---

Notice!

I’ve found that this book project has been showing up on more and more search engines lately and is also being directly linked to for the information it contains⁽¹⁾. I therefore find it necessary to warn all persons viewing this document that it is a work in progress, and as such it contains errors of all kinds, be them in experimental procedures that may cause harm, or in faulty reasoning that would get you slapped by nearly any chemistry instructor. Please for now take the information here with a grain of salt.

Most Importantly!

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 or following any of the procedures that are outlined within. Chemistry inherently possesses a degree of danger and you must understand this, wear gloves and more if the situation calls for it, your safety is in your own hands, not mine!

Also note that this project is open for contribution by any party on the internet. Simply submit a section to Rob.Vincent@gmail.com and it will be added into the text pending editing and such within a few weeks. Any person contributing will have their name mentioned in the credits. Thank you for reading this, and enjoy!

1	Although this document may be directly linked to, it will not work in that manner as I have hotlink protection for PDF documents, however directly linking to the html document is possible, still though I would prefer links be to the main book project page.

---

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)

pH

“pH”; everybody knows the term but what does it really stand for? And perhaps more importantly what is it good for?

pH stands for “potential of Hydrogen” (from the original German term “potenz”). It is a measure for the activity of hydrogen and because the activity of hydrogen in water equals the acidity of that water the pH effectively denotes the acidity of a solution. When hydrogen cations (H⁺ ions) are introduced into water they react with water to form the hydronium ion (also referred to as the oxonium ion) which is denoted as H₃O⁺(aq). The hydronium ion is the ion that gives acidic solutions their acidic nature. The direct opposite of the hydronium ion is the hydroxide ion (denoted as OH^-) which makes water alkaline. Water always contains H₃O⁺ ions and OH^- ions (hydroxide ions) but in pure water they are in equilibrium which means they cancel each other out for as the acidity of the water is concerned. When an acid is added to the water the equilibrium shifts to the acidic end of the spectrum which means more H₃O⁺ ions are present in the solution than OH^- ions. When a base is added the equilibrium shifts to the alkaline end of the spectrum which means more OH^- ions than H₃O⁺ ions are present in the solution. A pH from 0 to 7 means the solution is acidic (so more hydronium ions than hydroxide ions); a pH of 7 means the water is neutral (there are as many hydroxide ions as there are hydronium ions present in the solution) and a pH from 7 to 14 means the solution is alkaline (more hydroxide ions than hydronium ions).

The highest attainable pH at STP (standard temperature and pressure) is 14 and the lowest attainable pH at STP is 0. When the temperature and pressure do not conform perfectly to STP the minimum and maximum pH will vary accordingly. This is however not essential knowledge for hobby-chemists and as such I will not go into it any further.

Two kinds of acids and two kinds of bases.

There are two kinds of acids and bases; strong and weak. The strong versions do not form equilibriums in water but simply completely dissociate. The weaker versions will form an equilibrium in water and as such they will generally not be nearly as acidic or alkaline as the strong version.

Strong acids	In solution
Hydrochloric acid (HCl)	H3O^+(aq) + Cl^-(aq)
Sulphuric acid (H2SO4)	H3O^+(aq) + HSO4^-(aq)
Nitric acid (HNO3)	H3O^+(aq) + NO3^-(aq)
Perchloric acid (HClO4)	H3O^+(aq) + ClO4^-(aq)
Hydroiodic acid (HI)	H3O^+(aq) + I^-(aq)
Hydrobromic acid (HBr)	H3O^+(aq) + Br^-(aq)

Weak acids	In solution
Ethanoic acid (Acetic acid)	C2H3OOH(aq)
2-hydroxy-1,2,3-propanetricarboxylic acid (Citric acid)	C3H4OH(COOH)3(aq)
H3BO3 (Boric acid, aka B(OH)3 )	H3BO3(aq)
HF (Hydrofluoric acid)	HF(aq)
H3PO4 (Phosphoric acid)	H3PO4(aq)

The first difference that catches the eye is the fact that strong acids are denoted as free constituent ions and that weak acids are denoted as {weak acid}(aq). This is because the strong acids always completely dissociate in water whereas the weak acids only very partially dissociate in water.

