Chapter 14 – Acids and Bases
History of Acids & Bases
• Vinegar was probably the only known acid in
ancient times.
• Strong acids such as sulfuric, nitric and
hydrochloric acids were not discovered until
after the 12th century.
• Over the years, there have been many
attempts to define acids and bases.
Old Definitions of Acids and Bases
• At first, acids and bases were defined in terms of their observed
properties such as taste, effects on indicators and reactions with other
substances.
• In the 17th century, Boyle described the properties of acids in terms of
taste, their action as solvents and how they changed colour of certain
vegetable materials.
• He also noticed that alkalis (soluble bases) could reverse the effects of
acids.
• Lavoisier, in the 18th century, thought that acidic properties were due to
the presence of oxygen.
• In 1810, Davy suggested that the acid properties of substances were
associated with hydrogen and not oxygen.
• In 1887, Arrhenius defined acids as substances that produced hydrogen
ions (H+) in water while bases produced hydroxide ions (OH-) in water.
• According to his theory, when acids and bases react together, the H+ and
OH- form water according to the equation:
H+ + OH-  H2O
Arrhenius called this a neutralisation
Definitions cont…
• There were, however, limitations to these theories.
• Arrhenius’ definition for example was restricted to
acids and bases in water.
• One of the more useful definitions used today was
first proposed by the Bronsted and Lowry
• Bronsted and Lowry described the reactions of acids
as involving the donation of a hydrogen ion (H+).
• A hydrogen ion is a hydrogen that has lost its only
electron.
• In most cases, a hydrogen ion is a proton.
Bronsted-Lowry Acids and Bases
• According to the Bronsted-Lowry theory, a
substance behaves as an acid when it donates
a proton, ie H+ to a base.
• A substance behaves as a base when it accepts
a proton from an acid. Hence:
– Acids are proton donors and
– Bases are protons acceptors.
Bronsted-Lowry Acids and Bases
• As protons are exchanged from an acid to a base, this
definition explains why acids and bases react together.
• In an aqueous solution of hydrogen chloride, nearly all the
hydrogen chloride is present as ions – virtually no molecules
of hydrogen chloride remain.
• This solution is known as hydrochloric acid.
• In this reaction, each hydrogen chloride molecule has donated
a proton to a water molecule.
• According to the Bronsted-Lowry theory, the hydrogen
chloride has acted as an acid.
• The water molecule has accepted a proton from the hydrogen
from the hydrogen chloride, so has acted as a base.
HCl(g) + H2O(l)  H3O+(aq) + Cl-(aq)
Acid-base Conjugate Pairs
• Because HCl and Cl- can be formed from each other
by the loss or gain of a single proton, they are called
a conjugate acid/base pair.
• Similarly, H3O+ and H2O are also a conjugate pair.
• A conjugate pair is two species which differ by a
proton.
• For the reaction between HCl and H2O, the conjugate
pairs are shown as:
HCl(g) + H2O(l)  H3O+(aq) + Cl-(aq)
Blue = bases
Red = acids
The H+ ion in Water
• A hydrogen ion (or proton) in solution is
represented as H3O+(aq) or more simply
H+(aq) and is called the hydronium ion.
• The hydronium ion itself attracts more water
molecules and is further hydrated.
• However, these water molecules are not as
strongly attracted and their number is not
constant.
Some Common Acids & Bases
Amphiprotic Substances
• Some substances can behave as either acids or
bases, depending on what they are reacting with.
• These substances are given the name amphiprotic
substances.
• In equation 1 below, water readily accepts a proton
from sulfuric acid and acts as a base.
• In equation 2, water donates a proton to the oxide
ion and acts as an acid.
Eqn 1:H2SO4(aq) + H2O(l)  HSO4-(aq) + H3O+(aq)
Eqn 2: O2-(aq) + H2O(l)  OH-(aq) + OH-(aq)
Amphiprotic Substances cont…
• If the solute is a stronger acid than water, then
water will act as a base.
• If the solute is a stronger base than water,
then the water will act as an acid.
Amphiprotic Substances cont…
• When an amphiprotic substance is placed in water, it reacts as
both an acid and a base.
• For example, the hydrogen carbonate (HCO3-) ion reacts
according to the equations:
HCO3-(aq) + H2O(l)  H2CO3(aq) + OH-(aq)
HCO3-(aq) + H2O(l)  CO32-(aq) + H3O+(aq)
• Since HCO3- can act as both acid and base, it is amphiprotic.
• Although both reactions are possible for all amphiprotic
substances in water, generally one of these reactions occurs
to a greater extent.
• The dominant reaction can be identified by measuring the pH
of the solution.
Acid & Base Strength
• Experiments show that different acid solutions of the same
concentration do not have the same pH.
• Some acids donate a proton more readily than others.
• The strength of an acid is based on its ability to donate
hydrogen ions.
• The strength of a base is based on its ability to accept
hydrogen ions.
• Since aqueous solutions of acids and bases are most
commonly used, it is convenient to use an acid’s tendency to
donate a proton to water, or a base’s tendency to accept a
proton, as a measure of its strength.
Strong Acids
• Acids that ionise completely in solution are
called strong acids.
• Strong acids donate protons easily.
• Solutions of strong acids would contain ions
and virtually no unreacted acid molecules.
• The most common strong acids are
hydrochloric acid, sulfuric acid and nitric acid.
Weak Acids
• An acid that does not fully ionise is called a
weak acid.
• An example of a weak acid is ethanoic acid.
• Only a small proportion of ethanoic acid
molecules are ionised.
• A weak acid can be shown be the presence of
reversible arrows.
CH3COOH(l) + H2O(l)
CH3COO-(aq) + H3O+(aq)
Strong Bases
• The ionic compound sodium oxide (Na2O)
dissociates in water, releasing sodium ions
(Na+) and oxide ions (O2-).
