Q4. Give the conjugate base of each: H2O, NH4+, H2CO3, HS−.
OH−, NH3, HCO3− and S2−: one proton fewer each time, and one unit more negative charge.
Q5. In HCO3−(aq) + OH−(aq) ⇌ CO32−(aq) + H2O(l), is HCO3− acting as an acid or a base? And in HCO3−(aq) + H3O+(aq) ⇌ H2CO3(aq) + H2O(l)?
In the first it gives a proton to OH−, so it is an acid (conjugate base CO32−). In the second it takes a proton from H3O+, so it is a base (conjugate acid H2CO3). It is amphiprotic.
Q6. Is every Brönsted base a Lewis base? Is every Brönsted acid a Lewis acid?
Yes to the first: to accept a proton a base must supply an electron pair (usually a lone pair). Proton transfer can be described as electron-pair donation to the acidic proton, so introductory texts often include Brönsted acidity within Lewis acidity. Distinguish this proton-centred description from formation of a Lewis adduct with the intact acid molecule: HCl donates H+, while that proton accepts the base’s electron pair. And BF3, AlCl3 and Mg2+ are Lewis acids with no proton at all.
Q7. Write the Brönsted acid–base reaction between HNO3 and water, and name the conjugate pairs.
HNO3(aq) + H2O(l) → H3O+(aq) + NO3−(aq). Pairs: HNO3/NO3− and H3O+/H2O. HNO3 is a strong acid, so NO3− is a very weak base and the reaction is practically complete.
Q8. Why do water’s high dielectric constant and its lone pairs both matter when an acid such as HCl ionises?
Water acts as the Brönsted base, using an oxygen lone pair to take the proton and form H3O+; its high dielectric constant then weakens the attraction between H3O+ and Cl−, and hydration keeps the ions apart.