Strong base	In solution
KOH (Potassium hydroxide)	K^+(aq) + OH^-(aq)
NaOH (sodium hydroxide)	Na^+(aq) + OH^-(aq)
Ca(OH)2 (calcium hydroxide)	Ca^2+(aq) + OH^-(aq)
Ba(OH)2 (barium hydroxide)	Ba^2+(aq) + OH^-(aq)
Na(C2H5O) (sodium ethoxide)	Na^+(aq) + OH^-(aq) + C2H5OH(aq)

Weak base	In solution
K2SO4 (potassium sulfate)	2K^+(aq) + OH^-(aq) + HSO4^-(aq) çè 2K^+(aq) + SO4^2-(aq)
K2CO3 (potassium carbonate)	2K^+(aq) + OH^-(aq) + HCO3^-(aq) çè 2K^+(aq) + CO3^2-(aq)
KF (potassium fuloride)	K^+(aq) + OH^-(aq) + HF(aq) çè K^+(aq) + F^‑(aq)
Na2HPO4 (sodium biphosphate)	Na^+(aq) + OH^-(aq) + H2PO4^-(aq) çè Na^+(aq) + HPO4^2-(aq)
NH3 (ammonia)	NH4^+(aq) + OH^-(aq) çè NH3(aq)

The obvious difference here is that in the case of the weak bases they are denoted as equilibriums when in solution whereas the strong bases (which do not form equilibriums) are denoted simply as their respective constituent ions in solution (so with “(aq)” at the end).

Calculations.

The pH of a solution can be calculated as follows:

pH = -log[H₃O⁺]

please note here that [H₃O⁺] denotes the hydronium concentration and that the addition of the p (for potential) effectively means you take the negative logarithm of the value without the p (in this case negative logarithm of the hydronium concentration). So when I take 2 moles of the “strong acid” HCl(g) and add enough water to make a solution of 1L of water I will have a 2M solution of hydrochloric acid. The pH of that solution would be –log(2) = -0.30 M = -3 * 10^-1. Which not only means the solution would be rather acidic but also immediately shows that the pH of a solution might be realistically expected to be between ~ -1,5 and ~ 15,5 rather than between the limits that are predicted by the rule. Don’t worry it really doesn’t make much of a difference when you’re using these formulae.

An important conclusion that can be drawn from this formula is that one can easily calculate the hydronium concentration (and as such the amount of hydrogen cations in solution) by means of the following calculation:

10^-pH

So when a solution has a pH of 3,7 the [H₃O⁺] = 10^-3.7 = 0,0001995 = 2 * 10^-4 mole/L (can also be denoted as 2 * 10^-4 M).

Ka and pKa

The degree to which an acid will dissociate in water is denoted by means of its Ka. The Ka is basically the dissociation constant of the acid in water. Ka’s can be denoted in a simpler form by taking their negative logarithm to yield the pKa.

pKa = -log(Ka)

These formulae can be used to calculate the amount of hydrogen cations and the and the pH of a solution of a certain strong acid. The calculations for a base can be performed in the same way but using the Kb in stead and the result of the calculation would yield the [OH^-] which can be transformed into the pH by means of the following formulae:

-log[OH^-] = pOH

14,00 – pOH = pH

This is because pH + pOH for a certain solution at STP must always equal 14,00. Also the Ka of a base can be calculated by means of the following formula:

pKa + pKb = pKw (which means KaKb = Kw = 10^-14)

pKw = 14 (which means Kw = 10^-14)

10^-pKa = Ka

To calculate the pH of a solution of a weak acid (or the pH of the solution of a weak base) the following formulae can be applied:

This is because only a small portion of the weak acid will dissociate. You can calculate by calling [X^-] and [H₃O⁺] x and filling out the equation. Some simple calculus should yield the [H₃O⁺]. It is useful to have a calculator handy in that case because it will be rather laborious to calculate x.

