You may recall from
Chapter 3, that we discussed labeling each reagent in the equation for a
chemical reaction by its state. And that we included aqueous solution in the list of states, along with solid, liquid, and gas.
Some of you may have correctly inferred that chemistry taking place in
solution is a very important topic. And that among all the possible
solvents, water is the most important.
What Makes Water Such a Good Solvent? You may recall from Chapter 2 that water has
a bent geometry with the two carbon-hydrogen bonds intersecting at an
angel of about 105°. These bonds are covalent, meaning that they
involve the sharing of electrons between the oxygen atoms and the
hydrogens. However, the sharing is not
equal the electrons in an oxygen hydrogen bond are more strongly
attracted to the oxygen than to the hydrogen. As a result, the atoms in a
are not electrically neutral The oxygen has a small net negative charge
(less than one unit), while the hydrogens each have small net positive
charges. This is illustrated in Figure 4.1 from the text. Water is said to
be a polar molecule on account
of this unequal charge distribution, and it is this polarity that makes
water such a good solvent.
How Ionic Substances Dissolve in Water: Recall that an ionic substance such as the
salt, sodium chloride (NaCl) is held together with ionic bonds, i. e.,
electrostatic attractive forces between oppositely charged ions. Clearly
if an ionic solid is to dissolve in water, there must be attractive forces
between water and ions of either charge. But as we have just seen, the
needed charges do exist in the water molecule. A cation such as sodium ion
(Na+) will be attracted to the partial negative charge on the
oxygen atom of water, and an anion such as chloride (Cl-) will be
attracted to the partial positive charges on the hydrogens. If these
attractive forces are more favorable than the forces that hold the ionic
solid together, it will break up, and the ions will disperse throughout
the water. Each dissolved ion will be surrounded by water molecules.
How Sodium Chloride Dissolves in Water: The dissolution of sodium chloride in water
can be written as the equation:
(aq) + Cl- (aq)
In the equation, Na+
(aq) represents a hydrated sodium ion, or a sodium ion surrounded by water
molecules, with the partially negatively charges oxygen atoms oriented inward
toward the positively charged sodium ions. Similarly, Cl- (aq)
represents a chloride (ion) surrounded by water molecules with their hydrogens
oriented inward. Figure 4.2 from the text gives this pictorial representation
of the dissolution process. In the upper panel, we see two water molecules
tugging at a chloride with their hydrogens and two other water molecules
tugging on a sodium ion with their oxygens. The second panel shows two hydrated
ions, one each of sodium ion and chloride, and it shows other water molecules
tugging at still more ions from the solid lattice:
How Nonionic Substances Dissolve in Water: Water is not a solvent for all substances.
(How fortunate for us?) Many ionic solids, such as silver chloride (AgCl)
do not dissolve in water. The forces holding the solid AgCl lattice
together are too strong to be overcome by the forces favoring the
formation of the hydrated ions, Ag+(aq) and Cl-(aq).
Covalently bonded, non polar substances like fats do not dissolve in
(pure) water because there is no driving force to form hydrated fat
molecules. However, covalently bonded substances like sugars and ethanol
will readily dissolve in water. The reason is that ethanol and sugars,
like water, are polar because, like water, they contain hydrogen-oxygen
bonds. Figure 4.3a illustrates the polarity of the OH bond in ethanol,
and Figure 4.3b pictures the attraction between ethanol and water that
4.2Strong and Weak (and non) Electrolytes
Solvent: We speak of a substance that dissolves something as a solvent.
Solute: The something that is dissolved in a solvent is called the solute.
Homogeneous Mixture: Recall from Chapter 2 that a solution is
homogeneous. The solute is mixed throughout the solution, and the
composition of the solution is the same everywhere in the solution.
Variable Composition: The composition of a solution of A (the
solute) in B (the solvent) is not fixed. Two solutions can be prepared
with the same amounts of solvent but different amounts of the solute and
therefore have different compositions.
Electrical Conductivity: A useful property of solutions is their
electrical conductivity. Pure water does not conduct electricity, but many
aqueous solutions do.
Strong Electrolytes: Strong electrolytes are solutions whose
electrical conductivity is high. They are good conductors of electricity.
