Chapter 4: Aqueous Reactions and Solution Stoichiometry
4.1 Water
Water is by far the most common solvent on earth. It is sometimes called the “universal
solvent,” which of course it is not. If it were, what container could be used for storing it?
H2O is a bent molecule with a bond angle of about 105. Because it is not linear, the centers of
+ and – charge in a water molecule are located in different places, making the molecule polar.
Because of its polarity, water can dissolve a great many compounds (making aqueous solutions).
Water is best at dissolving ionic and other polar substances. Solubility refers to how much solute
will dissolved in a given amount of solvent. The solubility of some compounds in water is very
low (such as AgCl). Other compounds, like sugar, have very high solubility.
4.2 Electrolytes
Substances whose aqueous solutions conduct electricity are electrolytes. Electric current
involves the movement of charged particles (ANY charged particles, not just electrons). In
the case of electrolytes, ions in aqueous solution move in toward electrodes placed in the
solution, completing an electric circuit.
The degree of conductivity depends upon the number of dissolved ions in the solution.
Substances fall into 3 categories:
 Strong electrolytes: includes strong acids and bases, and all high solubility salts. These
substances ionize completely, producing lots of dissolved ions in the solution.
 Weak electrolytes: includes weak acids and bases and all low solubility salts. These
only ionize partially, or dissolve poorly, producing some but not many ions in solution.
 Non-electrolytes: includes non ionizing solutes, i.e. most covalent compounds. These
compounds, like sugars and alcohols do not ionize at all. No ions = no conductivity.
Be sure to examine the sample exercise 4.5 on p. 114 on identifying electrolytes.
4.3 Composition of Solutions
Solutions are homogeneous mixtures of a solvent and 1 or more solutes.
Many chemical reactions, such as those in your body cells, take place in water solution. To
perform stoichiometric calculations on these reactions, we must know 1) the nature of the
reaction (what type of reaction is happening) and 2) the amounts (or concentrations) of the
chemicals involved.
Chemists by far use molarity (M) most often as the preferred unit of concentration for chemical
reactions.
Molarity = M = moles of solute / Liters of solution
A 1.00 molar solution (1.00 M) contains one mole of dissolved solute in enough solvent to
produce a total volume of 1.00 L.
Ex: To make 1.00 L of a 0.150 M solution of sucrose in water:
1.00 L sln x 0.150 mol x 342 g = 51.3 g + enough water to bring total volume to 1 L.
1 L sln
1 mol
Pp. 134-6 have sample exercises for calculating molarity.
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Dilutions: It is very common for chemists to dilute concentrated solutions to lower
concentrations. For example, pure HCl is either a gas or a liquid under high pressure, very
dangerous to work with in either case for a normal lab situation. Instead, HCl comes in
concentrated water solution (12 M) from the manufacturers. So, anytime we wish to make up a
solution of HCl that is less than 12 M, we must dilute it from the 12 M stock solution. Such a
procedure requires the use of a volumetric flask and very accurate devices for measuring
volumes, like volumetric pipettes or burettes. A very simple but important mathematical
equation is used for calculating dilutions:
C1V1 = C2V2
Where C stands for concentration (usually molarity) and V stands for volume. The 1 is for the
beginning or “stock” solution, and the 2 represents the new diluted solution.
*When working with acids, it is a good idea to follow the AAA rule: "Always Add Acid to
water." The reason for this is that if water is poured into concentrated acid, the water can
actually get hot enough to boil and literally explode, spewing hot acid all over. NOT FUN!!
The equation often appears as M1V1 = M2V2 or MiVi = MfVf, but it can be used with any unit of
concentration.
4.4 Types of Chemical Reactions
**For purposes of the AP exam (and your year end final) reactions are generally divided
into 5 different categories:
 complexation
 redox
 acid/base
 precipitations
 organic/other
Of these, precipitations (a. k.a. ionic, metathesis, or double replacements), acid-base, and redox
reactions usually occur in solution.
4.5 Precipitation Reactions
Often when solutions of 2 ionic compounds (salts) are mixed together, a new +/- ionic
combination is produced that does NOT dissolve well in water. When this happens, solid ionic
crystals form which eventually fall to the bottom of the container and so the resulting solid is
called a precipitate.
The general solubility rules describe which ionic compounds are likely to be water soluble and
which are not. We will explore the rules in the laboratory. You should be intimately familiar
with the solubility rules.
Precipitation Reaction examples:
Rxns in solution that produce 1 or more low solubility products.
 HCl(aq) + AgNO3(aq)  HNO3(aq) + AgCl(s)
 Pb(NO3)2(aq) + Na2CrO4(aq)  PbCrO4(s) + 2NaNO3(aq)
Typical double exchange or double replacement rxns are sometimes called "metathesis" rxns.
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Metathesis Reactions (double replacement or ionic reactions)
Precipitations are common, but not the only type of ionic or double replacement reaction. These
have the general form: AX + BY  AY + BX.