E.g. a 0,1M solution of formic acid (HCOOH)

pKa = 3,79 è Ka = 1,6 * 10^-4

[HCOOH] = 0.1

Fill out the formula, do the math and the answer will prove to be 3.9 * 10⁻³. (hence the pH will be ~2,4)

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 H3COOH	Commonly known as vinegar, 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 available 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 HNO3	Not commonly available, which is a shame considering how useful it is. Nitrates of metals are all soluble so it provides a good ability to solvate a cation of your choice. It is a strong oxidizing acid, able to oxidize metals readily at room temperature evolving nitrogen oxides.
Sulfuric Acid H2SO4	The staple acid of at home chemistry. Very difficult to obtain in some countries but relatively easy to find in America. Most sulfates are soluble in water although there are some notable exceptions (CaSO4, BaSO4, and PbSO4). Concentrations vary widely form common battery acid (~30%) to additional acid anhydride dissolve in 100% H2SO4. Weakly oxidizing.
Sulfamic Acid NH2SO3H	Somewhat readily available and stronger then many organic acids. It forms may highly soluble salts, in the pure form it is a solid.
Boric Acid B(OH)3	Very weak acid. Borates are readily available in the cleaning industry, dehydrates easily to boric oxide. Boric acid readily forms boric esters, which burn to give beautiful colors.
Hydrofluoric Acid HF	Weak solutions are available over the counter for cleaning rims of cars and such (~3%). Hydrofluoric acid is very toxic and highly concentrated solutions can kill very rapidly if splashed on the skin. As an acid though it is somewhat weak compared to hydrochloric, as a pure compound it is a liquid near room temperature.
Cyanuric Acid HOCHC(OH)NC(OH)N*2H2O	Somewhat weak acid available for adjusting the pH of pools.
Phosphoric Acid H3PO4	Concentrated phosphoric acid is a fairly strong acid. It readily attacks metals forming phosphates, which are on the whole soluble in water. The more concentrated the solution the more syrupy it is until it becomes a solid. Phosphoric acid will not boil, it will continuously loose water even past where it is 100% dehydrating to other forms of phosphoric acid such as pyrophosphoric acid. Avalible for cleaning metal and for marine cleaning.

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 standard 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, 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 reduced 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. [Note, a picture of this reaction is shown in the picture in the preceding section]

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 

        The case to the left 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, usually has to be bought from a pyrotechnics supplier or made via electrolysis.
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.  Occasionally available during the summer months as fertilizer.
Nitric Acid  HNO3	Clear - Yellow/Green liquid.  Available in various concentrations, >70% show remarkable oxidizing capabilities, lower concentrations available over the counter for hydroponics.
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:

With molten salt oxidations one can force metals into oxidation states that would be very difficult to achieve in the aqueous phase and would be considerably less stable if formed in that way as well. The actual chemistry of such oxidations is usually complex but there are only two simple needs to perform most of these oxidations, an alkali metal hydroxide, and an oxidizing agent usually an alkali metal nitrate where the gaseous visages of the oxidizing anion might readily leave the melt and the remaining cation will not interfere. Several reactions go on in such melts, but mixtures involving potassium hydroxide make a good example:

2KOH Û H2O + K₂O

4KOH + 3O₂ Û 4KO₂ + 2H₂O

Now, peroxides and superoxides are strong oxidizing agents in their own right, but the oxide anion O^2- is incredibly basic as is the superoxide O^-1/2 (note that potassium peroxide is fairly unstable, and does not exist appreciably in the molten state). The oxidizing agent in the melt, usually something like potassium nitrate helps to drive this equilibrium, acting as a very convenient source of oxygen. At these temperatures things like potassium nitrate are very reactive, a cotton glove for instance, coming into contact with molten KNO₃ will burst into flames, but at room temperature KNO₃ could be safely handled with ones bare hands. All in all, the very basic electron rich environment can best stabilize a number of high oxidation state compounds, which can then be used in further chemical endeavors.

