Week of November 16, 2009

This experiment introduces a variety of chemical reactions which can occur in aqueous solutions. It should help develop a sense of when a chemical reaction is going to occur, and what some classic chemical reactions look, feel and smell like.

It is important that you record your observations in your notebook in sufficient detail. Your observations will give you a clue to the chemistry of a particular reaction. The balanced chemical reaction that you are asked to submit for each of your reactions should be consistent with your observations in the lab. If they are not, you may want to take a step back and ask: Why? If you cannot answer this question, it may be in your best interest to talk to your lab instructor, lecture professor, T.A. and/or tutor.

Reactions #10 requires some time for the chemical reaction to occur. In the interest of efficiency, you should set this reaction up at the beginning of the lab period, and observe the reaction periodically (and, of course, record your observations in your notebook).

Good Luck and don't save this lab write up until the night before it's due. It's a long one!!!

Helpful Hints for balancing the chemical equations of the reactions you did for experiment 10-003c.

metathesis: a reaction that proceeds without any of the reagents changing oxidation states, and

oxidation-reduction (or redox) reactions, that proceed with changes in oxidation states.

The challenge becomes to learn just what conditions will lead to a legitimate metathesis. For our purposes, we can say that there are three events that will cause a metathesis to occur:

This classic acid-base reaction is probably familiar, especially in its most common guise of a metal hydroxide plus an aqueous acid to give a metal salt and water: ("Acid plus base gives salt plus water".)

Metal oxides dissolve in water to generate basic solutions note the formation of OH-(aq) ions in reactions:

Many oxides of non-metals (such as nitrogen and sulfur) form acidic solutions (note the generation of H+(aq) ions) when they dissolve in aqueous solution:

Acid-base reactions are one of the most common reaction types in all of chemistry. As your study of chemistry continues, your model of what constitutes an "acid" and a "base" will expand, and you will come across many acid-base reactions that do not involve water. For now, however, it is convenient to consider acid-base reactions to be a type of metathesis in which H₂O_(l) is generated from the reaction of H⁺_(aq) and OH^-_(aq).

Let's look at an example that you will see in lab. If we mix the aqueous NaCl with an aqueous solution of AgNO₃, the mixture becomes opaque as it turns milky white. If you let this solution sit for a while (or accelerate the effect of gravity by using a centrifuge), the white solid will settle out to the bottom of the reaction flask; such a solid, that forms as the result of a reaction between two solutions, is call a "precipitate". The solid is AgCl(s), so the reaction must be:

Na^⁺(aq) + Cl^{^-} (aq) + Ag^⁺(aq) + NO₃^{^-}(aq)---> AgCl(s) + Na^⁺(aq) + NO₃^{^-}(aq)

The problem, of course, is to know which salts are soluble and which are insoluble. Unfortunately, there really is no way to figure this out from first principles, you just have to memorize which salts are relatively insoluble in water. Most general chemistry texts give rules-of-thumb which can help you, so you should check out your text. In brief, it is good to remember:

At first, this may seem to be a daunting list of rules, but it really isn't. It is well worth remembering.

Soluble Salts

1. All nitrate (NO₃^{^-}) and acetate (CH₃CO₂^{^-}) salts are soluble.

2. All alkali metal (Li^⁺, Na^⁺, K^⁺ etc.) and ammonium (NH₄^⁺) salts are soluble.

3. Acids (H^⁺ salts ?) are soluble.

4. Halide (Cl^{^-}, Br^{^-}and I^{^-}) salts are soluble. EXCEPT when combined with Ag^⁺ , Pb^²⁺, and Hg₂²⁺ cations. (thus the insolubility of AgCl, as discussed above).

5. Sulfate (SO₄2-) salts are soluble. EXCEPT when combined with Pb²⁺, Sr²⁺, Ba²⁺, and Hg₂²⁺. (Ag₂SO₄ and CaSO₄ are only "slightly soluble".)

Insoluble Salts

6. Sulfide (S₂^-) salts are virtually all insoluble. EXCEPT when combined with Group IIA, or IA metals or NH₄⁺ (see rules 2,3,4).

7. Hydroxides are insoluble, except for those in rules 2, 3 and 4. (Ca(OH)₂, Sr(OH)₂, and Ba(OH)₂ are "slightly soluble".)

8. Phosphates (PO₄^{^3-}), carbonates (CO₃^{^2-}) and sulfites (SO₃^{^2-}) are insoluble, except for those in rules 2, 3 and 4.

As noted above, metatheses can occur if mixing the electrolytes leads to the formation of a material which drops out of the aqueous solution (the precipitate). As we will discuss later in the course, it is this separation from the reaction solution that actually drives the reaction toward completion. If a relatively insoluble gas is generated during a metathesis, it will also leave the reaction area by bubbling away, and, just as with the formation of a precipitate, can drive a chemical reaction. There are several important reactions which illustrate this driving force.

