ACIDS AND BASES
Topic 10: HSC course
Core 10 – Acids and Bases
This topic develops the student’s understanding of acids and bases, the concept of
pH,
and applies this knowledge in volumetric analysis, where simple analytical
techniques
are also comprehended.

1. Acids and basesthe way people’s understanding of acids and bases has
changed from early concepts to the Arrhenius theory and to the
Bronsted–Lowry theory.

proton transfer as the common feature of all acid/base systems
(a) acids as proton donating species (molecule or ion)
(b) bases as proton accepting species (molecule or ion)
(c) diprotic and polyprotic acids
(d) amphiprotic species
(e) conjugate pairs

strengths of acids may be compared by the extent of their proton transfer to
water forming aqueous hydrogen ions

[H + ][OH - ] = 1 x 10 -14 at 25°C

pH related to [H + ]

calculations to produce Ka from pH and [H + ]

pH calculations
(a) [H + ] to pH
(b) pH to Ka

importance of pH in many natural systems — qualitative description of buffers
only

calculations involving pH, including the pH of mixtures of acids and bases.
2. Volumetric analysis

Volumetric analysis involving strong acid – strong base, weak acid – strong
base and strong acid – weak base
This analysis involves:
(a) the use of an appropriate reaction
(b) the preparation of a standard solution
(c) identification of the end point of the reaction
(d) appropriate experimental techniques
(e) care and accuracy in the use of equipment.

select an indicator on the basis of the relevant titration curve.
Chemistry 10: Acids and Bases
Fri 12 February 2016
1
3. Mandatory experiences

a quantitative study to determine the concentration of one of the reactants by
titration

a practical study of a variety of indicators.
4. Suggested experiences

a practical investigation of the production of an acid or base

a practical investigation of the pH in soils, colour of pigments, food and
cosmetics.
Key Points


Acids and bases can be defined in
slipperiness, turns litmus, red/purple
terms of observations or concepts
Conceptual definitions were proposed in
pH = -log10[H+]
the 19th century.
PROPERTIES OF ACIDS AND BASES
ARRHENIUS ACIDS AND BASES
Definitions of acids and bases based on
Svante Arrhenius proposed a conceptual
their characteristic properties are known
definition of acids as containing
as operational definitions. Eg:
positive H+ ions.


An Arrhenius acid is a substance
solution.
which provides H+ ions in aqueous
HCl –> H+ + ClACID + BASE –> SALT + WATER
A water soluble base is a
substance which produces
hydroxide ions as its only negative
Problems with Arrhenius’ definition
ions in solution.
Some compounds which do not contain
NaOH –> Na+, OH-
hydroxide ions are able to neutralise
acids.
Common Acids
Nitric
Sulfuric
Acetic / Ethanoic
Carbonic
Common Bases
Ammonium hyd.
Barium hyd.
Calcium hyd.
Formula
HCl
HNO3
H2SO4
BRØNSTED-LOWRY ACIDS AND
BASES
Brønsted-Lowry definitions:
H3CO3
An acid is a proton donor.
Formula
A base is a proton acceptor.
NaOH
KOH
NH4OH
CONJUGATE PAIRS
Acid solutions are formed when an acid
transfers a proton to a water molecule.
E.g.
Neutralisation and formation of salts
Chemistry 10: Acids and Bases
Fri 12 February 2016
HCl + H2O –> Cl- + H3O+ (aq)
Since the proton is transferred from an
acid to a base, acids and bases must exist
DIPROTIC AND POLYPROTIC ACIDS
in pairs called conjugate pairs.
Acids can be monoprotic, diprotic,
e.g.
HCN + H2O  CN- + H3O+
triprotic,... etc
eg
Conjugate pair 1:
HCN and CNConjugate pair 2:
H2O and H3O+
HCl
Monoprotic
H2SO4
Diprotic
H3PO4
Triprotic
If all 3 acids were 1 M concentration the
triprotic acid would have the highest
An acid-base reaction is a transfer of a
[H3O+] and hence, lowest pH.
proton from an acid to a base.
It depends on the number of H+ ions
donated.
The difference in a conjugate pair
between the acid and the base is 1
proton.
Note the number of H+ ions when doing
calculations. A diprotic acid has twice the
concentration of H+ ions.
Conjugate pairs
Acid
Base
Strengths of acids and bases
Strong Acid
HCl
Cl-
The strength of an acid is measured by
Weak Acid
CH3COOH
NH4+
CH3COONH3
the concentration of H+ ions. Since many
acid-base reactions are in equilibrium the
strength of the acid is determined by the
From the above table it can be seen that a
value of the equilibrium constant.
strong acid is the conjugate of a weak
base. Similarly a strong base will have a
weak acid as its conjugate


