: 28.1 g/mol Known oxidation state(s): -4, +4

Silver Atomic Symbol: Ag Atomic Number: 47 Atomic Weight: 107.9 g/mol Known oxidation state(s): +1, +2 (rare)

Sodium Atomic Symbol: Na Atomic Number: 11 Atomic Weight: 23.0 g/mol Known oxidation state(s): +1

Strontium Atomic Symbol: Sr Atomic Number: 38 Atomic Weight: 87.6 g/mol Known oxidation state(s): +2

Sulfur Atomic Symbol: S Atomic Number: 16 Atomic Weight: 32.1 g/mol Known oxidation state(s): -2, +2, +4, +6

Tantalum Atomic Symbol: Ta Atomic Number: 73 Atomic Weight: 181.0 g/mol Known oxidation state(s): +2, +3, +5

Technetium Atomic Symbol: Tc Atomic Number: 43 Atomic Weight: 99.0 g/mol Known oxidation state(s): +4, +5, +6, +7

Tellurium Atomic Symbol: Te Atomic Number: 52 Atomic Weight: 127.6 g/mol Known oxidation state(s): -2, +2, +4, +6

Terbium Atomic Symbol: Tb Atomic Number: 65 Atomic Weight: 158.9 g/mol Known oxidation state(s): +3, +4

Thallium Atomic Symbol: Tl Atomic Number: 81 Atomic Weight: 204.4 g/mol Known oxidation state(s): +1, +3

Thorium Atomic Symbol: Th Atomic Number: 90 Atomic Weight: 232.0 g/mol Known oxidation state(s): +4

Thulium Atomic Symbol: Tm Atomic Number: 69 Atomic Weight: 168.9 g/mol Known oxidation state(s): +3

Tin Atomic Symbol: Sn Atomic Number: 50 Atomic Weight: 118.7 g/mol Known oxidation state(s): +2, +4

Titanium Atomic Symbol: Ti Atomic Number: 22 Atomic Weight: 47.9 g/mol Known oxidation state(s): +3, +4

Tungsten Atomic Symbol: W Atomic Number: 74 Atomic Weight: 183.9 g/mol Known oxidation state(s): +2, +4, +5, +6

Uranium Atomic Symbol: U Atomic Number: 92 Atomic Weight: 238.0 g/mol Known oxidation state(s): +3, +4, +6

Vanadium Atomic Symbol: V Atomic Number: 23 Atomic Weight: 50.9 g/mol Known oxidation state(s): +2, +3, +4, +5

Xenon Atomic Symbol: Xe Atomic Number: 54 Atomic Weight: 131.3 g/mol Known oxidation state(s): +2, +4, +6, +8

Ytterbium Atomic Symbol: Yb Atomic Number: 70 Atomic Weight: 173.0 g/mol Known oxidation state(s): +2, +3

Yttrium Atomic Symbol: Y Atomic Number: 39 Atomic Weight: 88.9 g/mol Known oxidation state(s): +3

Zinc Atomic Symbol: Zn Atomic Number: 30 Atomic Weight: 65.4 g/mol Known oxidation state(s): +2

Zirconium Atomic Symbol: Zr Atomic Number: 40 Atomic Weight: 91.2 g/mol Known oxidation state(s): +2, +3, +4

Compounds

Teflon® Teflon®, is the brand name of a polymer produced by Dupont named PolyTetraFluorEthylene or PTFE for short. A Dupont researcher accidentally discovered this compound when he noticed there was no more pressure on his vessel, which contained tetrafluorethylene gas. He found a snow-white condensation product which proved to have exceptional chemical resistance.

Teflon is the compound of choice for the amateur chemist when he needs a very resistant and yet not extremely expensive material. The only problem with teflon is that it is a thermoplast and thus it weakens and eventually melts when heated too much. Compared to usual plastics its heat resistance is far higher, it can be safely employed between -200C and +250C.

Another noteworthy fact is that Teflon is insoluble in every solvent below 300C. Teflon should NOT be exposed to temperatures above 400C because it will decompose into several fluorocarbon nasties, which can severely damage your health.

Because of the exceptionally strong fluorine-carbon bond, Teflon resists the most aggressive chemicals, including fluorine gas or ozone. The only applications where it can't be employed are those where it comes into contact with very strong reducing agents, because of the fluorine content Teflon can act as an oxidizer in special circumstances. These reductants are mainly alkali and to a lesser degree alkali earth metals. The Air Force uses Teflon + Mg flares (although hard to ignite) to distract heat seeking missiles because it burns hotter than an aircraft exhaust, so be warned.

