ACIDS AND BASES
• Define acids and bases according to Arrhenius and LowryBrønsted theories
• Distinguish between strong acids/bases and weak
acids/bases with examples.
• Distinguish between concentrated acids/bases and dilute
acids/bases
• Write down the reaction equations of aqueous solutions of
acids and bases
• Identify conjugate acid-base pairs for given compounds
• Describe a substance that can act as either acid or base as
amphiprotic
• Write down neutralisation reactions of common laboratory
acids and bases
• Determine the approximate pH (equal to, smaller than or
larger than 7) of salts in salt hydrolysis.
• Motivate the choice of a specific indicator in a
titration
• Perform stoichiometric calculations based on
titrations of a strong acid with a strong base, a
strong acid with a weak base and a weak acid with a
strong base. Calculations may include percentage
purity
• Explain the pH scale as a scale of numbers from 0 to
14 used to express the hydrogen ion concentration.
• Calculate pH values of strong acids and strong
bases.
• Define the concept of Kw as the equilibrium constant
for the ionisation of water – the ionic product of
water (ionisation constant of water).
• Explain the auto-ionisation of water
DEFINITIONS
• Arrhenius theory:
An acid is a substance that produces hydrogen ions
(H+) in water.
A base produces hydroxide ions (OH-) in water.
• Lowry-Brønsted theory:
An acid is a proton (H+ ion) donor.
A base is a proton (H+ ion) acceptor.
• Ionisation: The process whereby covalent molecules
will produce ions in solution for the first time
• Dissociation: The process where ionic compounds
break up into their individual ions in solution
• Strong acid is an acid that ionises completely in water
to form a high concentration of H3O+ ions.
Examples: hydrochloric acid, sulphuric acid and
nitric acid.
• Weak acid is an acid that ionises incompletely in
water to form a low concentration of H3O+ ions.
Examples: ethanoic acid and oxalic acid.
• Strong base is a base that dissociates completely in
water to form a high concentration of OH- ions.
Examples: sodium hydroxide and potassium
hydroxide.
• Weak base is a base that dissociates/ionises
incompletely in water to form a low
concentration of OH- ions.
Examples: Ammonia, calcium carbonate,
potassium carbonate, calcium carbonate and
sodium hydrogen carbonate.
• Concentrated acids contains a large amount
(number of moles) of acid in proportion to the
volume of water.
• Dilute acids contain a small amount (number of
moles) of acid in proportion to the volume of
water.
TYPES OF ACIDS
• HCℓ(g) + H2O(ℓ) → H3O+(aq) + Cℓ-(aq)
(HCℓ is a monoprotic acid.)
• H2SO4(aq) + H2O(ℓ) → H3O+(aq) + HSO4− (aq)
HSO4− (aq) + H2O(ℓ) → H3O+(aq) + SO4−2 (aq)
(H2SO4 is a diprotic acid.)
• H3PO4(aq) + H2O(ℓ) → H3O+(aq) + H2PO4−(aq)
H2PO4−(aq) + H2O(ℓ) → H3O+(aq) + HPO4−2(aq)
HPO4−2(aq) + H2O(ℓ) → H3O+(aq) + PO4−3(aq)
(H3PO4 is a triprotic acid.)
conjugate acid-base pairs
• When the acid, HA, loses a proton, its conjugate
base, A-, is formed.
• When the base, A-, accepts a proton, its
conjugate acid, HA, is formed. These two are a
conjugate acid-base pair
• Protolysis is the acid-base reaction where proton
transfer occurs. That is when conjugate acid-base
pairs occurs.
• A strong acid has a weak conjugate base and a
strong base has a weak conjugate acid and vice
versa
CONJUGATE ACID-BASE PAIRS
AMPHOLYTE
• A substance that can act as either acid or base
can also be referred to as amphiprotic.
Substances such as H2O and HSO4- that can react as
both an acid and a base and therefore are
amphiprotic and are called ampholytes
Neutralisation reactions
The reaction between an acid and a base to form salt and water
Acid + base → salt + water
Hydrolysis is the reaction of a salt with water
1. Hydrolysis of the salt of a weak acid and a
strong base results in an alkaline solution, i.e.
the pH > 7. Examples of such salts are sodium
ethanoate, sodium oxalate and sodium
carbonate.
2. Hydrolysis of the salt of a strong acid and a
weak base results in an acidic solution, i.e. the
pH < 7. An example of such a salt is ammonium
chloride.
3. The salt of a strong acid and a strong base does
not undergo hydrolysis and the solution of the
salt will be neutral, i.e. pH = 7.
• The equivalence point of a titration is the point
at which the acid/base has completely reacted
with the base/acid.
• The endpoint of a titration is the point where the
indicator changes colour.
Auto-ionization of Water
• Self-ionization is a reaction in which two like
molecules react to give ions (amphiprotic
therefore can react with self)
– In the case of water, the following equilibrium is
established.
H 2O(l )  H 2O(l )
H 3O  (aq )  OH  (aq )
– The equilibrium-constant expression for this
system is:


[H 3O ][OH ]
Kc 
2
[ H 2O ]
• The concentration of ions is extremely small (equilbrium
lies to far left), so the concentration of H2O remains
essentially constant. This gives:


[H 2O] K c  [H 3O ][OH ]
2
constant
• We call the equilibrium value for the ion product [H3O+][OH-] the
ion-product constant for water, which is written Kw.


