ACIDS & BASES
Acids and Bases reactions occur in everyday life
and are essential for understanding our world.
How does pH value affect our environment?
Why is it important to monitor and maintain the pH of the water
in aquariums, soil and our blood?
What exactly is pH? How is it measured?
Milk of magnesia is a medicine
that usually relieves
uncomfortable gastrointestinal
symptoms within 30 minutes
and constipation within six
hours.
Why is the milk of magnesia
an antacid?
Keywords
Acidity
 Basicity (Monoprotic, diprotic, triprotic)
 Bronsted-Lowry Theory
- Proton donor/acceptor
- Acid-base Conjugate pair
- Amphiprotic
 Lewis Theory
- Lone pair electrons
- Dative/Coordinate bond

What is an acid?
A
solution that contains __________ ions (protons).
OLD THEORY
Weak acid like ethanoic acid does not have
the power to neutralise strong acid like
sodium hydroxide.
What is a base/alkali?
•
•
•
A base is a substance like __________and
______________that reacts with acid to form salt and water
only.
An alkali is a soluble base which in solution produces
________ ions.
Most bases are insoluble in water. 3 soluble bases are
NaO/NaOH, KO/KOH, CaO/Ca(OH)2
Both acids and alkalis are _____________.
What causes acidity?
• It is the __________________that give an acid its acidic
properties when they dissolve in water and ____________
into ions.
E.g. HCl gas is a covalent compound.
When dissolves in water, it forms HCl acid which dissolciate
to form ions.
What is basicity (proticity)?
Basicity
• refers to the no.of _____ atoms in one molecule of acid that
can be replaced by a _________.
• refers to the no. of _______that can be replaced by one
molecule of that acid.
E.g. HCl (monobasic),
H2SO4(dibasic),
H3PO4(tribasic)
Bronsted-Lowry theory
An acid is defined as a molecule or ion that acts as a
proton __________.
A base is defned as a molecule or ion that acts as a
proton __________.
Types of acids
Acids that have single proton to donate –
___________ (monobasic).
E.g. HCl(aq), HNO3(aq), HNO2(aq)
 Acids that have 2 protons to donate – __________
E.g. H2SO4(aq), H2SO3(aq), H2CO3(aq)
 H3PO4(aq) is ___________.

Hydrogen chloride gas dissolved in water (solvent)
HCl(g) + H2O(l)
H3O+(aq) + Cl-(aq)
The equation can be split into
(i) HCl(aq)
Cl-(aq)
acid
conjugate base
(ii) H2O(l) + H+(aq)
base
+
H+(aq)
H3O+(aq)
conjugate acid
Acidic behaviour is a transfer reaction in different solvents.
Acid-base conjugate pair
CH3COOH(l) + H2O(l)
acid
base
donates H+
NH3(g) + H2O(l)
H3O+ (aq) + CH3OOO-(aq)
conjugate
conjugate
acid
base
+
donates H
NH4+(aq) + OH-(aq)
Water is sometimes described as _______________ because
it can accept or donate a proton.
Competition between
acid/base and its conjugate
(i) HCl(g) + H2O(l)
acid
base
(ii) CH3COOH(l) + H2O(l)
acid
(i)
(ii)
base
H3O+(aq) + Cl-(aq)
conjugate acid
conjugate base
H3O+ (aq) + CH3OOO-(aq)
conjugate acid
conjugate base
Water is a much stronger base than chloride ion and has a stronger
tendency to accept proton.The equilibrium shifts more to the right.
Ethanote ion is a much stronger base than water molecule. The equilbrium
shifts to the left.
Gas-phase acid-base reaction
HCl(g) + NH3(g)
NH4Cl(s)
The Bonsted-Lowry model can be extended to gasphase acid-base reaction.
 It involves the transfer
of hydrogen ion from
hydrogen chloride to
ammonia.



