Self-Ionisation of Water
Kw is the dissociation constant of water
Kw = [H3O+][OH-] = 1 x 10^-14 at 298K
pH
pH is defined as the negative logarithm to the base 10 of the hydrogen ion concentration.
pH = -log[H+]
Prove that the pH of water is 7
Pure water is neutral >>> equal amounts of H3O+ and OH-. In 1L of pure water, there are 10^-7 moles each of H3O+ and OH- at 298K.
- [H3O+] = [OH-] = 10^-14 = x
- x² = 10^-14
- x= 10^-7
- [H+] = 10^-7
- -log(-7) = -(-7) = 7
Calculate the pH of 0.03 H2SO4
0.03 x 2 = 0.06 (x2>>>2 hydrogens)
-log(o.o6) = 1.22
pH of Bases
pH + pOH = 14
Find pH of 0.1M NaOH
pOH = -log(0.1) = 1
pH = 14-1 = 13
Solutions of Weak Acids and Bases
Solutions of weak acids and bases do not fully dissociate and therefore react:
AH + H2O <—-> A- + H3O+
Ka is the dissocation constant
Ka = [A-][H3O+]
[AH]
Remember the stronger the acid, the greater the value of Ka.
pH of a weak acid = -log(√Ka x Ma) [Ma = molarity]
Calculate the pH of a 0.001M solution ethanoic acid, Ka = 1.8 x 10 ^ -5
pH = -log(√1.8 x 10^-5 x 0.001) = 3.87
Acid-Base Indicators
Acid-base indicator is a weak acid ot base that has a different colour when it is dissociated into its ions than when it is undissociated.
When neutralisation occurs in a titration, pH changes slowly until the reaction is almost complete when a rapid change occurs.
Indicators change colour in a certain pH range.
Strong Acid V Strong Base
- Change: 4-10
- All indicators
Weak Acid V Storng Base
- Change: 7-10
- Phenolphthalein
Strong Acid V Weak Base
- Change: 4-7
- Methyl Orange
Weak Acid V Weak Base
- No sharp end
- No suitable indicator
Acid Indicator
HIn +H2O <—> In- + H3O+ (HIn is indicator)
Add:
- Acid – reaction goes from right to left
- Base – reaction goes from left to right
Base Indicator
XOH <—-> X+ + OH- (X is base)
Add:
- Base – reaction goes from right to left
- Acid – reaction goes from left to right
Remember adding same substances (Acid to acid or base to base) the reaction will go to the left