In order for a rxn to occur, one of the products must be: a low solubility compound (in which
case a precipitate forms) or a covalent molecule, which may take the form of water, a weak
acid or a gas. Another way to look at it would be to watch for the formation of a weak or
nonelectrolyte, or a gas.
Acid-base neutralizations are metathesis rxns. So are precipitations.
Some metathesis rxns form gases. The most common ones to watch for are:
CO2, HCN, H2S, and SO2 or SO3. NH3 is also a gas, but is highly water soluble, so gas
production is often limited. Of these perhaps the most common is the rxn of an acid with any
carbonate cpd. which always produces carbon dioxide.
4.6 Describing Reactions in Solution
There are 3 ways of representing a chemical rxn.
 Formula (or molecular) equations show all complete chemical formulas intact. This is
not a great term, since some of the materials may be ionic cpds.
 Complete ionic equations show all soluble salts, bases and strong acids in their ionized
forms. Covalent substances including weak acids are shown as molecular formulas.
 Net ionic equations show all ions and molecules that undergo significant changes during
the course of the rxn. Spectator ions, which undergo no major change and appear the
same on the product side of the ionic equation as they did on the reactant side are
removed from the net equation.
Examples:
 Formula eq: HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l)
 Complete ionic eq: H+(aq) + Cl- (aq)+ Na+(aq)+ OH-(aq)  Na+(aq) + Cl- (aq) + H2O(l)
Spectator ions have been underlined.
 Net ionic eq: H+(aq) + OH-(aq)  H2O(l)
4.7 Stoichiometry of Precipitation Reactions
As in Chapter 3, the principles of stoichiometry apply to reactions in solution.
However, instead of dividing grams by molar mass to get moles, we calculate moles of a
compound in solution by multiplying volume by molarity.
V x M = moles
L x (mol/L) = mol
You may also see it this way: L x (mol L-1) = mol
The problem solving strategy for solution stoichiometry problems is as follows:
1. Identify the species present in the solution and which ones react.
2. Write the balanced net ionic equation for the reaction.
3. Calculate the moles of the reactants (V x M).
4. Determine, if necessary, which is the limiting reactant.
5. Calculate the moles of product(s) formed based on the limiting reactant.
6. Convert moles to grams or other units, as required.
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4.8 Acid-Base Reactions (Neutralizations)
Traditionally, acids produce H+ ions in solution:
 HCl(g) dissolves in water  HCl(aq)  H+(aq) + Cl-(aq)
 H2SO4 (aq)  H+(aq) + HSO4-(aq) and
 HSO4-(aq)  H+(aq) + SO4-2(aq)
Bases traditionally produce OH- (hydroxide) ions in solution:
 NaOH(s) in water  Na+(aq) + OH- (aq)
Ammonia is also a base, even though it has no OH- in its chemical formula:
 NH3(g) in water  NH3(aq) + H2O(l)  NH4+(aq) + OH- (aq)
Neutralization rxn: Acid + Base  a salt + water
(**All ionic cpds are considered salts.)
 HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l)
A more modern and useful definition of acids and bases is the Brønsted-Lowry definition.
Acids are proton (H+ ion) donors.
Bases are proton acceptors.
Solving acid-base solution stoichiometry problems is similar to precipitation problems. The
problem solving steps are essentially the same. See example problems in Chpt. 4.
Titrations: A special and common laboratory type of acid-base reaction is the titration. Its
purpose is to determine the unknown concentration of an acid or base. A solution of precisely
known concentration (the titrant) is added to the solution of unknown concentration (the
analyte). Titrant is added until the equivalence point (or the stoichiometric point) is reached, as
indicated by an indicator color change, or a rapid shift in pH.
In order for a titration to be successful, the following conditions must be met:
1. The exact reaction between acid and base must be known, and rapid.
2. The equivalence point must be accurately marked by the indicator.
3. The precise volumes of the analyte and titrant must be known.
The titrant concentration is typically determined through a procedure known as standardization.
The unknown analyte concentration can be titrated as described above using moles, volume and
the stoichiometry of the balanced equation.
Another efficient way to calculate the analyte concentration is to use a concentration known as
normality (N). Normality is defined as
N = chemical equivalents/L of solution
What, you may ask, is a chemical equivalent? An equivalent is 1 mole of a “particle of interest.”
For acids, the particle of interest is the hydrogen ion, H+.
For bases, the particle of interest is the OH- ion.
So, for an acid, an equivalent is the amount of acid needed to produce 1 mole of protons. For a
base, an equivalent is the amount of base needed to either produce 1) 1 mole of hydroxide ions,
or 2) accept 1 mole of protons.
Example: 0.3 M HCl = 0.3 N HCl. Hydrochloric acid is monoprotic (1 H ion/molecule) so
moles and equivalents are equal.