Teflon is OTC available mainly as tape, for sealing pipe joints in plumbing, and as sheet, for baking without the use of fat. Teflon tape (if it's pure, it should be white) is a very good substitute for joint grease because it won't contaminate your distillate, yet it provides good sealing. Most teflon baking sheets are black, so it probably is not pure, but it is the material of choice for applications where elevated temperatures are needed. Note that some baking sheets are made out of ICFLON, an unknown propriety compound.

Technical Terms

13.0 Appendix  (Specific Procedures/Additional Experiments)

13.1 Salts (Nearly)Insoluble in cold water:

Calcium Oxylate (AS)	Ca(COO)2	6.7 x 10^-4g/100ml	Calcium Sulfate	CaSO4 or CaSO4 * 2H2O	.241 g/100ml
Barium Sulfate	BaSO4	2.22 x 10^-4 g/100ml	Silver Chloride	AgCl	8.9 x 10^-5 g/100ml
Silver Bromide	AgBr	8.4 x 10^-6 g/100ml	Silver Iodide	AgI	3 x 10^-7 g/100ml
Magnesium Hydroxide (AS)	Mg(OH)2	9.0 x 10^-4 g/100ml	Aluminum Fluoride	AlF3 of AlF3*3H2O	----
Barium Carbonate (AS)	BaCO3	2.0 x 10^-3 g/100ml	Barium Chromate (AS)	BaCrO4	3.4 x 10^-4 g/100ml
Barium Citrate	Ba3 (C6H5O7)2 *7H2O	4.06 x 10^-2 g/100ml	Barium Oxylate (AS)	Ba(COO)2	9.3 x 10^-3 g/100ml
Barium Phosphate (AS)	Ba3(PO4)2	------	Barium Sulfite	BaSO3	2.0 x 10^-2 g/100ml
Bismuth Hydroxide (AS)	Bi(OH)3	1.4 x 10^-4 g/100ml	Cadmium Carbonate (AS)	CaCO3	-----
Camium Hydroxide (AS)	Cd(OH)2	2.6 x 10^-4 g/100ml	Cadmium Oxalate (AS)	Cd(COO)2	3.3 x 10^-3 g/100ml
Calcium Carbonate (AS)	CaCO3	1.45 x 10^-3 g/100ml	Calcium Fluoride	CaF2	1.6 x 10^-3 g/100ml
Calcium Hydroxide (AS)	Ca(OH)2	1.8 x 10^-1 g/100ml	Calcium Phosphate (AS)	Ca3(PO4)2	2.5 x 10^-3 g/100ml
Calcium Metasilicate (AS)	CaSiO3	9.5 x 10^-3 g/100ml	Cesium Aluminum Sulfate	CsAl(SO4)2 * 12H2O	3.4 x 10^-1 g/100ml
Cobalt (II) Carbonate (AS)	CoCO3	-----	Cobalt (II) Chromate (AS) (OxA)	CoCrO4	-----
Cobalt (II) and (III) hydroxide (AS)	Co(OH)2 & Co(OH)3	-----	Copper (I) & (II) Carbonate (AS) (Am)	Cu2CO3 & CuCO3	-----
Copper (I) Halides (AS) (Am)	CuX	-----	Copper (I) & (II) Hydroxide (Am)	CuOH & Cu(OH)2	-----
Lead Phosphate (AS) (BS) (OxA)	Pb3(PO4)2	1.4 x 10^-5 g/100ml	Mercury (I) Chloride	Hg2Cl2	2.1 x 10^-4g/100ml
Lead (II) Chloride	PbCl2	6.73 x 10^-1 g/100ml	Mercury (II) Sulfide	HgS	1.0 x 10^-6 g/100ml
Lead (II) Sulfide (OxA) (AS)	PbS	1.2 x 10^-2 g/100ml	Copper (I) Sulfide (OxA) (Am)	Cu2S
Copper (II) Sulfide (OxA) (Am)	CuS	3.3 x 10^-5 g/100ml	Cadmium (II) Sulfide (OxA)	CdS	-----
Arsenic Sulfide (OxA)	As2S3	-----	Antimony Sulfide (OxA)	Sb2S3	-----
Tin (IV) Sulfide (OxA)	SnS2	-----	Aluminum Hydroxide (AS) (BS)	Al(OH)3	-----
Iron (II) Sulfide (OxA)	FeS	-----	Manganese (II) Sulfide	MnS	-----
