Chemistry 103A; Sections 5, 6, 7, 8; Lecture 9; 11 Sep 00

Chemistry 103A; Sections 5, 6, 7, 8; Lecture 9, 11 Sep 00

Electrical conductivity

If you can get enough ions in a water solution the solution will conduct electricity.

Ionic compounds whose aqueous solutions conduct electric are called electrolytes.

We distinguish between strong electrolytes and weak electrolytes

Strong electrolytes are compounds which produce a lot of ions in water solution. (All of the compounds shown above are strong electrolytes. They are "completely dissociated" into ions in aqueous solution.)
Weak electrolytes are compounds which produce a relatively small number of ions. An example of a weak electrolyte is acetic acid.
CH₃COOH(aq) ® H⁺(aq) + CH₃COO^- (aq)
For reasons which we will discuss later, acetic acid solutions in water do not produce very many ions. (You will even learn to calculate how many ions are produced.)
Compounds that do not conduct electricity in aqueous solution are called nonelectrolytes. These compounds do not give any ions at all in water solutions. Examples would be sugar, ethylene glycol (antifreeze), alcohol, etc.

Notice that in all cases where we get ions we get both positively charged ions and negatively charged ions. And we get them in numbers so that the total charge in the solutions is zero. That is, the solution is electrically neutral. The positive ions are called cations (because it was discovered that they were attracted to the cathode or the negative electrode).

The negative ions are called anions (because they are attracted to the anode).

Reactions in aqueous solutions

Consider the reaction of HCl and NaOH in water solution:

We can write this reaction as

HCl(aq) + NaOH(aq) ® NaCl(aq) + H₂O(l).

But since all of these compounds (except water) dissolved in water to form ions, we could - or should - write:

H⁺(aq) + Cl- (aq) + Na⁺(aq) + OH^- (aq)

® H₂O(l) + Na⁺(aq) + Cl^- (aq)

Notice that the interesting part is:

H⁺(aq) + OH^- (aq) ® H₂O(l)

The Cl^- (aq) and Na⁺(aq) occur on both sides of the equation. They can be canceled out. We call the Cl^- (aq) and Na⁺(aq) ions spectator ions. It is very common to just leave the spectator ions out when writing reactions in aqueous solution. (The spectator ions have to be there in order to keep the solutions electrically neutral, but their identity is not usually important to the chemistry taking place.) Another example:

2 HCl(aq) + Na₂CO₃(aq) ® CO₂(g) + H₂O(l) + 2 NaCl(aq)

This is short for

2 H⁺(aq) + 2 Cl^- (aq)+2 Na⁺(aq) + CO₃^2- (aq)

® CO₂(g) + H₂O(l) + 2 Na⁺(aq) + 2 Cl^- (aq)

When we strip out the spectator ions we get

2 H⁺(aq) + CO₃^2- (aq) ® CO₂(g) + H₂O(l)

This is the interesting (or important) part of the reaction.

The reaction that is left after we drop the spectator ions is called the net ionic reaction or sometimes just the net reaction.

There must always be spectator ions in an ionic reaction, but we don't always have to know what they are.

How do we know which ions to drop and which to keep?

Drop:

aqueous ions that appear on both sides

Keep:

insoluble ionic compounds (like AgCl)
gases (like CO₂)
water
nonionic compounds

How do we know which ionic compounds are insoluble?

There is a set of general solubility rules for ionic compounds in water on page 184 of the text.

Na⁺, K⁺, NH₄⁺ compounds all soluble
NO₃^- compounds all soluble
Cl^- , SO₄^2- compounds soluble except AgCl, BaSO₄, PbSO₄
CO₃^2- , PO₄^3- , S^2- , OH^- , O^2- compounds insoluble except Ba(OH)₂ (and, of course, compounds with Na⁺, K⁺, and NH₄⁺).

Many times we can reason by analogy from these rules. For example AgCl is insoluble so we might expect that AgBr and AgI would be insoluble too.

Let's go back, now, and look at the examples we gave before:

H⁺(aq) + Cl^- (aq) + Na⁺(aq) + OH^- (aq)

® H₂O(l) + Na⁺(aq) + Cl^- (aq),

2 H⁺(aq) + 2 Cl^- (aq)+2 Na⁺(aq) + CO₃^2- (aq)

® CO₂(g) + H₂O(l) + 2 Na⁺(aq) + 2 Cl^- (aq),

and another one we didn't see before:

AgNO₃(aq) + NaBr(aq) ® AgBr( ? ) + NaNO₃(aq)

Acids and Bases (in water solution)

There are three different schemes for defining acids and bases in chemistry which you will learn about in this course. These schemes get progressively more general (and more abstract), but they are all very useful in discussing chemical reactions.

For now we will only talk about the simplest definitions of acids and bases. this is the Arrhenius definitions of acids and bases.

In the Arrhenius scheme:

In Arrhenius acid is any compound that gives H⁺ ions in water solution.
An Arrhenius base is any compound that gives OH^- ions in water solution.

We will just call them simply acids and bases. You won't need to distinguish Arrhenius acids and bases from the other classifications until Chapter 17.

Examples of acids are HCl, HBr, HNO₃, H₂SO₄, HClO₄, etc. All of these compounds dissociate in water to give H⁺ ions.
Examples of bases are NaOH, Ba(OH)₂, NH₃, and etc. They all give OH- ions in water. (The last one looks weird, because we don't see any way to get OH^- ions out of it. The way it works is that the NH₃ reacts with H₂O to give NH₄OH, which then gives us the OH^- ions in solution.)

Acids and bases are frequently formed by dissolving oxides or hydroxides in water.

Metal oxides and hydroxides give bases when they dissolve in water.

