Acids, bases and alkalis
Acids react with metals, bases, and carbonates.
$$\text{acid}+\text{metal} \rightarrow \text{salt}+\text{hydrogen}$$
$$\text{acid}+\text{base} \rightarrow \text{salt}+\text{water}$$
$$\text{acid}+\text{carbonate} \rightarrow \text{salt}+\text{water}+\text{carbon dioxide}$$
A base is usually a metal oxide or hydroxide that neutralises an acid. An alkali is a soluble base. Bases react with acids and ammonium salts:
$$\text{base}+\text{ammonium salt} \rightarrow \text{salt}+\text{water}+\text{ammonia}$$
In aqueous solution, acids contain $\mathrm{H^+}$ ions and alkalis contain $\mathrm{OH^-}$ ions. Acids are proton donors; bases are proton acceptors. Strong acids such as hydrochloric acid dissociate completely, while weak acids such as ethanoic acid dissociate partially.
Indicators and pH
| Indicator | Acidic solution | Alkaline solution |
|---|---|---|
| Litmus | Red | Blue |
| Thymolphthalein | Colourless | Blue |
| Methyl orange | Red | Yellow |
Universal indicator gives a range of colours and allows comparison of pH. Red/orange indicates strong acidity, yellow indicates weaker acidity, green is approximately pH 7, and blue/purple indicates alkalinity. Lower pH means greater acidity and higher hydrogen-ion concentration. Higher pH means greater alkalinity and higher hydroxide-ion concentration. pH 7 is neutral.
Neutralisation is:
$$\mathrm{H^+(aq)+OH^-(aq) \rightarrow H_2O(l)}$$
Oxides
- Basic oxides, usually metal oxides such as $\mathrm{CuO}$ and $\mathrm{CaO}$, react with acids.
- Acidic oxides, usually non-metal oxides such as $\mathrm{SO_2}$ and $\mathrm{CO_2}$, react with bases.
- Amphoteric oxides such as $\mathrm{Al_2O_3}$ and $\mathrm{ZnO}$ react with both acids and bases to form salts and water.
Solubility rules
| Soluble | Insoluble |
|---|---|
| All $\text{Na}^+, \text{K}^+, \text{NH}_4^+, \text{NO}_3^-$ salts | All $\text{CO}_3^{2-}$ except $\text{Na}, \text{K}, \text{NH}_4$ |
| All $\text{Cl}^-$ except $\text{Pb}^{2+}, \text{Ag}^+$ | $\text{OH}^-$ except $\text{Na}, \text{K}, \text{NH}_4, \text{Ca}$ |
| All $\text{SO}_4^{2-}$ except $\text{Ba}^{2+}, \text{Ca}^{2+}, \text{Pb}^{2+}$ |
Preparing salts
Insoluble salts by precipitation
- Choose two soluble solutions that contain the required ions.
- Mix them to form a precipitate.
- Filter the mixture.
- Wash the residue with distilled water.
- Dry the solid.
Example:
$$\mathrm{AgNO_3(aq)+NaCl(aq) \rightarrow AgCl(s)+NaNO_3(aq)}$$
Net ionic equation:
$$\mathrm{Ag^+(aq)+Cl^-(aq) \rightarrow AgCl(s)}$$
Soluble salts
- Acid and alkali: use titration because both are soluble. Add indicator, record the titre, repeat until consistent results obtained, then repeat without indicator using the mean titre. Evaporate the solution until nearly saturated, cool it to crystallise the salt, filter the crystals, wash them with a little cold distilled water, and dry them between filter papers.
- Acid and excess insoluble metal, base, or carbonate: add the solid until no more reacts, filter off excess solid, evaporate the solution until nearly saturated, cool it, filter the crystals, wash them with a little cold distilled water, and dry them.
Use a fume cupboard and appropriate eye protection when preparing salts from corrosive acids or when gases are released.
Hydrated salts
A hydrated salt contains chemically combined water, for example $\mathrm{CuSO_4\cdot5H_2O}$ or $\mathrm{CoCl_2\cdot6H_2O}$. An anhydrous salt contains no water of crystallisation. Heating a hydrated salt can remove this water.