Acid Reactions and Salt Preparation

Predicting the products of acid reactions

Acids can be used to make salts: ionic compounds containing positive metal ions, or ammonium ions, and negative ions derived from an acid. Choosing the reactants determines which salt is made. The type of substance reacting with the acid determines the other products.

For metal salts, the first part of the name comes from the metal. The second part comes from the acid:

For example, zinc reacting with sulfuric acid produces zinc sulfate. Magnesium carbonate reacting with nitric acid produces magnesium nitrate. Naming the salt is only the first step: we must also identify any gas or water formed.

Acids reacting with metals

Some metals react with dilute hydrochloric acid or dilute sulfuric acid to produce a salt and hydrogen:

metal + acid → salt + hydrogen

Metals above hydrogen in the reactivity series, such as magnesium, zinc and iron, can undergo these reactions. Copper is below hydrogen and does not react with these dilute acids in this way. More reactive metals generally react more vigorously under the same conditions. Sodium and potassium are too reactive for this salt-preparation method because their reactions with acids are dangerous.

Zinc reacting with dilute sulfuric acid gives zinc sulfate and hydrogen:

$$\mathrm{Zn + H_2SO_4 \rightarrow ZnSO_4 + H_2}$$

Magnesium reacting with dilute hydrochloric acid gives magnesium chloride and hydrogen:

$$\mathrm{Mg + 2HCl \rightarrow MgCl_2 + H_2}$$

The hydrogen escapes as bubbles, while the metal is gradually used up. The coefficient 2 before $\mathrm{HCl}$ means that two units of hydrochloric acid are needed for each magnesium atom. It balances both the hydrogen and chlorine atoms.

Acids reacting with carbonates

Metal carbonates react with acids to form a salt, carbon dioxide and water:

metal carbonate + acid → salt + carbon dioxide + water

The carbon dioxide causes effervescence, meaning bubbling as gas escapes. This distinguishes a carbonate reaction from the reaction of an acid with a metal oxide or hydroxide, which produces salt and water without a gas.

For calcium carbonate and hydrochloric acid, the metal supplies the name calcium and the acid supplies chloride:

$$\mathrm{CaCO_3 + 2HCl \rightarrow CaCl_2 + CO_2 + H_2O}$$

Magnesium carbonate and nitric acid instead produce magnesium nitrate:

$$\mathrm{MgCO_3 + 2HNO_3 \rightarrow Mg(NO_3)_2 + CO_2 + H_2O}$$

These equations conserve atoms: there is the same number of each type of atom on both sides. The products differ from those of metal–acid reactions because the reactant is a carbonate compound, not the metal itself.

Deducing formulae from ion charges

An ion is an atom or group of atoms with an electrical charge. An ionic compound contains positive and negative ions in proportions that make its overall charge zero. Its formula gives the simplest whole-number ratio of those ions.

Useful ions for these reactions include:

To deduce a formula, write the ions and their charges, then find the smallest numbers of each ion that give equal total positive and negative charges.

• Magnesium chloride: one $\mathrm{Mg^{2+}}$ needs two $\mathrm{Cl^-}$ ions. The formula is $\mathrm{MgCl_2}$.

Brackets keep a compound ion together when more than one is needed. For example, ammonium sulfate is $\mathrm{(NH_4)_2SO_4}$: two ammonium ions balance one sulfate ion. The final compound formula has no overall charge written on it.

Preparing a soluble salt

Making a salt solution is not the same as obtaining pure, dry salt crystals. A useful preparation method is to react an acid with an insoluble solid, add enough solid to use up all the acid, and then remove the excess solid by filtration.

Copper(II) sulfate can be prepared from copper(II) oxide and dilute sulfuric acid:

copper(II) oxide + sulfuric acid → copper(II) sulfate + water

$$\mathrm{CuO(s) + H_2SO_4(aq) \rightarrow CuSO_4(aq) + H_2O(l)}$$

Copper(II) oxide is a black insoluble solid. As it reacts, a blue copper(II) sulfate solution forms. The state symbol $(aq)$ means dissolved in water.

Reacting and filtering

Wear eye protection. Avoid contact with the acid and copper compounds, and handle hot equipment carefully. Use gentle heating to reduce the risk of splashing; dispose of copper-containing materials as instructed by the teacher.

1. Measure a volume of dilute sulfuric acid into a beaker. Gently warm it using suitable heating apparatus, then remove the heat before adding the solid.
2. Add small spatula portions of copper(II) oxide, stirring with a glass rod after each addition. Warming speeds up the reaction, while stirring brings the reactants into contact.
3. Continue adding and stirring until some black solid remains without reacting. This persistent excess shows that the available acid has been used up.
4. Pour the mixture carefully through filter paper in a funnel into a conical flask. Keep the liquid below the top edge of the paper so that solid cannot bypass the filter.

The excess copper(II) oxide stays on the paper as the residue. The blue copper(II) sulfate solution passes through as the filtrate. Keep this filtrate: dissolved salt passes through filter paper, so filtration cannot remove it from the water.

Evaporating, crystallising and drying

Transfer the filtrate to an evaporating basin. Heat it gently over a water bath to evaporate some of the water. Evaporation removes solvent and makes the salt solution more concentrated.

Stop heating when the solution is close to saturation, with crystals beginning to appear at the edge. Do not boil it dry: strong heating can cause spitting and damage the crystals.

Leave the solution to cool. As it cools, less copper(II) sulfate can remain dissolved, so solid crystals form. Separate the crystals from the remaining liquid by filtration, then gently blot them with filter paper and leave them to air-dry.

Each stage has a different job: reaction makes the salt, the first filtration removes excess reactant, evaporation removes some water, crystallisation forms solid salt, and drying removes liquid from the crystals. This excess-solid method works because the unused reactant is insoluble and can be filtered out. A soluble alkali cannot be removed this way; an accurately measured acid–alkali reaction requires the titration approach.

Reaction patterns

• Metal + dilute hydrochloric or sulfuric acid → salt + hydrogen. Suitable metals include magnesium, zinc and iron; copper does not react in this way.
• Carbonate + acid → salt + carbon dioxide + water. Carbon dioxide causes effervescence.
• Metal oxide + acid → salt + water.

Naming salts