The melts used for these oxidations are fairly corrosive; vessels of nickel, platinum, and silver are best. Glass is out of the question, as the strong bases present will attack it. However for the roughest oxidations, disposable vessels of steel or commonly available pipefittings can work. These will contaminate products obtained but compounds formed under these conditions are not going to be very pure anyways. Here are a few examples of high oxidation state compounds that can be made by these methods:

Ferrate [FeO₄]^2- : Ferrates will decompose almost instantly if in acid solution, quickly in neutral solutions, and slower in basic solutions. Kept free of moisture and stored without access to air, ferrates will keep for several weeks or months. They are made by fusing ferric oxide with KOH and an oxidizing agent and are purple/red in color. Ferrates can be precipitated from an aqueous solution as the slightly soluble barium salt or by concentrated potassium hydroxide solution.

Bismuthate [BiO₃]^- : Bismuthate as with many other high power oxidizing agents is available from chemical suppliers usually only as a purity of 85% or so, further refinement being unnecessary due to some of the brute force type oxidations done with it. Out of these four this is the second most commercially available oxidizer listed. It can oxidize manganese ions in solution to permanganate and is usually found for sale as the sodium salt. It can be prepared by fusing bismuth trioxide with potassium hydroxide as long as the mix is exposed to air.

Chromate [CrO₄]^2- : This is the most widely available oxidizer listed here. Chromates are toxic and should be handled with care. They are also the weakest oxidizer on this list. They are usually formed by fusing chromium (III) oxide, acidification of a solution of chromate will lead to the formation of dichromate which will usually precipitate if the concentration is high enough and the temperature lowered afterward, dichromates being more useful then chromates. Industrially this process is used to make dichromate by fusing chromite ore (FeCr₂O₄) with potassium hydroxide in the presence of oxygen, after oxidation the mixture is dissolved in water and acidified, the chromate being converted to dichromate and the ferrate going to soluble Fe³⁺.

Manganate [MnO₄]^2- : Manganates are green in color and the product of fusing manganese dioxide. They are fairly unstable and upon addition to water and acidification yield a solution of permanganate. Subsequent filtering and crystallization allowing for the production of permanganate at home, which is useful as well.

When an oxidation is completed there are two courses of action, if the product is stable to the atmosphere it can be poured onto a sheet of steel and allowed to cool quickly, then broken with a hammer and stored. However if it is not then it must be covered in the crucible while covered and once cooled chipped out and stored. In either case these melts are very strong oxidizing agents and must never come into contact with anything organic or flammable. In addition these melts can NEVER be poured directly into water while in the molten state as they can very often explode. Temperature control is not a major issue with most molten oxidations but things should still never be heated too strongly, molten oxidations work best in the range from 375 – 550 °C and the amount of time to hold the reactants there depends strongly on what you are trying to oxidize and the amount you have in the mixture. If you desire to produce an oxidizing agent that is very unstable at high temperatures considering an eutectic mixture can help greatly, a 50/50 mixture of NaOH/KOH has a significantly lower melting point then either component alone.

Aside from the inherent risks of holding oxidizing mixtures at high temperatures, the use of nitrates can lead to the formation of nitrogen oxides which are extreme hazards therefore the actual heating step should be preformed while you are not in the company of the reaction vessel. Additionally the subsequent dissolution and acidification of some of these mixtures can also lead to nitrogen oxide release if the reaction yielded a large amount of nitrites as can often be the case. Not to mention that the oxidation reactions also have the possibility of going awry, if excessive frothing or sparks start to come from a reaction mixture that is your key to exit the area. One final note, chlorates and perchlorates can be used for these oxidation reactions, however acidification of a solution of chlorate can yield explosive quantities of chlorine dioxide and additionally the reaction itself can run away in the presence of certain metal oxides.

