Acids, Bases and Buffers

Module 1 of F325

5.1.3 Acids, Bases and Buffers

Bear with me this is my least favourite topic!

  • Created by: Samantha
  • Created on: 23-04-14 10:49

Acids and Bases

Bronsted-Lowry Acid: a proton, H+, donor

Bronsted-Lowry Base: a proton, H+, acceptor

1 of 26

Strong and Weak Acids

Strong Acids

An acid that completely dissociates in solution

e.g. HCl, HNO3, H2SO4, HBr, HI, HClO4 (chloric (VII))

Weak Acids

An acid that partially dissociates in solution

CH3COOH(aq) ↔ H+(aq) + CH3COO-(aq)

  • Small concentrations of H+(aq) and CH3COO-(aq) compared with CH3COOH(aq)
  • Equilibrium position to the left
2 of 26

The Role of H+ in the Reactions of Acids

Acid + Water

The acid dissociates, releasing H+ ions into the solution

e.g. HCl(g) + aq → H+(aq) + Cl-(aq)

Different acids release different numbers of protons. Depends on their formulae & bonding.

HCl is monobasic. Each molecule releases one proton

H2SO4 is dibasic. Each molecule releases 2 protons

H2SO4 (aq) → H+ (aq) + H2SO4- (aq)

HSO4(aq) ↔ H+ (aq) + (SO4)2- (aq)

H3PO4 is tribasic. Each molecule releases 3 protons

3 of 26

The Role of H+ in the Reactions of Acids

Acid-Base Reactions

Aqueous acids take part in typical acid-base reactions with carbonates, bases and alkalis

Acid neutralised

Water formed as one of the products

H+ responsible for reactions

Ions only dissociated in solution

Solid ionic compounds shown undissociated

Only cancel species that do not change

Release of H+ may be incomplete. As the H+ react, the acid will release more.

Reaction stops when all H+ have been released

4 of 26

The Role of H+ in the Reactions of Acids

Acid + Carbonate 

2HCl(aq) + CaCO3(s) → CO2(g) +CaCl2(aq) + H2O(l)

Ionic Equation

2H+(aq) + 2Cl-(aq)CaCO3(s) → CO2(g) + Ca2+(aq) + 2Cl-(aq) +H2O(l)

Cancel species that do not change (including physical states)

This means cancel species that occur on both sides of the arrow that have the same state. In this case it means 2Cl-

2H+(aq) + CaCO3(s) → CO2(g) + Ca2+(aq) + H2O(l)

If the carbonate is in solution (CaCO3(aq)) the final ionic equation will be even simpler because the carbonate is dissociated and its ions cancel.

5 of 26

The Role of H+ in the Reactions of Acids

Acid + Base

2HCl(aq) + CuO(s) → CuCl2(aq) + H2O(l)

Ionic Equation

2H+(aq) + 2Cl-(aq) + CuO(s) → Cu2+(aq) + 2Cl-(aq) + H2O(l)

Cancel species that do not change (including physical states)

2H+(aq) + CuO(s) → Cu2+(aq) + H2O(l)

6 of 26

The Role of H+ in the Reactions of Acids

Acid + Alkali

HCl(aq) + NaOH(aq) → H2O(l) + NaCl(aq)

Ionic Equation

H+(aq) + Cl-(aq) + Na+(aq) + OH-(aq) → H2O(l) + Na+(aq) + Cl-(aq)

Cancel species that do not change (including physical states)

H+(aq) + OH-(aq) → H2O(l)

7 of 26

The Role of H+ in the Reactions of Acids

Acid + Metal

This is a redox reaction

It does not follow the typical acid-base reaction structure

2HCl(aq) + Mg(s) → MgCl2(aq) + H2(g)

Ionic Equation

2H+(aq) + 2Cl-(aq) + Mg(s) → Mg2+(aq) + 2Cl-(aq) + H2(g)

Cancel species that do not change (including physical states)

2H+(aq) + Mg(s) → Mg2+(aq) + H2(g)

Water is not formed as one of the products

8 of 26

Conjugate Acid-Base Pairs

Acids will only release a proton if there is a base to accept it

Hydronium Ions, H3O+

HCl(aq) + H2O(l) → H3O+(aq) + Cl-(aq)



9 of 26

Conjugate Acid-Base Pairs

Acid-base pairs are a set of two species that transform into each other by gain or loss of a proton.