• The oxide ions react completely with the
water, accepting a proton to form hydroxide
ions (OH-).
• The oxide ion is an example of a strong base.
• Strong bases accept protons easily.
Weak Bases
• Ammonia is a covalent molecular compound that
ionises in water by accepting a proton.
• This ionisation process can be represented by the
equation:
NH3(aq) + H2O(l)
NH4+(aq) + OH-(aq)
Only a small proportion of ammonia molecules ionise.
This is shown in the equation by the presence of
reversible arrows.
Ammonia is a weak base in water.
Polyprotic Acids
• Some acids are capable of donating more than one
proton from each molecule and are said to be
polyprotic.
• The number of hydrogen ions an acid can donate
depends on the structure of the acid.
• Monoprotic acids: can donate only one proton and
include HCl, HF, HNO3, CH3COOH.
• Diprotic acids: can donate two protons and include
H2SO4, H2CO3,
• Triprotic acids: can donate three protons and include
H3PO4, H3BO3.
Polyprotic Acids cont…
• Polyprotic acids do not donate all protons at
once, but do so in steps when reacting with a
base.
• Sulfuric acid (H2SO4) is diprotic, meaning it has
two protons that it can donate to a base.
• A diprotic acid ionises in two stages, for example:
STAGE 1: H2SO4(l) + H2O(l)  HSO4-(aq) + H3O+(aq)
STAGE 2: HSO4-(aq) + H2O(l)  SO42-(aq) + H3O+(aq)
Polyprotic Acids cont…
• When added to a base stronger than water, a weak acid will
ionise to a greater extent.
• For example, a strong base such as OH- will accept a second
proton from H2SO4 and the second and third proton from
H3PO4.
• Similarly a weak base will ions to a greater extent if added to a
strong acid.
• Sometimes there are more hydrogens in a molecule than can
actually be donated.
• For example CH3COOH contains four hydrogen and yet will
only donate one.
• Only the hydrogen involved in the polar OH- bond is donated.
• In general each hydrogen ion that is donated by an acid
molecule is involved in a polar bond.
Relative Strengths of Acid Base
Pairs
Strength vs. Concentration
• It is important that the terms strong and weak
are not confused with the terms concentrated
and dilute.
• Concentrated and dilute describe the amount
of acid or base dissolved in a given volume of
solution.
• The terms strong and weak describe how
readily an acids donates, or base accepts a
proton.
Strength vs. Concentration cont…
Qualitative vs. Quantitative
• Terms such as concentrated and dilute, or
weak and strong are qualitative, or descriptive
terms.
• Solutions can be more accurately described by
stating concentration in mol/L or g/L.
• This is a quantitative description.
Acidic, Basic and Neutral Solutions
• The acidity of a solution is a measure of the
concentration of hydrogen ions present.
• The higher the concentration of hydrogen ions, the
more acidic the solution.
• Water has the ability to act as either an acid or a
base.
• Pure water undergoes self ionisation to a small
extent with allows it to conduct electricity slightly.
• This can be represented by the equation:
H2O(l) + H2O(l)
H3O+(aq) + OH-(aq)
Acidic, Basic and Neutral Solutions
cont…
• Acidic solutions contain a greater
concentration of H3O+ than OH-.
• Neutral solutions contain equal
concentrations of H3O+ and OH-.
• Basic solutions contain a lower concentration
of H3O+ than OH-.
Measuring Acidity
• [H3O+] x [OH-] = 10-14M2
• Pure water is neutral so [H3O+] = [OH-]
• If either the [H3O+] or [OH-] in an aqueous
solution is increased, the other must decrease
proportionally.
• At 25°C, a solution is:
Acidic if [H3O+]>10-7M and [OH-]<10-7M
Neutral if [H3O+] = 10-7M = [OH-]
Basic if [H3O+]<10-7M and [OH-]>10-7M
Acidity Example
• In a 5.6x10-6M HNO3, solution at 25°C, calculate the
concentration of:
a. H3O+ ions
HNO3 is a strong acid and ionises completely to
produce 5.6x10-6M of H+ ions.
b. OH- ions
[H3O+] x [OH-] = 10-14
5.6x10-6 x [OH-] = 10
[OH-] = 10-14/5.6x10-6
[OH-] = 1.79 x 10-9M
The pH Scale
• This scale is a useful way of indicating the
acidity of a solution.
• pH = -log10[H3O+]
• The pH of a solution decreases as the
concentration of hydrogen ions increases.
• Acidic solutions have a pH<7
• Basic solutions have a pH>7
• Neutral solutions have a pH=7
Calculating pH Example 1…
• What is the pH of a solution in which [H+] =
0.0135M
pH = -log[H+]
pH = -log(0.0135)
pH = -(-1.87)
pH = 1.87
Calculating pH Example 2…
• What is the pH of a 0.0050M of Ba(OH)2?
Step 1: Find concentration of H+
Ba(OH)2(aq)  Ba2+(aq) + 2OH-(aq)
Ba(OH)2 is completely dissociated in water and each mole of Ba(OH)2 dissociates to
release 2 moles of OH- ions
So, [OH-] = 2 x [Ba(OH)2]
= 2 x 0.0050
= 0.010M
Since [H+] x [OH-] = 10-14
[H+] x 0.010 = 10-14
[H+] = 10-14 / 0.010
[H+] = 10-12
Step 2: Calculate the pH
pH = -log[H+]
= -log(10-12)
= 12
Calculating the Concentration of H+
in a solution of a given pH
[H+] = 10-pH
If the pH is 5.00, what is the [H+]?
[H+] = 10-5
= 0.0001M