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0.3 M H2SO4 = 0.6 N H2SO4. Sulfuric acid is diprotic, so each mole of compound has 2
equivalents. Normality is 2 times the molarity.
It works the same for bases.
For NaOH, normality = molarity, but Mg(OH)2, normality = 2 x molarity.
Once the normality is known, a simple equation solves the titration problem:
NaVa = NbVb
N = normality, V = volume, a = acid, b = base. Simple.
Acids, Bases and Salts
Traditional acids are covalent cpds that produce H+(aq) ions in solution. The Brönsted-Lowery
definition of an acid is a "proton donor."
Ex: HCl(aq)  H+(aq) + Cl- (aq)
There are seven strong acids you should memorize. All other acids are considered weak.
"Strong" means 100% ionization of the 1st available proton in solution.
HCl
hydrochloric
HNO3
nitric
HBr
hydrobromic
HClO4 perchloric
HI
hydroiodic
HClO3 chloric (not very stable)
H2SO4 sulfuric
Traditional bases produce hydroxide ions, OH-(aq), in solution. The Brönsted-Lowery
definition of a base is as a proton acceptor. The strong bases include the hydroxides of the
column IA and IIA metals. They dissociate in water to produce aq. hydroxide ions.
There are a number of weak bases as well. Some of these produce few hydroxide ions due to
low solubility, such as Pb(OH)2. Others produce hydroxide ions by the hydrolysis of water
molecules. One such common weak base is ammonia:
NH3(aq) + H2O(l)  NH4+(aq) + OH-(aq)
Neutralization rxns occur when an aqueous acid and base react together. The net rxn for all such
neutralizations is the same:
H+(aq) + OH-(aq)  H2O(l)
A salt is the other product of an acid-base neutralization rxn. A salt may be defined as any ionic
compound. Even hydroxides may be considered salts, although they are also bases. There
are many salts.
The two most highly manufactured industrial salts in the U. S. are Na2CO3 sometimes called
soda ash, used in making glass, and NH4NO3, used in fertilizers.
4.9 Oxidation-Reduction Reactions
Reactions in which one or more electrons are transferred from one specie to another are called
oxidation/reduction reactions, or redox for short.
Example: 0
+1-1
+2-1
0
Fe(s) + 2HCl(g)  FeCl2(s) + H2(g)
e- transferred from Fe to H+
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Half Reactions
0
Fe  Fe2+ + 2eOxidation
0
2H+ + 2e-  H2 Reduction
**LEO goes GER
Losing electrons is oxidation, and gaining electrons is reduction.
Any element or ion, whether alone or part of a compound has an oxidation state, and can be
assigned an oxidation number. Sometimes these numbers represent real ionic charges, but
frequently they do not.
There is a simple set of rules for assigning oxidation numbers. They are:
1. Atoms of a free element are 0.
2. The oxidation number of a monatomic ions = the ionic charge.
3. Fluorine is -1 in all of its compounds.
4. The sum of the oxidation number of a polyatomic ion = the ionic charge.
5. The sum of the oxidation numbers of a neutral compound = 0
6. Oxygen in a compound is -2, except in peroxides when it is -1, or with F.
7. Hydrogen in a compound in +1, except in the metal hydrides of Columns IA and IIA
when it is -1.
Oxidation (losing electrons) results in an increase in oxidation state.
Reduction (gaining electrons) results in a decrease in oxidation state.
Elements getting oxidized or reduced is the passive view of things. We can also look at the
active side of things. An element that gives electrons away can be thought of as a reducing
agent. Likewise, an electron taker is an oxidizing agent.
4.10 Balancing Oxidation-Reduction Reactions
All reactions must be balanced with respect to atoms and charge.
Balancing equations by the half reaction method.
Often, all of the ions present at the beginning of a redox will not be represented, due to the fact
that they are spectators and are not important to the redox process.
Also, redox rxns may be accompanied by the terms "in acid solution or acidic environment" or
"in base or basic environment."
 If acid, use H+ and H2O to balance the equation.
 If base, use OH- and H2O to balance the equation.
For reactions in acid:
1. Write the separate half reactions for the oxidation and reduction.
2. For each half-reaction,
a. Balance all the elements except O and H by inspection.
b. Balance O using H2O.
c. Balance H using H+.
d. Balance the net charge for each half-reaction using e-.
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3. If necessary, multiply one or both half-reactions by an integer to equalize the number of
electrons transferred in the 2 reactions.
4. Add the half-reactions together and cancel identical species appearing on both sides of
the equation.
5. Check that the elements and charges are balanced.
For reactions in base:
Follow the same set of steps as above, except add the following adjustment.
1. Balance the half reactions as though they were in acid (using H+ and H2O).
2. Add enough OH- ions to neutralize any H+ ions to both sides of either half-rxn.
3. This keeps the half-rxns balanced and changes any H+ ions to water.
4. You will automatically now have the OH- and H2O on the correct sides of the half-rxns.
We will do examples in class.
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