HNO3 (acid) ↔ H+ + NO2-(base)

Acid-base equilibria involve two acid-base pairs

HNO2(aq) + H2O(l)  H3O+(aq) + NO2-(aq)

Forward Reaction:

The acid HNO2(aq) releases a proton to form its conjugate base NO2-(aq)

The base H2O(l) accepts the proton form its conjugate acid H3O+(aq)

Reverse Reaction:

The acid H3O+(aq) releases a proton to form its conjugate base H2O(l)

The base NO2-(aq) accepts the proton to form its conjugate acid HNO2(aq)

10 of 26

The Acid Dissociation Constant

The weak acid HA has the following equilibrium in solution:

HA(aq) ↔ H+(aq) + A-(aq)

The Acid Dissociation Constant:

Ka = [H+(aq)][A-(aq)] / [HA(aq)]

The terms 'stong' and 'weak' describe the extent of dissociation of an acid given by the Ka value:

  • Large Ka value = large amount of dissociation, therefore the acid is strong
  • Small Ka value = small amount of dissociation, therefore the acid is weak
11 of 26

Ka and pKa

Ka is made more manageable by converting to pKa

pKa = -log10Ka

Ka = 10^-pKa

Low Ka = High pKa

High Ka = Low pKa

Smaller pKa = stronger acid

12 of 26


pH = -log[H+(aq)]

[H+(aq)] = 10^-pH

The difference between each successive whole number pH value on the pH scale is a factor of 10

e.g. the difference between pH 3 and pH 5 is a factor of 10 x 10 = 100

A ph change of 1 changed [H+(aq)] by 10x

An acid with pH of 2 contains 1000x the [H+(aq)] of an acid with pH of 5

13 of 26


Low pH = Large [H+(aq)]

High pH = small [H+(aq)]

14 of 26

The Ionic Product of Water

In the ionisation of water: 

H2O(l) ↔ H+(aq) + OH-(aq)  The equilibrium lies to the left

Kc = [H+(aq)] [OH-(aq)] / [H2O(l)]

Kc x [H2O(l)] = [H+(aq)] [OH-(aq)]

[H2O(l)] = concentration of water, constant value

Kc and [H2O(l)] are both constants which are multiplied together to give Kw, the ionic product of water

15 of 26

The Ionic Product of Water

 Kcontrols the balance between [H+(aq)and [OH-(aq)] in all aqueous solutions

At 25°C the pH value of 7 is the neutral point at which [H+(aq)and [OH-(aq)] are the same and equal to 10^-7 moldm^-3

This applies to water and neutral solutions

In water and neutral solutions: [H+(aq)] = [OH-(aq)] 

In acidic solutions: [H+(aq)[OH-(aq)]

In alkaline solutions:  [H+(aq)[OH-(aq)]

We can find the concentrations of H+ and OH- in solutions with different pH values

Use [H+(aq)[OH-(aq)] = 1.00 x 10^-14 mol^2 dm^-6

16 of 26

Buffer Solutions

A buffer solution is a mixture that minimizes pH changes on addition of small amounts of acid or base

It is a mixture of a weak acid, HA, and a conjugate base, A-

Can be made from a weak acid and the salt of a weak acid e.g. ethanoic acid CH3COOH and sodium ethanoate CH3COONa. A weak acid can be partially neutralised by an aqueous alkali to give a solution of the salt and the excess of weak acid.

In the CH3COOH / CH3COONa system:

  • the weak acid partially dissociates, CH3COOH(aq) ↔ H+(aq) + CH3COO-(aq)
  • the salt completely dissociates, generating the conjugate base, CH3COO-,    CH3COONa(aq) → CH3COO-(aq) + Na(aq)

The equilibrium mixture contains a high concentration of the undissociated weak acid, CH3COOH(aq), and its conjugate base, CH3COO-(aq). The high concentration of the conjugate base pushes the equilibrium to the left, so there is a low concentration of H+ ions.

17 of 26

How A Buffer Acts

The weak acid and the conjugate base control pH

HA(aq) ↔ H+(aq) + A-(aq) 

The weak acid removes added alkali. The conjugate base removes added acid.

Addition of acid:

[H+(aq)] is increased. The conjugate base, A-, reacts with the H+ ions. The equilibrium shifts to the left, removing most of the added H+ ions.

Addition of alkali:

[OH-(aq)] is increased. A small concentration of H+ ions already in the solution react with the OH- ions.

H+(aq) + OH-(aq) → H2O(l) 

 HA dissociates, shifting the equilibrium to the right to restore most of the H+ ions that have reacted

18 of 26

pH and [H+(aq)] for Strong Monobasic Acids

Nearly complete dissociation in water. [H+(aq)] of a strong acid = [HA(aq)]

pH is calculated using pH = -log[H+(aq)]

HCl, concentration 1.22 x 10^-3 mol dm^-3

What is its pH?

