5.1.3: Acids, Bases and Buffers Flashcards
Bronsted Lowry acid
a species that is a proton donor
Bronsted Lowry base
a species that is a proton acceptor
Why use the pH scale rather than simply [H⁺]
- [H⁺] deals with negative indices over a very wide range
* pH scale makes numbers used to measure acid concentration more manageable
Strong alkali
A fully ionic soluble base
Weak Bronsted Lowry acid
a proton donor that only partially dissociates
How to calculate pH strong acid?
1) [HA] = [H⁺] because strong acid fully dissociates
2) pH = -log₁₀[H⁺]
How to calculate pH weak acid?
NEED Ka
1) Ka = [H⁺] [A⁻] / [HA]
2) Ka = [H⁺]² /[HA]
3) Sub in Ka and [HA]
4) pH = -log₁₀[H⁺]
Assumptions when calculating pH weak acid?
- [H⁺] = [A⁻] i.e. water dissociation does not increase the [H⁺]
- [HA] is unaffected by dissociation, because so little dissociation occurs i.e. you keep all the [HA] you started with
- The Ka value is small, approx. 1x10⁻⁵
Calculating pH of strong base
1) Kw = [H⁺] [⁻OH] = 1 x 10⁻¹⁴
2) [H⁺] = 1 x 10⁻¹⁴ / [⁻OH]
3) pH = -log₁₀[H⁺]
pKa =
–log₁₀Ka
Kw is known as
the ionic product of water
Calculating pH of diluted solution
1) Find [H⁺] in original solution (acid)
2) [H⁺] x (new volume / total volume )
3) pH = -log₁₀[H⁺]
Buffer
a solution of a weak acid and its conjugate base. Resists change in pH on addition of small quantities of acid/alkali
What happens when you add H⁺ to a buffer?
- There is a large reservoir of A⁻
* So H⁺ reacts and pH doesn’t change much
What happens when you add ⁻OH to a buffer?
- ⁻OH reacts with H⁺ to give H₂O
- There is a large reservoir of HA, so equilibrium will shift to replace H⁺
- The alkali can also react with HA
- So the HA decreases, leading to equilibrium shift to replace the HA
