Chemistry of Aqueous Solutions 🌊

Define Arrhenius acid

1. Arrehenius acid is a substance that ionises in water to produce a solution containing H+ ions.

Define Arrhenius base

1. Arrhenius base is a substance that ionises in water to produce a solution containing OH- ions.

Define Bronsted-Lowry acid

1. Bronsted-Lowry acid is a proton donor.

Define Bronsted-Lowry base

1. Bronsted-Lowry base is a proton acceptor.

Define Lewis acid

1. Lewis acid is an electron pair acceptor (electrophile).

2. They have vacant and energetically accessible orbitals to accept a pair of electrons from a base to form a dative bond.

Define Lewis base

1. Lewis base is an electron pair donor (nucleophile).

Define pH

1. pH is defined as the negative logarithm to base 10 of hydrogen ion concentration.

Define pOH

1. pOH is defined as the negative logarithm to base 10 of hydroxide ion concentration.

State when the effect of auto-ionisation of water has to be considered

1. If [H+] or [OH-] is very low e.g. 10-6 moldm-3, the autoionisation of water must be considered when calculating pH of strong acids or bases.

2. [H+] from water= 1.0 × 10-7 moldm-3.

3. [OH-] from water= 1.0 × 10-7 moldm-3.

Define degree of dissociation

1. Degree of dissociation of a weak acid or base is the fraction of the acid or base that has dissociated into ions at any particular dilution.

Explain why the pH at equivalence point between a strong acid (HCl) and strong base (NaOH) is 7

1. HCl + NaOH—> NaCl + H2O.

2. NaCl—> Na+ + Cl-.

3. Na+ does not undergo salt hydrolysis with water. Na+ has low charge density and hence low polarising power and does not hydrolyse water.

4. Cl- is the conjugate base of strong acid HCl. Cl- does not undergo salt hydrolysis with water as it is a weaker base than water.

5. Therefore, [H+]= [OH-]= 1 × 10-7 moldm-3 and pH> 7.

Explain why the pH at equivalence point between a strong acid (HCl) and weak base (NH3) is less than 7

1. HCl + NH3–> NH4Cl.

2. NH4Cl—> NH4+ + Cl-.

3. NH4+ is the conjugate acid of weak base NH3. NH4+ undergoes salt hydrolysis with water as it is a stronger acid than water.

4. NH4+ + H2O (reversible reaction) NH3 + H3O+.

5. Cl- is the conjugate base of strong acid HCl. Cl- does not undergo salt hydrolysis with water as it is a weaker base than water.

6. Therefore, [H+]> 10-7 moldm-3 and pH< 7.

Explain why the pH at equivalence point between a weak acid (CH3CO2H) and strong base (NaOH) is more than 7

1. CH3CO2H + NaOH—> CH3CO2Na + H2O.

2. CH3CO2Na—> CH3CO2- + Na+.

3. CH3CO2- is the conjugate base of weak acid CH3CO2H. CH3CO2- undergoes salt hydrolysis with water as it is a stronger base than water.

4. Na+ does not undergo salt hydrolysis with water. Na+ has low charge density and hence low polarising power and does not hydrolyse water.

5. CH3CO2- + H2O (reversible reaction) CH3CO2H + OH-.

6. Therefore, [OH-]> 1 × 10-7 moldm-3 and pH> 7.

Explain why the pH at equivalence point between a weak acid (HCO2H) and weak base (NH3) depends upon the extent to which ion hydrolysis

1. HCO2H + NH3–> HCO2NH4

2. HCO2NH4–> HCO2- + NH4+.

3. HCO2- is the conjugate base of weak acid HCO2H. HCO2- undergoes salt hydrolysis with water as it is a stronger base than water.

4. HCO2- + H2O (reversible reactionHCO2H + OH-.

5. NH4+ is the conjugate base of weak bae NH3. NH4+ undergoes salt hydrolysis with water as it is a stronger acid than water.

6. NH4+ + H2O (reversible reaction) NH3 + H3O+

7. Therefore, if Ka(NH4+)> Kb(HCO2-), [H+]> 1 × 10-7 moldm-3 and pH< 7.

Define buffers

1. Buffers are solutions that resist changes in pH upon additions of small amounts of acid or alkali.

Explain how acidic and alkaline buffers are prepared

1. Acidic buffers are prepared by mixing a weak acid and salt of weak acid and strong base.

2. Alkaline buffers are prepared by mixing a weak base and salt of weak base and strong acid.