When sulfide salts (e.g. Na₂S, CaS, MnS, CoS...) are exposed to H+(aq) ions, H2S gas is formed. Since H₂S(g) is not very soluble in water, it bubbles out of the solution, driving the reaction forward:

Consider the reaction of NH₄Cl with NaOH (both strong electrolytes). Both the Na⁺ and Cl^- will be spectator ions, and the net ionic equation for this reaction is reaction:

The generation of water is an immediate clue that this is an acid-base reaction, but it has the additional driving force that NH3(g) is not very soluble in water, and bubbles away from the reaction area.

This important class of reactions involves the release of a gas (CO₂) and an acid-base reaction. Consider, for example, the reaction of limestone (CaCO₃) with nitric acid (HNO₃):

Since nitric acid (HNO₃) is a strong electrolyte, the nitrate ion is a spectator ion and may be eliminated from both sides of the equation. The net ionic equation then becomes:

As written, this reaction is incomplete, however, as we don't know whether "H₂CO_3(aq)" actually exists. We do know that whatever is formed in the above reaction rapidly decomposes in several different ways; for this discussion we need consider only its decomposition to CO_2(g) and H₂O_(l)(reaction.

Combining these two reactions, the overall reaction between limestone and any acid becomes:

Two powerful driving forces make this reaction occur: the formation of water (it’s an acid-base reaction) and the release of CO_2(g) which escapes from the aqueous solution. Similar reactions are also observed for other carbonates and :

Metatheses occur when H₂O_(l) is formed by the reaction of H⁺(aq) and OH^-(aq) (an acid-base reaction), or when one of the products of the reaction leaves the reaction area (the solution). This separation from the aqueous solution can occur when one of the products is an insoluble precipitate, or is a relatively insoluble gas.

More commonly called "redox" reactions, these oxidation-reductions are reactions in which electrons are transferred from one atom (or molecule) to another atom or molecule. The electron transfers result in a change in the oxidation number (oxidation state) we assign to some of the atoms involved in the reaction, so it is easy to identify a redox reaction. Just look for a change in the oxidation state of any atom in the reaction; if such a change occurs, it is a redox reaction.

This requires that you be able to assign oxidation states to elements within molecules and ions. In brief, the following rules may be used to determine the oxidation states of most elements:

Some rules for determining oxidation numbers

1a. Elements in their normal state have oxidation number of 0.

1b. The oxidation number of a monoatomic ion is the charge on the ion.

1c. The sum of the oxidation states of the elements within a molecule or ion equals its charge.

2a. Alkali metals (Li, Na, K, Rb, Cs) have an oxidation state of +1

2b. Alkaline earth metals (Be, Mg, Ca, Sr, Ba) have an oxidation state of +2

3. Hydrogen (H) has an oxidation state of +1.

4. Oxygen (O) has an oxidation state of -2.

5. Fluorine (F) has an oxidation state of -1.

6. The remaining halogens (Cl, Br and I) have oxidation states of -1.

Rule 1 (a, b and c) is always true; rule 2 is generally true, except when it conflicts with rule #1. (For example, rule 2 says that sodium has an oxidation state of +1; this is certainly true for the Na+ ion, but not for elemental sodium, which has an oxidation state of 0 (rule 1).

Rule #3 is true, except when it conflicts with rules 1 and/or 2. For example, NaH cannot have both +1 oxidation states for Na and H (that would conflict with rule 1c), so rule 2 has precedence over rule 3, and Na is assigned the oxidation state of +1; using rule 1c, the H is forced to an oxidation state of -1. A similar ranking gives oxygen an oxidation state of 1 in H₂O₂ (rather than the expected value of 2), as rule 3 holds sway over rule 4.

Being able to determine the oxidation state of an element in a molecule or ion is an essential skill for anyone studying chemistry. It is strongly recommended that you invest an hour or so to make sure you can do it with ease. Your text will have many examples and practice problems to help you; take advantage of them.

In a redox reaction, one of the reactants loses electrons; the reactant that loses electrons is said to have been oxidized, and its oxidation state will increase to a more positive value. In this process, another of the reactants will accept electrons; the act of accepting electrons is called reduction, and a molecule (or ion) that has absorbed electrons has been reduced. In contrast to oxidation, reduction causes a decrease in the value of the oxidation number. This transfer of electrons between molecules accounts for the other name for redox reactions: electron-transfers.

An important point to remember about oxidations and reductions: they must occur in concert they cannot occur separately. If one chemical loses its electrons (getting oxidized in the process), it donates them to another species which is reduced as a result. A reactant which causes another to be reduced is called a reducing agent, or reductant. Conversely, a chemical which accepts electrons (getting reduced in the process) must have ripped the electrons off of some other species, oxidizing it. Reagents which cause other chemicals to be oxidized are said to be oxidizing agents or oxidants. This all makes sense, except it can lead to the sometimes confusing business that oxidizing agents get reduced and reducing agents get oxidized. This is summarized below:

To help you identify common oxidants and reductants, check the following tables:

