CHAPTER 17

Equilibria involving acids and bases

Acidity of Solutions

Lowry-Brønsted theory Acids are proton donors Bases are proton acceptors.

An acid base reaction involves the transfer of a proton from an acid to a base.

Some substances can both donate and accept protons, these are amphiprotic solutions.

The ionisation constant of water

Water will react with itself in a process called self-ionisation:H2O(l) + H2O(l) H3O+(aq) + OH-(aq)

Write the equilibrium constant for this reaction.

The concentration of water is usually fairly constant so we can write the equilibrium law the ionisation of water as:

Kw = [H3O+][OH-]

Where Kw is the ionisation of water

The ionisation constant of water

This law applies to both pure water and all aqueous solutions.

In pure water at 25°C, chemists have found that the concentration of both H3O+ and OH- ions is 10-7M.

The value of Kw at 25°C can be calculated:

Kw = [H3O+][OH-] = 10-7 x 10-7 = 1.0 x 10-14 M2

Acidic and Basic Solutions

In acidic solutions the concentration of H3O+ ions will be greater than 10-7 M at 25°C.

The opposite is true for basic solutions.In summary at 25°C:

In pure water and neutral solutions: [H3O+] = [OH-] = 10-

7 M In acidic solutions: [H3O+] > 10-7 M and [OH-] < 10-7 M

In basic solutions: [H3O+] < 10-7 M and [OH-] > 10-7 M

The higher the concentration of hydronium ions in a solution the more acid the solutions is.

pH

pH = -log10[H3O+]

[H3O+] = 10-pH

Remember for basic solutions we are given the OH- concentration

Can anyone remember how we can convert this to a H3O+

concentration?Lets look at the worked

examples on page 287

How is pH affected by temperature?

Remember:Kw = [H3O+][OH-] = 10-7 x 10-7 = 1.0 x 10-14 M2 at 25°CBut what happens if the temperature is not 25°CIonisation of water is endothermicSo:

As the temperature increases the equilibrium constant increases

This causes Kw to rise which results in a decrease in pH

If the temperature decreases it favours a reverse reaction which is exothermic Kw will decrease and the pH will increase

Your Turn

Page 288Questions 1 - 3

Acidity Constants

For the reactionHCl(aq) + H2O(l) H3O+(aq) + Cl-(aq)The expression for the equilibrium constant can be

written as:

Water is the solvent in aqueous solutions and its concentration is virtually constant.

So we can remove water from the equilibrium constant and call it Ka instead

K =[H3O+][Cl-][HCl][H2O]

Acidity Constant

Ka is the acidity constant.

The value of Ka for HCl is 107 M at 25°C.This means that in HCl solutions, most of the

HCl has been converted to product.This is why HCl is classified as a strong acid.

Ka =[H3O+][Cl-]

[HCl]

Acidity Constants

These ideas can be generalised to solutions of any weak acid represented by HA:

HA(aq) + H2O(l) A-(aq) + H3O+(aq)

[H3O+] = [A-] and [HA] does not change during the ionisation

Calculations Involving Acidity Constants

Calculate the pH and percentage hydrolysis of a 0.50 M ethanoic acid solution, given that the Ka for ethanoic acid is 1.75 x 10-5 M

Your Turn

Page 292Question 4 and 5

Buffers: Using equilibrium to resist change

Buffers are solutions that can absorb the addition of acids or bases with little change of pH.

They can be made most easily by mixing a weak acid and a salt of its conjugate base.

Buffers are especially important in environmental and living systems where they maintain a delicate chemical balances essential to life.

Consider the equation for ethanoic acid in water as a buffer. What happens to the equilibrium when you add a strong

acid or a strong base???

pH in the body

A number of reactions that occur in the body involve acid-base reactions.

Without a means of controlling acidity the pH of body fluids could fluctuate from extremely basic to extremely acid.

Our bodily systems can only operate in a small range.

Turn to page 292 for an example.

Your Turn

Page 292Question 6End of Chapter questions are due