Zinc Sulfide (OxA)	ZnS	-----	Nickel (II) Sulfide (OxA)	NiS	-----
Cobalt (II) Sulfide (OxA)	CoS	-----	Strontium Phosphate	Sr3(PO4)2	-----
Zinc Hydroxide (AS) (BS)	Zn(OH)2	-----	Chromium (III) Hydroxide (AS) (BS)	Cr(OH)3	-----
Iron (III) Hydroxide	Fe(OH)3	-----	Copper (II) Oxylate (Am)	Cu(COO)2	2.5 x 10^-3 g/100ml
Bismuth (III) Sulfide (OxA)	Bi2S3	-----	Gold Sulfide (OxA)	Au2S	-----
Iron Ferricyanide	Fe3 [Fe(CN)6]2	-----	Lead (II) Bromide (AS)	PbBr2	4.5 x 10^-1 g/100ml
Lead (II) Chromate (BS) (AS)	PbCrO4	5.8 x 10^-6 g/100 ml	Lead (II) Carbonate (BS) (AS)	PbCO3	1.1 x 10^-4 g/100ml
Lead (II) Hydroxide (BS) (AS)	Pb(OH)2	1.5 x 10^-2 g/100ml	Lead (IV) Oxide (AS)	PbO2	-----
Lead (II) Oxylate	Pb(COO)2	1.6 x 10^-4 g/100ml	Lead (II) Sulfate	PbSO4	4.25 x 10^-3 g/100ml
Magnesium Carbonate (AS)	MgCO3	-----	Magnesium Fluoride (OxA)	MgF2	7.6 x 10^-3 g/100ml
Magnesium Oxylate (BS) (AS)	Mg(COO)2	7.0 x 10^-2 g/100ml	Magnesium Phosphate (AS)	Mg3(PO4)2	2.0 x 10^-2 g/100ml
Manganese (II) Fluoride (AS)	MnF2	-----	Manganese (II) Hydroxide (AS)	Mn(OH)2	2.0 x 10^-3 g/100ml
Manganese (II) Oxylate (AS)	Mn(COO)2	-----	Mercury (I) (AS) and (II) Bromide	Hg2Br2 & HgBr2	------
Mercury (I) and (II) Carbonate (AS)	Hg2CO3 & HgCO3	-----	Mercury (II) Phosphate (AS)	Hg3(PO4)2	------
Molybdenum (II) and (III) Bromide	MoBr2 & MoBr3	-----	Molybdenum (II) and (III) Chloride	MoCl2 & MoCl3	-----
Molybdenum Sulfides	Mo2S3 and MoS2	-----	Nickel (II) Carbonate (AS)	NiCO3	9.3 x 10^-3 g/100ml
Nickel (II) Fluoride	NiF2	2.0 x 10^-2 g/100ml	Nickel (II) Hydroxide (AS) (Am)	Ni(OH)2	1.3 x 10^-3 g/100ml
Nickel (II) Oxylate (AS)	Ni(COO)2	-----	Nickel (II) Phosphate (AS)	Ni3(PO4)2	-----
Potassium Perchlorate	KClO4	7.5 x 10^-1	Silver Carbonate (Am)	Ag2CO3	3.2 x 10^-3 g/100ml
Silver Oxide (AS)	Ag2O	1.3 x 10^-3 g/100ml	Silver Phosphate (Am) (AS)	Ag3PO4	6.5 x 10^-4 g/100ml
Silver Sulfate (Am) (AS)	Ag2SO4	5.7 x 10^-1 g/100ml	Tin Phosphate	Sn3(PO4)2	-----
Zinc Carbonate (AS) (BS)	ZnCO3	1.0 x 10^-3 g/100ml	Zinc Cyanide (BS)	Zn(CN)2	5.0 x 10^-4 g/100ml

(AS) = Increased solubility in acids (BS) = Increased solubility in bases (OxA) = Soluble in oxidizing acidic conditions (Am) = Can be rendered soluble in the presence of ammonia

14.0 Credits (Free Distribution Clause)

· Sections on Glassware and Pyrex partially written by and based off writings by 'Quantum (from USA)' from MSDB.

· Suggestions regarding topics were provided by ‘Magpie (from USA)’ from MSDB.

· The electrolysis section would have been near impossible without ‘Tacho (from Brazil)’ who wrote most everything, a member of the MSDB.

· Numerous parts, especially the distillation sections, were written by ‘Vulture (from Belgium)’ a member of the mad science discussion board.