Laboratory Experiment 3.1*. Ion Exchange Reactions

Lecture

Reagents: solutions of sodium hydroxide, sulfuric acid, sodium chloride, sodium carbonate, iron(III) chloride, potassium nitrate, phenolphthalein.

Carry out reactions between the following pairs of substances:

1) sodium hydroxide (in the presence of phenolphthalein) and sulfuric acid;
2) iron(III) chloride and sodium hydroxide;
3) sulfuric acid and sodium carbonate;
4) potassium nitrate and sodium chloride;
5) between iron(III) hydroxide, obtained in experiment 2, and sulfuric acid.

Write the equations of the reactions carried out in molecular form, and in full and net ionic forms. Indicate the effects you observed.

Draw a conclusion about the necessary conditions for ion exchange reactions to occur.

Theoretical background

Ion exchange reactions are chemical processes in solutions of electrolytes in which substances exchange ions. They proceed only when a precipitate, a gas, or a weak electrolyte (for example, water) is formed as a result. Otherwise, exchange reactions are reversible.

Definition

An ion exchange reaction is an interaction between solutions of acids, bases, and salts in which ions change places.
An important feature: the oxidation state of the elements does not change, since these are not oxidation-reduction reactions, but purely exchange reactions.

Conditions for occurrence

For a reaction to proceed, at least one condition must be met:

Formation of a precipitate (for example, Fe(OH)₃).
Evolution of a gas (for example, CO₂ upon the interaction of carbonates with acids).
Formation of a weak electrolyte (for example, H₂O upon the neutralization of an acid and a base).

If none of the conditions is met (for example, mixing solutions of NaCl and KNO₃), the reaction does not proceed.

Laboratory Experiment 3.1*. Ion Exchange Reactions

Forms for writing reactions

Ion exchange reactions are conventionally written in three forms:

Molecular form — shows the substances in full. Example:

H_{2} S O_{4} + N a_{2} C O_{3} \to N a_{2} S O_{4} + H_{2} O + C O_{2} ↑

Full ionic form — all strong electrolytes are written out as ions.

2 H^{+} + S O_{4}^{2 -} + 2 N a^{+} + C O_{3}^{2 -} \to 2 N a^{+} + S O_{4}^{2 -} + H_{2} O + C O_{2} ↑

Net ionic form — only the particles that actually take part in the reaction remain.

2 H^{+} + C O_{3}^{2 -} \to H_{2} O + C O_{2} ↑

Rules for writing ion exchange reactions

When writing ionic equations, one must always be guided by the solubility table of acids, bases, and salts in water, that is, always check the solubility of the reagents and products, noting this in the equations.
Keep in mind that double exchange reactions between salts with the formation of precipitates proceed in all cases where the solubility of the reagents is higher than the solubility of one of the products.
To obtain a sparingly soluble substance, one should always choose highly soluble reagents and use sufficiently concentrated solutions.

Laboratory Experiment 3.1*. Ion Exchange Reactions

Signs of an ion exchange reaction

Formation of a precipitate (brown Fe(OH)₃).
Evolution of a gas (CO₂).
Change in indicator color (phenolphthalein becomes colorless upon neutralization).
Dissolution of a precipitate upon reaction with an acid.