Oxidant	Usual Product	Element Reduced	Change in Oxidation States
O₂	H₂O	O	0 ---> -2
Cl₂, Br₂, I₂	Cl^-, Br^-, I^-	Cl, Br, I	0 ---> -1
Cr₂O₇^2-	Cr³⁺	Cr	+7 ---> +3
MnO₄^-	Mn²⁺	Mn	+7 ---> +2
S₂O₈^2-	SO₄^2-	S	+7 ---> +6
NO₃^-	NO	N	+5 --->+2
ClO^-	Cl^-	Cl	+1 --->-1
H₂O₂	H₂O	O	-1 --->-2
Ce⁴⁺	Ce³⁺	Ce	+4 --->+3
Ag⁺	Ag	Ag	+1 ---> 0
Fe³⁺	Fe²⁺	Fe	+3 ---> +2
H₂O or H₃O⁺	H₂	H	+1 ---> 0

Reductant	Usual Product	Element Oxidized	Change in Oxidation States
Zn	Zn²⁺	Zn	0 ---> +2
Cu	Cu²⁺	Cu	0 ---> +2
Mg	Mg⁺	Mg	0 ---> +2
Fe	Fe²⁺, Fe³⁺	Fe	0 --->+2, +3
Cr	Cr³⁺	Cr	0 ---> +3
Na	Na⁺	Na	0 ---> +1
Fe²⁺	Fe³⁺	Fe	+2--->+3
Ti³⁺	Ti⁴⁺	Ti	+3 ---> +4
SO₂,(SO₃^2-^{^-})	SO₄^2-	S	+4 ---> +6
SO₄^2-	S₂O₈^2-	S	+6 ---> +8
I^-, Br^-,Cl^-	I₂, Br₂,Cl₂	I, Br, Cl	-1 ---> 0
H₂O₂	O₂	O	-1 ---> 0
C₂O₄^2-	CO₂	C	+3 ---> +4

Elements in high oxidation states (Mn(+7) in MnO₄^-, Cr(+6) in Cr₄O₇^-, Ce⁴⁺, N(+5) in NO₃^-, etc.) are generally oxidizing agents

Elements in low oxidation states, and especially metals in their elemental state (oxidation state of zero) are generally reducing agents.

Some species (H₂O₂ in particular) can behave as both a oxidant or as a reductant, depending on what it is reacted with. React it with a strong oxidant, and it will behave as a reductant; react it with a good reductant, and the H₂O₂ will behave as an oxidant.

The driving force for a redox reaction depends on the tendency of the oxidizing and reducing agents to react with each other. For now, we need only recognize that some oxidizing agents are very powerful, while others are significantly weaker.

One type of redox reaction is so important on our planet that it deserves a separate section. The terms "oxidant" or "oxidation" are derived from oxygen (O₂), the gas that allows us to live in our atmosphere. Outside of biochemistry, some of the most familiar reactions of O₂ involve reactions with compounds containing only carbon and hydrogen (hydrocarbons), or carbon along with hydrogen and oxygen in a 2:1 ratio (carbohydrates). We refer to such reactions as combustions. In a combustion the reducing agent (the hydrocarbon or the carbohydrate)is oxidized by O₂ to generate CO₂(g) and H₂O

a) the combustion of natural gas (CH₄): CH₄(g) + 2 O₂ (g) ---> CO₂(g) + 2 H₂O(g)

b) the combustion of octane (C₈H₁₈) in an automobile engine: 2 C₈H₁₈(g) + 17 O₂ (g) ---> 8 CO₂(g) + 18 H₂O(g)

c) a campfire (combustion of cellulose, C₆H₁₀O₅): C₆H₁₀O₅(s) + 3 O₂ (g) ---> 6 CO₂(g) + 5 H₂O(g)

As we all know, these combustions are rapid chemical reactions, and they release a great deal of heat. Oxygen gas will react with many other reducing agents, and it is easy to expand our definition of combustion to include these reactions. For example, magnesium and aluminum react with some vigor in atmospheric oxygen:

These reactions do not release CO₂(g) and H₂O, but they are rapid reactions, and release heat (light, noise, sparks etc.), so it is hardly a stretch to call these reactions "combustions". In the Falklands War (1982) between Britain and Argentina, the British lost several ships to missiles which ignited their aluminum hulls. The combustion of these aluminum ships was apparently more intense and furious than the burning of wooden warships in previous centuries.

Other reactions of oxygen are less spectacular, but also lead to the formation of metal oxides. The most important is the reaction of metallic iron with oxygen, in the presence of water: 2 Fe(g) + 3/2 O₂(g) + n H₂O ---> Fe₂O₃*n H₂O (s)

The product, a hydrated iron oxide, is also know as "rust"; this oxidation represents the corrosion of iron which attacks bridges, automobiles, buildings and anything else made of steel that is exposed to our atmosphere. Such corrosions do release heat, but the reactions usually take place so slowly that they are not generally referred to as combustion reactions.

molecular equation	net ionic equation
CaS(s) + 2 HCl(aq) ---> H₂S (g) + Ca²⁺(aq)+ 2 Cl^{^-}(aq)	CaS(s) + 2 H⁺(aq) --->H₂S (g) + Ca²⁺(aq)
Na₂S(s) + H₂SO₄(aq) --->H₂S (g) + 2 Na⁺(aq)+ SO₄^2-(aq)	Na₂S(s) + 2 H⁺(aq) ---> H₂S (g) + 2 Na⁺(aq)