4.5 Reducing Agents

The most common class of reducing agents one runs across are usually the active metals. In pyrotechnics aluminum, magnesium, and occasionally zinc are used with strong oxidizing agents such as perchlorates and nitrates to give spectacular exothermic reactions. However on a more controlled level these metals can also be used to give reliable reactions even in reactions involving aqueous reactants. Additionally there are a number of organic reducing agents which are most popular in the field of organic chemicstry. Both organic and strictly inorganic reducing agents are of great utility in the chemistry lab however they are more difficult to obtain then oxidizing agents in most cases and in the case of the metals, usually difficult to get into a workable form.

Reducing Agent	Source	Example of Use
Aluminum Al	Pyro suppliers, scrap yards, foil	Powdered aluminum is a powerful solid state reducing agent (Thermite reactions), it can also work to reduce cations in the aqueous phase and can be amalgamated with mercury for organic reductions.
Magnesium Mg	Scrap yards, camping suppliers (fire starter), pyro suppliers, cell phone pieces and in some other high end applications (bike frames etc.)	Powdered magnesium is a stronger reducing agent then aluminum, it reacts slowly with water reducing it. Magnesium can reduce a number of inorganic compounds such as NaOH works in thermite type reactions.
Hydrogen H2	Reaction of a strong mineral acid with an active metal (aluminum, magnesium, iron)	Reductions with hydrogen gas usually take place at elevated temperatures when concerning inorganics and under high pressures with organics, it is very useful but difficult for an amateur to use.
Lithium Li	Purchased from a chemical supplier, some batteries contain lithium, home electrolysis in non-aqueous medium or of electrolysis of lithium chloride/bromide eutectic.	Lithium metal finds use either alone or in compounds for the reductions of organic compounds, lithium dissolved in ammonia or n-butyl lithium being incredibly strong organic reducers.
Sodium Na Potassium K	Purchased from chemical supplier, made at home via electrolysis or reduction of salts.	Considering the high reactivity for each of these they find little practical use in but are occasionally called upon. A liquid eutectic is formed between these two elements with is unbelievably reactive.
Carbon C	Graphite, sugar carbon (formed by heating sugar till it decomposes), coal, charcoal	Carbon is an excellent reductant however its use requires very high (>900C) temperatures.
Sulfite SO3^2-	Some OTC sources, bubbling sulfur dioxide into basified water.	Sulfites are oxidized to sulfates and in the process function as weak reducing agents.
Citric (Ascorbic) acid C6H8O7*H2O	Vitamin C tablets, sold as citric acid for flavoring, extracted from lemon juice	Weak aqueous reducing agent, works better at higher temperatures, good for making metal powders of somewhat nobel metals (e.g., nickel, chromium, silver, etc.)

Runaways can happen with reductions as well as oxidations.
 

 
Chemistry can be fun and educating in many ways but one should always remember to be careful. Sometimes one just forgets to do the background work first, before moving to do the actual experiment.. That happened to me some time ago.
 