Na₂O + H₂O ® 2 NaOH
Ca(OH)₂(aq) ® Ca²⁺(aq) + 2 OH^- (aq)

Nonmetal oxides give acids in water solutions.

SO₃ + H₂O ® H₂SO₄
N₂O₅ + H₂O ® 2 HNO₃

(Of course, there are other sources of acids than nonmetal oxides. Compounds like, HCl, HBr, and so on, are gases which dissolve in water to give strong acids (strong electrolytes - lots of ions).

The smelly gas, H₂S, dissolves in water to give an acid, but it is a weak acid (weak electrolyte).

Classes of Reactions in Aqueous Solutions

There are many possible chemical reactions. Fortunately, we can find ways to classify reactions by features they have in common. One class of reactions is called exchange reactions.

Consider the reaction:

NaCl(aq) + KNO₃(aq) ® KCl(aq) + NaNO₃(aq).

Notice that the ions have swapped partners. That is what we mean by an exchange reaction. However, this particular reaction has a problem. If we write it in terms of aqueous ions we can see what the problem is.

Na⁺(aq) + Cl^- (aq) + K⁺(aq) + NO₃^- (aq)

® K⁺(aq) + Cl^- (aq) + Na⁺(aq) + NO₃^- (aq) If we now look to strip out the spectator ions we see that all of the ions are spectator ions. So this reactions is not a reaction at all, it is just a mixture of ions.

The way to force a reaction to take place is to have one of the compounds removed from solution. This will happen if one of the products is an insoluble compound, or water, or a gas.

Precipitate formation:

BaCl₂(aq) + Na₂CO₃(aq) ® BaCO₃(s) + 2 NaCl(aq)

If we recognize that everything except the solid barium carbonate is a soluble ionic compound, and strip out the spectator ions we get the net ionic reaction:

Ba²⁺(aq) + CO₃^2- (aq) ® BaCO₃(s)

When we mix aqueous barium chloride and aqueous sodium carbonate we get a precipitate of barium carbonate. The other ions are just along to maintain electrical neutrality.

Another example of a reaction forming a precipitate:

2 AgNO₃(aq) + K₂S(aq) ® Ag₂S(s) + 2 KNO₃(aq).

Water formation

Acid-base reactions always form water (in the Arrhenius picture of acids and bases). For example:

HClO₄(aq) + NaOH(aq) ® NaClO₄(aq) + H₂O(l).

Stripping out the spectator ions from this reaction leaves:

H⁺(aq) + OH^- (aq) ® H₂O(l),

which is the same for all acid base reactions.

Gas Formation

We have already seen some reactions which form CO₂. For example,

K₂CO₃(aq) + H₂SO₄(aq) ® K₂SO₄(aq) + CO₂(g) + H₂(l).

Notice that the net ionic reaction is the same as our previous reaction giving CO₂. It doesn’t matter if the CO₃^2- came in with Na⁺ ions or with K⁺ ions

Let’s look at a different reaction.

Na₂S(aq) + 2 HBr(aq) ® H₂S(g) + 2 NaBr(aq).

What is the net ionic reaction?

Oxidation-Reduction Reactions (Redox Reactions)

There is another very important class of reactions where we generally strip out the spectator ions and just look at the net ionic reaction.

These are oxidation-reduction reactions.

Before we can talk about these we have to define what we mean by "oxidation" and by "reduction." We also have to define "oxidation number" or "oxidation state."

The concept of oxidation originally came from substances reacting with oxygen (well, duh!). Things reacting with oxygen were said to be oxidized.

For example, iron will react with oxygen to give iron (III) oxide.

4 Fe(s) + 3 O₂(g) ® 2 Fe₂O₃(s).

The iron, of course is oxidized.

Elemental iron can be recovered from the Fe₂O₃ by an appropriate chemical reaction,

Fe₂O₃(s) + 3 H₂(g) ® 2 Fe(s) + H₂O(l).

We say that the iron in the Fe₂O₃ has been "reduced."

It was soon discovered that you don’t necessarily need oxygen to produce the same change in the iron. The oxygen converted elemental iron into Fe₃⁺ ions, but this can be done in other ways without oxygen. For example,

2 Fe(s) + 3 Cl₂(g) ® 2 FeCl₃(s)

also produces Fe³⁺ ions.

In order to define oxidation and reduction accurately we must first define the terms oxidation number or oxidation state. (I will use these two terms interchangeably.)

The easiest way to define oxidation number is to think of it as the hypothetical ionic charge on the element.

If the element has formed a monatomic ion then the oxidation state IS the ionic charge.

However, the element may be incorporated into a molecule or molecular ion and not really have an ionic charge of its own. Nevertheless, we look at the element as ask what its ionic charge would be if it were an ion.

Here are some examples:

An element in its elemental form has an oxidation number of ZERO.
O₂, N₂, Fe, P₄, S₈, Cl₂, etc. all have an oxidation number zero.
Monatomic ions have an oxidation number equal to their ionic charge.
Cl^- , O^2- , Na⁺, Ca²⁺, Sn⁴⁺, H⁺, etc.

When an element is incorporated into a molecule or a molecular ion we have to figure out its oxidation state from the things it is bonded to.
Oxygen in a compound is almost always - 2 hydrogen in a compound is almost always +1

Here are some examples which we will work out in class:

HNO₃

SO₄^2-

PO₄^3-

Cr₂O₇^2-

B₂O₃.

Now we can define what we mean by oxidation and reduction.

Oxidation is an increase in oxidation number.
Oxidation always produces a loss of electrons.
Other tell-tale signs,
Gain of one or more O atoms
Loss of one or more H atoms

Reduction is a decrease in oxidation number.
Reduction always gives a gain of electrons.
Other tell-tale signs,
Loss of one or more O atoms.
Gain of one or more H atoms.