Weak Electrolytes: Weak electrolytes are solutions that conduct
electricity, but not very well. They are still called conductors, but
only poor conductors.
Non Electrolytes: Non electrolytes do not conduct electricity.
A device for detecting electrical
conductivity in solutions is shown in Figure 4.4 in the text:
Arrhenius Theory of Electrical Conductivity: The first modern theory to explain electrical
conductivity in solutions was advanced by Svante Arrhenius, at the time a
doctoral student in physics at the University of Uppsala, Sweden. He
deduced the presence of charged particles in an electrically conductive
solution, and he postulated that the extent of the conductivity depended
on the number of ions present. Hence the solute in a strong electrolyte
would produce many ions and conduct electricity well, while the solute in
a weak electrolyte would produce relatively few ions and not conduct
electricity very well. In a non electrolyte, the solute would produce no
ions, hence no electrical conductivity for these solutions.
Strong Electrolytes: The conductivity of a strong electrolyte is
high, as illustrated in Figure 4.4a. This conductivity is produced by
solutes that are completely ionized in solution. Three classes of strong
Soluble Salts: Any salt that readily dissolves in water
produces a strong electrolyte when it is so dissolved. A good example is
sodium chloride (NaCl). See Figure 4.5.
Strong Acids: According to Arrhenius, an acid is a substance that ionizes in
aqueous solution to generate H+ ions (hydrogen
ions or protons). If the ionization is complete or nearly complete, the
acid solution is a good conductor of electricity, and the acid is regarded
as a strong acid. Some examples
of strong acids in aqueous solution are:
Sulfuric acid (H2SO4)
deserves a closer look. When sulfuric acid dissolves in water, the first
hydrogen dissociates completely to form protons in aqueous solution, but the
second hydrogen remains bonded to the sulfate. Thus aqueous sulfuric acid
contains mostly protons (H+(aq)) and hydrogen sulfate (HSO4-(aq))
Strong Bases: According to Arrhenius, a base is a substance that ionizes
in aqueous solution to produce hydroxide ions (OH-).
Bases that ionize completely are regarded as strong bases. The two most common strong bases are sodium
hydroxide and potassium hydroxide:
Weak Electrolytes: Many substances will form ions in aqueous solution,
but the extent of ionization is slight. For example, acetic acid (the
essence of vinegar), will readily dissolve in water, but only around 1% of
acetic acid molecules will ionize to form hydrogen ions and acetate ions.
The remaining acetic acid molecules will remain as electrically neutral
acetic acid molecules, even though they are totally dissolved in the
solution. The result is a solution that is electrically conductive, but
much less so than a comparative solution containing a strong electrolyte.
The most common weak electrolytes are:
Weak Acids: A weak acid is an acid that is only partially ionized in aqueous
solution. Thus an aqueous solution of acetic
acid (HC2H3O2) contains some hydrogen
ions (H+(aq)) and some acetate ions (C2H3O2-(aq)),
but most of the solute particles are undissociated acetic acid molecules
Weak Bases: Ammonia (NH3) is the most common of the weak bases.
It is a base because its aqueous solutions contain hydroxide ions (OH-(aq)).
It is a weak electrolyte because only a small fraction of ammonia
molecules form ions. Most of the ammonia remains as neutral ammonia molecules.
NH3 (aq) + H2O (l)
Nonelectrolytes: Nonelectrolytes are substances that dissolve
in water, but do not generate any ions. Their solutions do not conduct electricity
because of their total lack of ions. Common examples are:
also known as cane sugar.
4.3The Composition of Solutions: Many important chemical reactions occur in
solution. We need to be able to perform stoichiometric calculations on these
Stoichiometry in Solutions: We need to know:
The forms that
the reactants and products take in solution, and
The amounts of
chemicals present in the solutions.
In the previous chapter,
we learned to count the molecules in a reaction by weighing them, but when we
deal with solutions, we need a different way of counting. The most convenient
way to measure a liquid is to measure its volume. Thus, for example, if we had
prepared a 1.000 liter solution containing 1.000 mol of sodium nitrate and then
drew out a 50.00 mL portion of that solution, we would know that the 50.00 mL
portion contains 0.05000 mol of sodium nitrate.
