- Bronsted-Lowry acids are proton donors
- Bronsted-Lowry bases are proton acceptors
Strong and Weak Acids
- Strong acids dissociate (or ionise) almost completely in water as nearly all of the H+ ions are released
e.g. HCl(g) + H2O(l) à H+(aq) + Cl–(aq)
- Strong bases dissociate (or ionise) almost completely in water too
e.g. NaOH(s) + H2O(l) à Na+(aq) + OH– (aq)
- In both Strong Bases and Acids, the equilibrium lies over to the right
- Weak acids dissociate minimally in water, as only a few of the H+ ions are released
CH3COOH(aq) ⇋ CH3COO–(aq) + H+(aq)
- Weak bases dissociate minimally in water:
NH3(aq) + H2O(l) ⇋ NH4++ OH–(aq)
- In both weak Bases and Acids, the equilibrium lies over to the left
Buffer Solution
- Buffers are solutions which can resist changes in acidity or alkalinity
- When a small volume of acid/alkali is added to them their pH remains at a constant
- They are based on an equilibrium reaction which will move in the direction to remove either additional hydrogen ions or hydroxide ions
Acidic Buffers
- Acidic buffers are made from weak acids
- The dissociation of a weak acid is an equilibrium reaction
HA(aq) ⇋ H+(aq) + A–(aq)
Acid Buffer: adding alkali
- If a little alkali is add, the OH– ions react with the HA to produce water molecules and A‑:
HA(aq) + OH–(aq) ⇋ H2O(aq) + A–(aq)
- This removes the added OH–so the pH trends to remain almost the same
Acid Buffers: adding acid
- If H+ is added, the equilibrium will shift to the left as the H+ ions combine with the A– ions to produce undissociated HA.
Calculations of Buffers
- Different buffers can be made which will maintain different pH’s. when weak acid dissociates:
HA(aq) ⇋ H+(aq) + A–(aq)
- The expression of this being:
[H+(aq)][A–(aq)]
[HA(aq)]
- Using this expression, the pH of the buffer can be calculated
pH Scale
- The acidity of a solution varies on the concentration of H+ (aq) and is measured along the pH scale
pH = -log10[H+(aq)]
Calculating Hydrogen Ion Concentration from pH
To find the hydrogen ion concentration from a given pH then the inverse of the pH formula is used: [H+] = 10-pH
Calculating Hydroxide Ion Concentration from pH
- The Concentration of OH– ions is proportional to the concentration of the base
- However, to find the pH the [H+] must be known, therefore the ionic product of water(Kw) is used Kw = [H+][OH–]
- If the [OH–] for a strong aqueous base and Kw at a certain temperature are both known then the [H+] can be determined and therefore the pH
Strong Monoprotic Acids
- Monoprotic refers to each molecule of an acid which will release 1 proton when it is dissociated
- Examples of monoprotic acids are Hydrochloric and Nitric acid
Strong Diprotic Acids
- Diprotic acids are those which release 2 protons when they dissociate
- Examples of diprotic acids are Sulfuric acid and carbonic acid
Indicators for Titrations
- The equivalence point is where there is an equal proportion of hydrogen ion added to hydroxide ions
- The End point is the volume of an alkali or acid which when added together the indicator changes colour
- Strong acids and Strong Base
- Methyl Orange and Phenolphthalein changes colour within the equivalence point and therefore suitable
- Phenolphthalein Weak Acid and Strong Base
- Phenolphthalein
- Strong Acid and Weak Base
- Methyl Orange will change sharply at the equivalence point
- Weak Acid and Weak Bases
- Neither indicator is suitable
- No indicator would be viable for the equivalence point over the two pH units
Calculating pH of a weak acid using Ka: The Dissociation of Weak acids
- Weak acids do not ionise fully in a solution and therefore the [H+] is not proportional to the acid concentration
- If a weak acid of HA is dissociated:
HA(aq) ⇋H+(aq) + A–(aq)
- The equilibrium constant (Kc) is formed:
Kc = [H+(aq)][A–(aq)]
[HA(aq)]
- For a weak acid, this is usually given the Ka symbol known as the acid dissociation constant
Ka = [H+(aq)][A–(aq)]
[HA(aq)]
- The greater the value of Ka the further the equilibrium is to the right, as more acid which is dissociated and the stronger it is
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