AQUEOUS SOLUTIONS
An aqueous solution is one in which a solute is dissolved in water (the
solvent).
Aqueous solutions are very common in chemistry because water is plentiful
and will dissolve a wide range of substances.
Species present in solution are surrounded by water molecules. We indicate
this by writing X2(aq) or X(aq).
Properties of Water
Pure water:
 is a clear, colourless liquid
 has a freezing point of 0C and a boiling point of 100C (at a pressure of
1atmosphere)
 is a poor conductor of electricity
 has a pH of 7
The water molecule
 is polar, with an unusually high dipole. This means that water molecules
are attracted to the charged particles of compounds. (diagram below)
 is hydrogen bonded.
Ions in salt crystal e.g. NaCl
Na+(aq)
Cl¯(aq)
Properties of aqueous solutions
Electrical conductivity:
Pure water is a very poor electrical conductor. An electrolyte is defined as a
substance which when dissolved in water, forms a solution which conducts
electricity e.g. sodium chloride. Sucrose dissolves in water, but the solution
does not conduct electricity; sucrose is a non-electrolyte. Arrhenius realised
that an electrolyte is made of ions. He also found materials that are weak
electrolytes – they conduct electricity better than pure water but not as well as
a sodium chloride solution of the same strength.
Strong electrolytes
NH4Cl NaCl
NaOH Ca(OH)2
CH3COONa
Weak electrolytes
CH3COOH
NH3
Non-electrolytes
C6H12O6
C12H22O11
CH3CH2OH
Freezing point depression:

In 1883 the French chemist Raoult discovered that the freezing point of
water decreases by 1.86 C for each mole per litre of a non-electrolyte
dissolved.

Arrhenius investigated the freezing point depression of electrolyte
solutions. He discovered that a 1.0 mol L-1 of sodium chloride has a
freezing point of –3.7 C (twice –1.86), while the same concentration of
calcium chloride solution freezes at –5.6 C (three times –1.86).

The freezing point is related to the number of particles in the solution.
Strong electrolytes dissociate into ions in solution – two ions for NaCl,
three ions for CaCl2. Weak electrolytes such as ethanoic acid have a
freezing point depression slightly more than 1.86 C mol-1L-1, which
indicates that only very few of the solute molecule have dissociated into
ions.

The freezing point of a solution can be used to determine the molar mass
of non-electrolytes. It can also indicate whether a solute is dissociated
significantly or not.
pH
pH = -log10[H3O+]
Pure water contains a very small number of ions.
2H2O(l)  H3O+(aq) + OH-(aq)
Pure water has a pH of 7. This means that [H3O+] = 10-7mol L-1
and from the equation we see that [OH-] = 10-7mol L-1 also.
Neutral solutions such as water, NaCl solution and sucrose solution all have
pH of 7, but other substances dissolve in water to form acidic or basic
solutions.
The Bronsted-Lowry theory of acids and bases defines an acid as a proton
donor and a base as a proton acceptor. For every acid there is a conjugate
base.
E.g.
HCl(g) + H2O(l)  H3O+(aq) + Cl-(aq)
NO3-(aq) + H2O(l)  HNO3(aq) + OH-(aq)
In aqueous solution water is always present and can act as an acid or a base.
In the examples above the acid is in bold type and its conjugate base is
underlined. Since water can act as both proton donor and acceptor it is said to
be amphiprotic.
NaOH dissolves in water to form a basic solution due to the OH- ions. Sodium
carbonate also forms a basic solution when it dissolves in water. Why?
CO32-(aq) + H2O(l)  HCO3-(aq) + OH-(aq)
The carbonate ions react with water – undergo hydrolysis. This is because
the carbonate ion (CO32-(aq)) is the conjugate base of the acidic, hydrogen
carbonate ion (HCO3-(aq)).
Two other ions, which hydrolyse with water, are the ammonium ion and the
ethanoate (acetate) ion.
Complete the equations for the hydrolysis of each of these.
CH3COOH(aq) + H2O(l)
+
NH4
(aq)+
H2O(l)
Whenever these ions, or the conjugates of other weak acid or bases, are
present in solution, hydrolysis can occur and affect the pH of the solution.
The composition of aqueous solutions:
The physical properties of a solution (conductivity, freezing point and pH) help
us determine what species are present in an aqueous solution of a given
concentration.
All aqueous solutions contain:
 The substance dissolved (either molecular form or as ions)

Liquid water , H2O(l)

H3O+(aq) and OH-(aq) (from the ionisation of water)
Concentration of water molecules: there are 1000 g of water in 1 L of
water, so the concentration of H2O(l) in moles per litre for dilute solutions is:
c(H2O) =
=
c
M
1000 gL1
18 gmol 1
= 55.56 mol L-1
The concentration of the H3O+(aq) and OH-(aq) depends on the solution.
Relative Concentrations of Species Present in Aqueous Solutions.
Examples: (a) NaCl
A 1 mol L-1 solution of sodium chloride is a good conductor (i.e. a strong
electrolyte) and has a pH of 7 (i.e. neutral). The freezing point indicates that
there are 2 moles of ions present in 1 litre of solution. The solution therefore
contains:





55.5 mol L-1 H2O(l)
1.0 mol L-1 Na+(aq)
1.0 mol L-1 Cl- (aq)
10-7 mol L-1 H3O+(aq)
10-7 mol L-1 OH-(aq)
Na+
H2O
Cl- H3O+ OH-
(b) NH4Cl
A 1 mol L-1 solution of ammonium chloride is a good conductor (i.e. a strong
electrolyte) and has a pH of 5 (i.e. acidic). The freezing point indicates that
there are 2 moles of ions present in 1 litre of solution. The solution therefore
contains:






55.5 mol L-1 H2O(l)
1.0 mol L-1 NH4+(aq)
1.0 mol L-1 Cl- (aq)
10-5 mol L-1 H3O+(aq)
10-5 mol L-1 NH3(aq)
10-9 mol L-1 OH-(aq)
NH4+ Cl-
H2O
H3O+ OH- NH3
This is because the ammonium ion has hydrolysed to form ammonia and
hydronium ions:
+
NH4 (aq) + H2O(l)
+
NH3(aq) + H3O
(aq)
and since the ionic product of water is constant, as the hydronium ion
concentration goes up, the hydroxide ion concentration must come down.
Kw = [H3O+][OH-] = 1 x 10-14