I had previously prepared few moles of nitrotoluene and planned to reduce some of it to toluidines. So, I remembered the standard Sn and Fe reductions of aromatic nitro compounds with HCl and did a few calculations on the amounts of reactants required. Ended up using 1mol of nitrotoluene and the appropriate amounts of 40 micron, hydrogen reduced Fe powder and 37% aqueous HCl. The nitrotoluene was mixed well with the Fe powder and a little distilled water in a 500ml erlenmeyer flask. I was first going to use a magnetic stirrer to efficiently keep the iron powder suspended but remembered that iron is magnetic, luckily before I dumped in the stirbar.. HCl was put into an addition funnel and was added drop by drop to the mixture, which I swirled continuously. I expected a fast temperature rise but there was none. HCl addition was continued a bit faster and then the temp finally started to rise. Stopped the addition for a moment and swirled strongly. After continuing the addition the temp didn't rise much more so I decided that I’d add half of the remaining HCl now and the rest after ten minutes or so. That was big mistake. Soon the temp started to rise fast. I swirled the flask as strongly as I could but it didn't seem to help. I took the thermometer off as it was nearing the limit (100C). I put the flask in to an ice bath and swirled vigorously. The flask started to feel _very_ hot at that time and I was ready to dump it in the bath, but I was too late. Suddenly the mix started to boil and shoot itself out of the flask. I had to let go off the flask as it was so hot and then the reaction got so vigorous that it shot all the remaining liquid out in a geyser like fashion on the floor. I quickly took a bottle of water and poured it in and over the flask. After I had the reaction tamed I decided to go outside and take the annoying gasmask off for a while. I immediately smelled the slightly irritating smell of nitrotoluene and realised that all the reaction mixture had probably boiled out of the flask with the steam.. The whole neighbourhood smelled of nitrotoluene for a few hours, as it wasn't windy. The flask seemed ruined but I got it eventually cleaned with some HCl and sodium ethoxide solution. The worst part was cleaning the mixture off the coarse concrete floor.. And the smell stayed for ages in my lab.

That certainly thought me a lesson to always find about different reactions and the possible mistakes that could be made during it. And not to start doing a new reaction with that much material. Starting from mill moles is much more recommended to get a feel for the reaction, before scaling up!

4.6 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.

The main use of desiccants is to remove water, usually from a liquid or gas to render that liquid or gas largely free of water in reactions where water might inhibit a desired reaction, interfere with a reaction, or cause extraneous byproducts. Drying agents can form very strong bonds to water, strong enough to take water from a chemical bond, such as the reaction between concentrated sulfuric acid and sugar, where the sulfuric acid will remove the water from the molecule C₆H₁₂O₆ leaving behind just carbon in a very pleasing visual display, phosphorus pentoxide, and hot NaOH are the only other drying agents on this list that posses this strength of drying power, agents of this type are referred to as dehydrating agents.

Chemical Name /Formula /Formula of hydrate	Form of anhydride / Form of hydrate	Details on Agent
Sulfuric Acid  H2SO4 H2SO4*xH2O	Dehydrating Acid Heavy 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 remains liquid, will dehydrate sugar to carbon. Great for drying liquids.
Phosphorus Pentoxide P2O5 H3PO4 or H3PO3	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 MgSO4*7H2O	Drying Agent Solid Forms stable hydrate	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 CaCl2*2H2O (But will go further)	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 CuSO4*5H2O	Drying Agent Solid (Colorless) Green/Blue when hydrated	Widely available for killing roots in sewer lines or preparable by dissolving copper in hot sulfuric acid, it can be made anhydrous by heating. 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.
Calcium Sulfate CaSO4 CaSO4*2H2O	Drying Agent Solid White Forms a stable hydrate	Available over the counter as the semi-hydrated ‘Plaster of Paris’ or as drywall, which is the dihydrate. Made anhydrous by heating is possesses decent drying abilities but its ready availability and low cost make it somewhat desirable, color changing versions are available from chemical supply companies.
Magnesium Perchlorate Mg(ClO4)2 Mg(ClO4)2*6H2O	Drying Agent Solid but liquid when hydrated	One of the kings of the drying agent world, magnesium perchlorate possesses exceptional drying ability. However it can explode when exposed to solvent vapors or intense heat and therefore it has fallen into disuse, concentrated sulfuric acid or phosphorus pentoxide often substituted for it.
Sodium Hydroxide  NaOH NaOH*xH2O	Dehydrating Agent Solid but liquid when hydrated	Available over the counter, store bought NaOH contains some impurities and is of a variable composition of sodium oxide and water. It is a good drying agent for taking the last bit of water out of liquids in which it is insoluble but does not have the ability to dry large amounts well, just small amounts of a liquid/solid/gas thoroughly.
Calcium Oxide CaO Ca(OH)2	Drying Agent Solid forming a stable hydrate	Widely available for adjusting the pH of soil, the calcium hydroxide thus formed is only slightly soluble, although regenerateable through heating it is more often then not simply used once.

