REVIEWTOPIC 8/18: ACIDS AND BASES
A. KEY DEFINITIONS (MEMORIZE!)
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A Bronsted-Lowry acid is a proton (H+) donor and a Bronsted-Lowry base is a proton acceptor
A Lewis acid is an electron pair acceptor and a Lewis base is an electron pair donor.
B. YOU SHOULD KNOW:
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Acids turn universal indicator and litmus red
Acids are neutralized by bases
Acids form a salt and water with metal oxides
Acids form a salt, water and carbon dioxide gas with carbonates and hydrogen carbonates
Acids react with metals to produce a salt and hydrogen gas
When bases dissolve in water they are called alkalis
Weak acids are only partially dissociated and because of this they have low concentrations of ions, so they do not
conduct electricity as well as strong acids, which are fully dissociated into their ions
Strong acids can be distinguished from weak acids of the same concentration by comparing their conductivities
Strong acids can be distinguished from weak acids of the same concentration by comparing their reactions with
magnesium of calcium carbonate – the strong acid will react more rapidly, leading to quicker evolution of gas
Strong acids will generally have a lower pH, which can be measured using a pH meter or universal indicator
The pH scale gives a measure of relative acidity or basicity, and on this scale pH 1 is a strong acid, pH 7 is neutral
and pH 14 is strongly alkaline
The causes of acid deposition and the reactions leading to the formation of acid rain
Kw is the ionic product of water, given by: Kw = [H+] x [OH-] or Kw = [H3O+] x [OH-] and at 298 K its value is 1.0 x 10-14
pKw is given by pKw = -logKw which at 298 K is equal to 14
a weak acid is represented by HA and dissociates like this(do not forget the equilibrium arrows!):
HA(aq)
H+(aq) + A-(aq) and the acid ionization constant, Ka, is given by:
Ka =
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a weak base is represented by B and reacts with water like this (do not forget the equilibrium arrows):
B(aq) + H2O(l)
BH+(aq) + OH-(aq) and the base dissociation constant, Kb, is given by:
Kb =
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[ H  ( aq) ][ A ( aq) ]
[ HA( aq) ]
[ BH  ( aq) ][OH  ( aq) ]
[ B( aq) ]
a low Ka or high pKa means that an acid is weak- so, for example, an acid with a pKa of 5.2(Ka = 6.3 x 10-6) is weaker
than one with a pKa of 2.1(Ka = 7.9 x 10-3) *Remember: (↓Ka = ↑pKa = weaker the acid)…and vice versa
a low Kb means that the base is weak
weak acids have low Ka values, and strong acids do not have Ka values…but they would be extremely large
an acidic buffer contains a weak acid and its salt (salt contains conjugate base of the weak acid)
a basic buffer contains a weak base and its salt (salt contains conjugate acid of the weak base)
buffers resist changes in pH when a small amount of acid or base is added
some salts of metals are acidic because the charge density on the cation is so high that it causes the hydrolysis of
water to happen, for example: [Fe(H2O)6]3+(aq)
[Fe(H2O)5(OH)]2+(aq) + H+(aq)
the acidity and basicity of salts formed by neutralization as shown in the table:
ACID
BASE
SALT FORMED
Strong
Strong
Neutral
Weak
Strong
Basic
Strong
Weak
Acidic
Weak
Weak
Neutral(approx.)
the shapes of the eight possible titration curves that can be produced by adding a strong or weak base to a strong
or weak acid, or vice versa
the starting pH will depend on the concentration and level of dissociation of the acid/base
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if you titrate a weak acid with a strong base, at the half neutralization point (i.e. at half the volume needed to
neutralize the acid, in other words, half-way to the equivalence point), the pH of the acid will equal the pKa of the
acid; similarly, if a weak base is titrated with a strong acid, pKb = pOH at half-way to neutralization
 the equivalence point will not always be at pH 7, because of salt hydrolysis
 acid-base indicators are weak acids and have different colours in their undissociated form (HIn) and dissociated
form (In-)
 the colour of an indicator will depend on the pH of the solution it is in (if the pH is the same and the pK a, there will
be equal quantities of HIn and In-) HIn(aq)
H+(aq) + In-(aq)
colour 1
colour 2
C. YOU SHOULD BE ABLE TO:
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Identify Bronsted-Lowry acids and bases in a reaction, and remember that the conjugate pair will differ by H +, for
example CH3COOH and CH3COOIdentify Lewis acids and bases in a reaction, and remember that all Bronsted-Lowry acids and bases are also Lewis
acids and bases
Remember that substances that are only acids and bases according to the Lewis theory will not contain H + and are
often transition metal or aluminum compounds in solution (so, for example. HCl is a Bronsted-Lowry acid and also
a Lewis acid, but AlCl3 is only a Lewis acid as it does not have a H+ to donate).
Work out the change in [H+] when the pH changes, as each pH unit change is equivalent to a 10-fold change in [H+]
concentration, so an increase in pH of 1 means a decrease in [H+] of 10, and a decrease in pH of 3 means an
increase in [H+] of 1000
Identify which substances are more acidic or alkaline from their pH values
Calculate the value of [H+] and [OH-] using Kw
Remember that the dissociation of water is endothermic, so increasing the temperature above 298 K will lead to
greater dissociation, and so more [H+] and lower pH (the water is still neutral, even though pH<7, because [H+] =
[OH-])
Solve problems using Kw = Ka x Kb and pKw = pKa + pKb and pKw = pH + pOH (i.e. 14 = pH + pOH)
Calculate the pH of a buffer solution using: Ka =
[ H  ( aq) ][ A ( aq) ]
by rearranging the equation and
[ HA( aq) ]
remembering that [H+] is no longer equal to [A-] or similarly using: Kb =
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[ BH  ( aq) ][OH  ( aq) ]
[ B( aq) ]
Identify Lewis acids and basis in a reaction, and remember that all Bronsted-Lowry acids and bases are also Lewis
acids and bases
Remember that substances that are only acids and bases according to the Lewis theory will not contain H+ and
Lewis acids are often transition metal or aluminum compounds in solution.
Suggest a suitable indicator by looking at the data booklet values (table 16), and the equivalence point of the
titration, as the indicator must change colour in this range
D. BE PREPARED
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Practise writing equations with hydrochloric (HCl), sulfuric (H2SO4), nitric (HNO3) and ethanoic (CH3COOH) acids
with a variety of metals, and basic compounds such as NaOH or CuO.
Rain water is naturally acidic due to CO2 in the air. When the pollutants SO2 and NOx react with rain water, this
leads to “acid rain”. Acid rain reacts with limestone buildings, eroding them, and changes the pH of soil and lakes.
Strong acids are HCl, H2SO4 and HNO3, and examples of weak acids are carbonic acid (H2CO3) and the carboxylic
acids such as ethanoic acid, CH3COOH. The strong bases (alkalis) are the hydroxides of group 1 (NaOH and so on)
and barium hydroxide, Ba(OH)2, and weak bases are ammonia (NH3) and the amines, methylamine (CH3NH2), for
example.
When calculating Ka and Kb it is usual to make assumptions – these are that the acid or base has not dissociated,
and hence the concentration given is the one you use in your calculation, and that the temperature is 298 K.
A buffer can be produced by partial neutralization of the acid/base or by adding a salt to the acid/base.
An example of an acidic buffer is CH3COOH/CH3COONa and an example of a basic buffer is NH3/NH4Cl
Indicators cannot be used to judge the equivalence point in a titration between a weak acid and weak base
because there is no sharp change in pH, so no sharp change in colour. `