K w  [H 3O ][OH ]
• At 25 oC, the value of Kw is 1.0 x 10-14.
• Like any equilibrium constant, Kw varies with
temperature. Kw means water + water and
basis of acid/base scale in aqueous solutions
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• Because we often write H3O+ as H+, the ion-product
constant expression for water can be written:


K w  [H ][OH ]
• Using Kw you can calculate the concentrations of H+
and OH- ions in pure water.
H 3O  (aq )  OH  (aq )
H 2O(l )  H 2O(l )
• Thus, the concentrations of H+ and OH- in pure water
are both 1.0 x 10-7 mol.dm-3. Baseline for what we
call a neutral solution with water as solvent.


K w  [ H 3 O ][OH ]
7
7
K w  (1.0 X 10 )(1.0 X 10 )
K w  1.0 X 10
14
Solutions of Strong Acid or Base
• By dissolving substances in water, you can
alter the concentrations of H+(aq) and OH(aq).
– In a neutral solution, the concentrations of H+(aq)
and OH-(aq) are equal, as they are in pure water.
[ H 3 O  ]  [OH  ]  1.0  10 7 mol.dm 3
– In an acidic solution, the concentration of H+(aq) is
greater than that of OH-(aq).
[ H 3 O  ]  1.0  10 7 mol.dm 3  [OH  ]
– In a basic solution, the concentration of OH-(aq) is
greater than that of H+(aq).
[ H O  ]  1.0  10 7 mol.dm 3  [OH  ]
3
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Solutions of Strong Acid or Base
• At 25°C, you observe the following
conditions.
– In an acidic solution, [H+] > 1.0 x 10-7 mol.dm-3.
– In a neutral solution, [H+] = 1.0 x 10-7 mol.dm-3.
– In a basic solution, [H+] < 1.0 x 10-7 mol.dm-3.
Realize these definitions for acid/base/neutral
solutions is based as water as solvent. If
different
solvent
or
temperature,
the
concentration would be different.
The pH Scale
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Calculating pH
• Although you can quantitatively describe the acidity of a
solution by its [H+], it is often more convenient to give acidity
in terms of pH (power of the hydrogen ion). Easier to see
larger value: 10-7 vs 10-8
– The pH of a solution is defined as the negative
logarithm of the molar hydrogen-ion concentration.
– Basically changing 1.0 x 10-7 mol.dm-3 to log
scale.
1.0 number indicates where between the 10-6 and
10-7 --> 7.00

pH   log[H ]
20
• For a solution in which the hydrogen-ion
concentration is 1.0 x 10-3 mol.dm-3, the pH
is:

pH   log[ H ]
3
pH   log(1.0  10 )  3.00
• A sample of orange juice has a hydrogen-ion
concentration of 2.9 x 10-4 mol.dm-3. What is
the pH?

pH   log[ H ]
4
pH   log( 2.9  10 )
pH  3.54
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• The pH of human arterial blood is 7.40. What
is the hydrogen-ion concentration?

pH   log[ H ]
7.40  log[ H  ]

[ H ]  anti log( 7.40)

[ H ]  10

7.40
8
[ H ]  4.0  10 mol.dm
3
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Calculate pH given concentration of
base
• Calculate the pH in 0.10 mol.dm-3 NaOH.
NaOH + HCl
NaCl + H2O
Option A
1.0 × 10−14 = 𝐻3 𝑂+ 0.10
−14
1.0
×
10
𝐻3 𝑂+ =
0.10
-3
𝐻3 𝑂+ = 1.0 × 10−13 mol.dm
Option B
𝑝𝑂𝐻 = −𝑙𝑜𝑔 𝑂𝐻−
= −𝑙𝑜𝑔 0.10
= 1.0
𝑝𝐻 + 𝑝𝑂𝐻 = 14
p𝐻 + 1.0 = 14
𝑝𝐻 = 14 − 1.0
𝑝𝐻 = 13
𝑝𝐻 = −𝑙𝑜𝑔 𝐻3 𝑂+
= − log 1.0 × 10−13
=13
23