Molarity: We can say that the above solution of sodium nitrate contains
1.000 moles of sodium nitrate in each 1.000 liters. Does this not look
like a useful way to characterize a solution? So we define molarity as:
We can quantitatively
describe our example solution as a 1.000 molar aqueous solution of sodium
nitrate, and we can abbreviate that to 1.000 M NaNO3.
Sample Exercise 4.1 (p. 134): What is the molarity of a solution
prepared by dissolving 11.5 g of solid NaOH in enough water to make up 1.5
L of solution? (The molar mass of NaOH is 40.00 g/mol.) We define some
familiar symbols: m represents mass, MM represents molar mass, and n
represents number of moles. And we need a new one: V
Sample Exercise 4.2: (p. 134): Calculate the molarity of a solution
prepared by dissolving 1.56 g of gaseous HCl in enough water to make 26.8
mL of solution. (The molar mass of HCl is 36.46 g/mol.)
Ion Concentrations: The above definition of molarity is based
upon how the solution was prepared. What is the volume of the solution?
How much solute did you weigh? It does not necessarily describe what is
actually in the solution. For example, our 0.192 M solution of NaOH in
Sample Exercise 4.1 does not contain NaOH molecules. The actual species
and their molarities are:
0.192 M sodium ions.
0.192 M hydroxide ions.
Suppose we had a
0.192 M solution of sodium
sulfate (Na2SO4). The species in solution would be:
0.192 M sulfate ions.
0.384 M sodium ions (because each mole
of sodium sulfate generates 2 moles of sodium ions).
Sample Exercise 4.3 (pp. 134-5): What are the concentrations of
each type of ion in the following solutions:
0.50 M Co(NO3)2
1 M Fe(ClO4)3
Moles of Solute in a Given Volume: If we know the molarity of a solution, we can
calculate how many moles of each of its components there are in a given
volume. We need simply to determine the molarities of each component and
multiply them by the volume. Remember that the units of the answer will be
Sample Exercise 4.4 (p. 135): How many moles of Cl-
ions are there in 1.75 L of 1.0 x 10-3M ZnCl2?
Standard Solutions: A standard
solution is a solution whose concentration is accurately known. The
procedure for preparing a standard solution is illustrated in Figure 4.10
form the text:
Sample Exercise 4.6 (pp. 136-7): A chemist needs 1.00 L of an
aqueous solution of K2Cr2O7 (potassium
dichromate) to analyze the alcohol contents of some samples of wine. How
much K2Cr2O7 must be weighed out to
prepare the solution?
The number of
moles of K2Cr2O7 required is:
Then the mass of
K2Cr2O7 (molar mass = 294.20 g/mol) that
must be weighed out is:
Dilution: Instead of preparing large volumes of solutions with relatively
low concentrations of solute, chemists will often
instead prepare a smaller volume of a much more concentrated solution and
dilute portions of it as needed. For example: Suppose the chemist in
Sample Exercise 4.6 had needed 20 L of 0.200 M K2Cr2O7. Rather than
perform the steps in Figure 4.10 twenty times over, she might instead have
prepared 1.00 L of a 4.00 M
solution. Then each time she needed some 0.200 M solution for an analysis, she could accurately measure (by
pipette) 50.00 mL of the 4.00 M
solution into another 1 liter volumetric flask and add water to the 1.00 L
mark. The key to dilution calculations is that:
Moles of solute before dilution = Moles of solute
In a dilution problem,
you will know or be given information to calculate 3 of the 4 quantities in the
Example from page 137: We need 500. mL of 1.00 M acetic acid (HC2H3O2) from
a 17.4 M stock solution. What
volume of stock solution do we need? (Work out on whiteboard.)
Sample Exercise 4.7 (pp. 138-9): What volume 16 M sulfuric acid (H2SO4)
must be used to prepare 1.5 L of a 0.10 M H2SO4
solution? (The solution in the text is a bit convoluted. Well work it out
directly on the whiteboard.)
4.4Types of Chemical Reactions in Solution
These topics take up the rest of Chapter 4. We will
cover them next week.