Strong acids have weak conjugate bases.
Weak acids have strong conjugate acids.
(i) HCl(g) + H2O(l)
acid
base
conjugate acid
(ii) CH3COOH(l) + H2O(l)
acid
H3O+(aq) + Cl-(aq)
base
conjugate base
H3O+ (aq) + CH3OOO-(aq)
conjugate acid
conjugate base
If HA is a strong acid in water,
+
 HA is a successful donor of H in water
 the reverse reaction hardly happens
 A- is a poor acceptor of H+
 Ka (dissociation constant) is big
Equilibrium lies to the right. Strong acid , weak conjugate base
HA + H2O
H3O + + A-
Weak acid , strong conjugate base. Equilibrium lies to the left
Common acids & conjugate bases
in order of strengths
Lewis theory


A Lewis acid is defined as a substance that can
accept a pair of _________ from another atom to
form a ________ covalent bond.
A Lewis base is defined as a substance that can
__________ a pair of electrons to another atom to
form a dative (coordinate) covalent bond.
B:
Lewis _____
H+  +BH
Lewis ______
Examples

Reaction between ammonia, NH3 and proton

H3N:
H+  +NH4
Reaction between NH3 and BF3.
H
H
H
F
B
N
F
H
 H
F
H
F
B
N
F
F
BF3 is a good Lewis ______ as there are _______electrons around the central
boron atom which leaves room for 2 more electrons.
Other common Lewis acid includes AlCl 3 and transition metal ions in aqueous
solution.

Reaction between a water molecule and proton
H2O:
H+  H3O+
Lewis bonding
In complex ions formed by transition metals
The 6 water molecules, each donate a lone pair electrons from
oxygen of their water molecules to the empty 3d orbitals of
iron.
What does each water molecule and iron(III) ion act
as in the reaction above?
Dative (Coordinate) bond


A dative covalent bond is always formed in a Lewis
acid-base reaction.
For a substance to act as a base, it must have space
to accept the lone pair of electrons.
Strong and weak acids and bases
Strong acid
 When strong acid (HA)
dissolves, virtually all acid
molecules react with the
water to produce
hydronium ions (H3O+).
HA + H2O(l)  H3O+(aq) + A-(aq)
0%
100%
Examples : HCl, H2SO4,HNO3, HClO4
or
HA  H+(aq) + A-(aq)
0% 100%
Strong and weak acids and bases
Weak acid
 When a weak acid dissolves
in water, only a small % of its
molecules (typically 1%) react
with water molecules to
release hydrogen or
hydronium ions. The
equilibrium lies on the lefthand side of the equation.
HA + H2O(l)
99%
H3O+(aq) + A-(aq)
1%
or
HA
H+(aq) + A-(aq)
99% 1%
Examples : CH3COOH, aqueous carbon dioxide
Distinguish between strong and weak acids
0.1 mol dm-3 HCl(aq)
0.1 mol dm-3 CH3COOH (aq)
[H+(aq)]
0.1 mol dm-3
0.0013 mol dm-3
pH
1.00
2.89
Electrical conductivity
high
low
Relative rate of reaction with
magnesium
fast
slow
Relative rate of reaction with
calcium carbonate
fast
slow
Base on the information above, how do we distinguish between strong
and weak acids of the same concentration (e.g. HCl and
CH3COOH)?
How to distinguish between strong and weak
acids?




A weak acid has a lower concentration of __________ and
hence a higher _____ than a stronger acid of the same
concentration.
A weak acid, because of its lower concentration of hydrogen
ions, will have much poorer _____________ than a stronger
acid of the same concentration.
Weak acids react more _______ with reactive metals, metal
oxides, metal carbonates and metal hydrogencarbonates than
strong acids of the same concentration.
Strong and weak acids can also be distnguished by measuring
and comparing their enthalpies of neutralisation.
What is the difference between the strength (strong and weak)
and the concentrated (concentrated or dilute)?
Strong and weak acids and bases
Strong base
 A strong base undergoes almost 100% dissociation/ionisation
when in dilute aqueous solution.
BOH B+(aq) + OH-(aq)
0%
100%
Examples : NaOH, KOH, Ba(OH)2
Strong and weak acids and bases
Weak base
 All bases are weak except the hydroxides of groups 1 and 2.
 Weak bases are composed of molecules that react with water
molecules to release hydoxide ions. In general for a weak
molecular base, BOH