The activity of a dehydrating agent, the ability of it to pull water from its surroundings is not usually something to be directly gauged, but there is a great difference in each drying agents ability to pull water and trap it. For example, sulfuric acid when concentrated will char wood and turn sugar to coal. Whereas calcium chloride will do neither. These aspects can make a great difference in their usage and what reaction they may be unsuited for.

4.7 Poisonous Reagents 

A poisonous reagent might be easily classified as a chemical that requires only minimal unintentional contact to cause adverse effects. Such a definition is better suited for this class then simply a chemical that can cause harm, after all, there is a lethal dose for table salt and alcohol, so it is better to only classify those chemicals that could easily cause harm to oneself thought an accident as poisonous. From here poisonous chemicals are further divided into two categories, not independent from one another. Those chemicals that are cumulative poisons, and those that are not.

A cumulative poison is a poison whose presence in the body is not immediately eliminated and it accumulates in the system, i.e., it would be easy for a person to take in more of this poison, however infrequently, then their body will expel. This can be referred to as the half-life of a substance. Examples of cumulative poisons are lead salts, fluoride, radioactive strontium, and others. Cumulative poisons can also have almost no half life in the body, but are instead cumulative in the effects they cause, long term lung damage can result from inhalation of even minute amounts of some chemicals and upon repeated exposures that damage might become severe enough to cause emphysema or other conditions.

One mistake people take into account when handing a potentially poisonous substance and assessing its lethality is to consider the time frame over which it is lethal. Poisons that can kill in minutes such as hydrogen sulfide and hydrogen cyanide are often viewed with considerable more trepidation then other gasses like nitrogen dioxide, simply because nitrogen dioxide may not kill instantly. It still caries with it significant danger, all three of them do, and it is not a matter of will it kill you instantly or eight hours from now, it’s a matter of if the chemical you are dealing with is dangerous and taking the necessary precautions to ensure that should an accident happen the least harm will befall you. A poison is a poison, and aside from physical differences they should all be treated with the same careful consideration.

Let’s say, for example you have dissolved silver in excess nitric acid and currently have the beaker sitting in the middle of an open table. Now, you have chosen to precipitate the silver and simultaneously neutralize the excess nitric acid with sodium carbonate. You make your carbonate into an aqueous solution and add it drop wise slowly to avoid excess spattering. The next day you wake up from a restful sleep to find that your arms are covered with tiny black dots, and so is your face. Invisible drops of silver chloride solution were thrown from the beaker, carried by the wind, and otherwise deposited on your person. Had that been a highly toxic chemical you actually may have been beginning to feel the effects, a similar neutralization of a barium salt solution with excess carbonate may well make you sick within a few hours.

Always pay careful attention to chemicals that you work with that may be poisonous. And do not use them unless you feel you have to. Vapors can travel surprisingly far, heating poisonous solids may release similar vapors, some reactions may cause the breakdown of chemicals into more poisonous alternatives. The best ways to deal with a poisonous substance depends on its current form.

My poisonous substance is a liquid or is in solution:

If your substance is in water, its poisonous ability is slightly lessened, water has a weak ability to penetrate, chemicals that do however are lipophilic. Examples are carbon tetrachloride, DMSO, ethanol, ether, poisons solvated in these pose a greater hazard then those chemicals alone usually. However, should your substance be a liquid that is comprised of heavy metals directly bonded to carbon, organo-metiallic compounds, I cannot stress enough the danger involved, these will give heavy metals the most direct path into your body and straight to your brain. Wear gloves and don’t cause the solution to foam, avoid heating of solutions that contain liquid poisons unless they are completely enclosed within a glassware setup.

My poisonous substance is a solid:

As l