The more electronegative the atom
AMPHIPROTIC SPECIES
with H+ is, the more likely the
= double life
molecule will dissociate fully,
Can act as either L-B acid or L-B base,
therefore the higher Ka value.
depending on what it is put with.
HCO3- + H3O+ –> H2CO3 + H2O:
e.g HCN is not as electronegative
as HCl. Therefore HCl ionises full
acts as L-B base.
HCO3- + OH- –> CO32- + H2O
and HCN only partially. (The H in
HCN can share more of the
acts as L-B acid.
electron than with the HCl, i.e.
more tightly held.).
Chemistry 10: Acids and Bases
Amphiprotic Molecules
Fri 12 February 2016
[H+] = 10-pH
[H3O+] = 1 x 10-14
Water
Acts as a base if put with an acid, eg
H2SO4. Conjugate acid/base pair:
[OH-]
H3O+/H2O.
Acts as an acid if put with a base, eg
NH3.Conjugate acid/base pair H2O/OH-.
e.g.
The Ammonia Molecule
[Ba(OH-)2] = 0.05 M
Acts as a base when put with water, eg
Therefore [OH-] = 2 x 0.05 M = 0.1 M
NH3 + H2O –> NH4+ + OH- . Conjugate
pH = 13
acid/base pair: NH4/NH3.
Finding pH of Weak Acids
Acts as an acid when put with OH-.
1. Using Ka values
THE pH SCALE
2. Given % ionisation
Pure water: pH = 7
3. Given pH, calculate Ka.
-14
Kw = 1 x 10
1. Using Ka Values
pH = -log10 [H+]


Equation (Ka = 7 x 10-4
Ka Expression
7 x 10-4 =
Use fudge factor (since x is very small)
x2
HF + H2O  H3O+ + FInitial 0.1
Final

-
=0.1-x
-
-
x
x
0.1
2. Given % ionisation
0.1
Q. 0.01 M is 4% ionised
Equation
3. Given pH, Calculate Ka
0.01 x 0.04 = 0.0004
Q. pH 3.8, 0.1 M
[H+] = 1.6 x 10-4
pH = 3.4

Ka Expression
pH range 1 - 13 (equal: Equivalence pt.
at pH 7)
TITRATION CURVES
Shape:
Strong acid vs. strong base

First part: pH stays low because
is good for titration: 6-8 best, 5-9
OH- ions from base are neutralised
OK. Jump in middle

by H3O+ ions, but there is an
excess of H3O+ which keeps pH
high pH. Flat curve
Strong Acid/Weak base
low. Flat curve

Neutral: Any indicator in this range
Chemistry 10: Acids and Bases
Last part: OH- in excess, therefore
Fri 12 February 2016

pH range: 1 - 10 ( Equivalence pt.
OH- ions are in equilibrium with
shifted down)
XOH. (Therefore pH higher)

Steeper curve (pH >7) because

pH range: 4 - 13 (Equivalence pt.
Weak Acid/Strong base
shifted up)

Has steeper angle (pH < 7) as less
H+ ions are in solution and more
are in HX solution. (Equilibrium
reaction).
PH of Salt Solutions
Some salt solutions have a pH greater or
less than 7 because ions present can
react with the water.
Eg
NH4Cl + H2O –> NH3 + H3O+ + ClCO32- + H2O –> HCO3- + OHHPO42- + H2O –> OH- + H2PO4-
Buffer Solutions
Buffers usually consists of a solution of
weak acid in the presence of one of its
salts
eg Ammonia solution and ammonium
Chloride
NH3 + H2O  OH- + NH4+
NH4Cl –> NH4+ + ClSee workbook for more explanation
Q.
1. What happens when acid base is
added?
2. What is the salt for?
eg2
CH3COOH + H2O  CH3COO- + H3O+
Na+CH3COO- --> Na+ + ClChemistry 10: Acids and Bases
Fri 12 February 2016