BOH + (aq)
B+(aq) + OH-(aq)
The equiibrium lies on the left side of the equation.
Examples : aqueous ammonia, ethylamine, caffeine, bases
of nuclei acids
The pH indicator




scale that measures the strength of an acid and
alkali.
pH of a substance is measured when it is dissolved
in water.
pH stands for “power of hydrogen”
[H+] = 1 x 10-n moldm-3 ( n = pH number)
The pH Scale
pH probe and meter
An accurate method of measuring pH value.
A pH probe is dipped into the solution being tested
and the pH value is then read directly from the
meter.
pH Calculation


pH is a measure of the concentration of H+ ions in a
solution.
pH = -log10[H+(aq)]
Example:
If the concentration of H+ is 2.50 x 10-3 moldm-3 , what is
the pH?
pH = -log (2.50 x 10-3)
= 2.60
Example:
Calculate the concentration of H+ of a solution that
has a pH = 3.2.
-log[H+] = 3.2
log [H+] = -3.2
[H+] = 6.31 x 10-4
Example:
Calculate the concentration of H+ and hence the pH
of a 1.00 x 10-3 moldm-3 NaOH
[OH-] = 1.00 x 10-3 moldm-3
[H+] x OH-] = 1.00 x 10-14
[H+] =
Example:
(a) What is the pH of 10cm3 of 0.1 moldm-3 HCl?
pH = -log (0.1) = 1
(b) If 90cm3 of water is added to the acid, what happens to the
pH?
Total volume = 100cm3
In 10cm3 solution, concentration of H+ is 0.1 moldm-3
In 100cm3 solution, concentration of H+ is 0.01 moldm-3
pH = -log (0.01) = 2
(c) If the solution from (b) is diluted by a factor of 105 , what is the
approximate pH?
The pH will increase to 6.
Buffer

A buffer resists changes in _____ when small
amounts of acid and alkali are added to it.
Acidic Buffer

An acidic buffer solution can be made by mixing a
weak ______ together with the _______ of the
weak acid and a strong _______.
(1) CH3COOH(aq)
H +(aq) + CH3COO-(aq)
(2) CH3COONa(aq)
Na+(aq) + CH3COO-(aq)
Acidic Buffer

If an acid is added, the extra H+ from the acid
react with the excess ethanoate ions in (2) and are
removed from the solution as ethanoic acid
molecules (these have no effect on the pH). Hence
the pH stays the same.
CH3COO-(aq) + H +(aq)
CH3COOH(aq)
new
Acidic Buffer

If an alkali is added, the OH- from the alkali react
with the hydrogen ions from (1) removing them from
the right hand side. There is, however, a large
reservoir of ethanoic acid on the left hand side of
this equilibrium able to dissociate and make more
hydrogen ions, restoring the pH.
CH3COOH(aq)+OH -(aq) CH3COO-(aq)+H2O(l)
Acidic Buffer
Acidic buffers are often made by taking a solution of
a strong base and adding excess weak acid to it, so
that the solution contains the salt and the unreacted
weak acid.
CH3COOH(aq)+NaOH(aq) CH3COONa(aq)+ H2O(l) +CH3COOH(aq)
limiting agent
salt
excess weak acid
buffer
Alkali Buffer
An alkali buffer with a fixed pH greater than 7 can
be made from a weak base together with the salt
of the base with a strong acid.
 E.g. Ammonia and ammonium chloride
NH4Cl(aq)  NH4+(aq) + Cl-(aq)
NH3(aq) + H2O(l)
NH4+(aq) + OH-(aq)

Alkali Buffer
(1) NH4Cl(aq)  NH4+(aq) + Cl-(aq)
(2) NH3(aq) + H2O(l)
NH4+(aq) + OH-(aq)
If H+ ions are added they will combine with OH- ions from (2)
to form water and more of the ammonia will dissociate to
replace them
Adding more OH- ions that can react with the free ammonium
ions (from equilibrium 1) producing more ammonia (as in
equilibrium 2) and effectively being removed from the
system. The ammonia molecules have no effect on pH an
therefore the pH remains the same.
In both cases, the hydroxide ion concentration and the
hydrogen ion concentration remain constant.