HCl is a strong acid-it completely dissociates

HCl(aq) →  + Cl-(aq)

[H+(aq)] = [HCl(aq)] = 1.00 x 10^-3 mol dm^-3

pH = -log[H+(aq)]

pH = -log(1.22 x 10^-3)

pH = 2.91

19 of 26

pH and [H+(aq)] for Weak Monobasic Acids

Weak monobasic acids partially dissociate:

HA(aq) ↔ H+(aq) + A-(aq)

Ka = [H+(aq)] [A-(aq)] / [HA(aq)]

At equilibrium: [H+(aq)] = [A-(aq)] so [H+(aq)] [A-(aq)] = [H+(aq)]^2

[HA(aq)] - [H+(aq)] as only a few HA molecules have dissociated

Ka[H+(aq)]^2 / [HA(aq)] - [H+(aq)]

Assume [HA(aq)] at equilibrium = [HA(aq)] at the start

= [H+(aq)]^2 / [HA(aq)]

[H+(aq)] = √([HA(aq)] x Ka)

pH = -log[H+(aq)]

20 of 26

Ka for a Weak Acid

We need:

  • pH
  • concentration

0.030mol dm^-3 HCOOH

pH 2.66

Calculate Ka

[H+(aq)] = 10^-pH

= 10^-2.66

= 2.19 x 10^-3 mol dm^-3

Ka = [H+(aq)]^2 / [HA(aq)]

= (2.19 x 10^-3) / 0.030 = 1.6 x 10^-4 mol dm^-3

21 of 26

pH of a Strong Base

We need [H+(aq)] which depends on the concentration of the base and the ionic product of water, Kw

Strong base: NaOH(aq)  Na+(aq) + OH-(aq)

Completely dissociated in aqueous solution

[OH-(aq)] = [base(aq)] so [OH-(aq)] = [NaOH(aq)]

Kw = [H+(aq)][OH-(aq)] so [H+(aq)] = Kw / [OH-(aq)]

pH = -log[H+(aq)]

A base is a proton acceptor. The strength of a base is a measure of its dissociation to generate OH- ions.

Bases that dissociate in water to release hydroxide ions are called alkalis. Strong bases are alkalis.

22 of 26

pH of a Buffer Solution

For a buffer consisting of a weak acid, HA and its conjugate base A-

 Ka = [H+(aq)] [A-(aq)] / [HA(aq)]

[H+(aq)] = Ka x ([HA(aq)] / [A-(aq)])

Only a small proportion of HA dissociates:

[HA(aq)] equilibrium = [HA(aq)] undissociated

Salt of the weak acid is ionic. Dissociates completely in aqueous solution.

CH3COO-Na+(aq)      CH3COO-(aq) + Na+(aq)

CH3COO-(aq) = CH3COO-Na+(aq)

[H+(aq)] used to calculate pH

pH = -log [H+(aq)]

23 of 26

The Carbonic Acid-Hydrogencarbonate Buffer System

H2CO3(aq)       H+(aq) + HCO3-(aq)

Carbonic acid is the weak acid
Hydrogencarbonate is the weak base

An increase in H+ ions is removed by HCO3-(aq)
The equilibrium shifts left

An increase in OH- ions is removed by H2CO3(aq)
A small concentration of H+ ions reacts with OH- ions
H+(aq) + OH-(aq)      H2O(l)
H2CO3 dissociates further
Equilibrium shifts right

Most materials released into the blood are acidic. HCO3- ions remove these by being converted to H2CO3. This is converted to aqueous CO2 by an enzyme. The dissolved CO2 is converted to CO2(g) in the lungs and exhaled. Amount of CO2 controlled by the rate of breathing

[H+(aq)] = 10^-7.40 = 3.98 x 10^-8 mol dm^-3
(Ka / [H+(aq)]) / ([HCO3-(aq)] / [H2CO3(aq)])
(4.3 x 10^-7) / (3.98 x10^-8) = 10.8 : 1 

24 of 26

Acid-Base Titration pH Curves

Equivalence point: the point in a titration at which the volume of one solution has reacted exactly with the volume of the second solution.

An acid-base indicator is a weak acid. It can be represented by Hln. It has one colour in its acid form (Hln) and another colour in its conjugate base form (ln-). Methyl orange is red in its acid form and yellow in its conjugate base form.

When there are equal amounts of weak acid and conjugate base [Hln] = [ln-], the indicator is at its end point. The colour of methyl orange at its end point is orange.

Strong acid-Strong base: phenolpthalein, methyl orange and bromothymol blue are suitable indicators
Strong acid-Weak base: Methyl orange a suitable indicator
Weak acid-Strong base: phenolpthalein suitable
Weak acid-Weak base: no indicators used in practice

A suitable indicator changes colour within the pH range of the vertical section of the curve

25 of 26

Enthalpy Change of Neutralisation

Standard Enthalpy Change of Neutralisation: the enthalpy change that accompanies the neutralisation of an aqueous acid by an aqueous base to form one mole of H2O(l), under standard conditions.

HCl(aq) + NaOH(aq)      NaCl(aq) + H2O(l)

H+(aq) + OH-(aq)      H2O(l)

Na+ and Cl- are spectator ions

26 of 26


No comments have yet been made

Similar Chemistry resources:

See all Chemistry resources »See all Acids, bases and salts resources »