How to write ionic equations

1. Ionic equations: overhearing reactions underwater

Many reactions in the O Level syllabus happen in water, like neutralisation and precipitation. To sound cheem, we say that they occur in aqueous solutions. In this set of notes, we will learn how to describe the underwater dramas with ionic equations.

H⁺(aq) + OH^–(aq) ⟶ H₂O(l)

Ionic equations are a special type of chemical equations. Like the ionic equation for neutralisation shown above, they give the limelight to the dissolved ions that actually undergo chemical changes during aqueous reactions.

2. Dissolving ionic compounds, dissociating ions

In the solid state, the ions of an ionic compound are held closely together by ionic bond. We usually describe these forces as “strong”. Yet, for soluble ionic compounds, their toughness is not without weakness. When we dissolve them in water, the water molecules have a special ability to overcome the electrostatic forces of attraction, separating the ions. The separated ions are mobile, mixing around with the water molecules.

The cheem word that we use for separate is dissociate. For example, we say that when sodium chloride dissolves in water, the sodium ions and chloride ions dissociate. In ionic equations, we will rewrite NaCl (aq) as Na⁺ (aq) and Cl^– (aq) separately, to emphasise that the ions are dissociated.

Dissociation is the separation of ions when a soluble ionic compound dissolves in water.

On the other hand, insoluble ionic compounds do not dissolve in water. This means that their ions are not dissociated.

3. The basics of writing ionic equations: knowing your solubility table

In ionic equations, we express dissolved ionic compounds as their dissociated ions, like Na⁺ (aq) and Cl^– (aq). This means that we have to know what is soluble, and what is not.

Compounds of…	Solubility Rules (of Thumb)
Group I ions: Sodium ion, Na+ Potassium ion, K+	All are soluble
Ammonium ion, NH4+	All are soluble
Nitrate ion, NO3–	All are soluble
Halide ions*, X–: Chlorides, Cl– Bromides, Br– Iodides, I–	All are soluble, except: Silver halides, AgX Lead(II) halides, PbX2
Sulfate ions, SO42-	All are soluble, except: Barium sulfate, BaSO4 Calcium sulfate, CaSO4 Lead(II) sulfate, PbSO4
Carbonates, CO32-	All are insoluble, except: Sodium carbonate, Na2CO3 Potassium carbonate, K2CO3 Ammonium carbonate, (NH4)2CO3
Hydroxide, OH–	All are insoluble, except: Sodium hydroxide, NaOH Potassium hydroxide, KOH

These are rules of thumb. The actual solubility depends on temperature and the relative volume of water used. Obviously, even if sodium chloride is soluble, we cannot dissolve 1000 kg of it in 1 small drop of water.

Nevertheless, they are still useful in helping us to predict the solubility of common ionic compounds under normal circumstances. If the table above feels overwhelming, just remember that S.P.A.N. (sodium, potassium, ammonium, or nitrate-containing) compounds are always soluble.

4. Translating chemical equations into ionic equations

We can write ionic equations by starting from their chemical equations. Then, we rewrite the soluble ionic compounds as their dissociated ions. This is when you realise that some ions do not react: they remain dissociated in solution. These are the spectator ions, which we cancel out. Therefore, the net ionic equation will show the actual chemical change, without the spectator ions.

Ionic Equations for Neutralisation

Let’s demonstrate this by writing the ionic equation for the neutralisation of hydrochloric acid with sodium hydroxide.

• STEP 1: Write the chemical equation

  HCl(aq) + NaOH(aq) ⟶ NaCl(aq) + H₂O(l)

• STEP 2: Rewrite by separating the soluble ionic compounds into their dissociated ions

  H⁺(aq) + Cl^–(aq) + Na(aq) + OH^–(aq) ⟶ Na⁺(aq) + Cl^–(aq) + H₂O(l)

  Note that water is not separated because it is not an ionic compound. As a simple molecule, the atoms do not separate and remain bonded as a discrete molecule of H₂O.

  However, HCl is separated. While pure HCl (g) is a simple molecule, it is an acid that undergoes ionisation in water to form hydrogen ions. Therefore, dissolved HCl (aq) should be treated as a soluble ionic compound, with dissociated ions.

• STEP 3: Cancel out common ions, which are the spectator ions

  H⁺(aq) + Cl^–(aq) + Na⁺(aq) + OH^–(aq) ⟶ Na⁺(aq) + Cl^–(aq) + H₂O(l)

  Sodium and chloride ions remain dissolved in the solution. They are the spectator ions, that watch and do not undergo the chemical reaction themselves.

• STEP 4: Write the net ionic equation for neutralisation

  H⁺(aq) + OH^–(aq) ⟶ H₂O(l)

Ionic Equations for Precipitation

Let’s practise again by looking at the precipitation of insoluble silver iodide from silver nitrate and sodium iodide.

• STEP 1: Write the chemical equation

  AgNO₃(aq) + NaI(aq) ⟶ AgI(s) + NaNO₃(aq)

  As silver iodide is insoluble, it precipitates as a solid. Therefore, its state symbol is (s) instead of (aq).

• STEP 2: Rewrite by separating the soluble ionic compounds into their dissociated ions

  Ag⁺(aq) + NO₃^–(aq) + Na(aq) + I^–(aq) ⟶ AgI(s) + Na⁺(aq) + NO₃^–(aq)

  Note that AgI is not separated because it is an insoluble ionic compound.

• STEP 3: Cancel out common ions, which are the spectator ions

  Ag⁺(aq) + NO₃^–(aq) + Na(aq) + I^–(aq) ⟶ AgI(s) + Na⁺(aq) + NO₃^–(aq)

  Sodium and nitrate ions remain dissolved in the solution. They are the spectator ions, that watch and do not undergo the chemical reaction themselves.

• STEP 4: Write the net ionic equation for precipitation

  Ag⁺(aq) + I^–(aq) ⟶ AgI(s)

5. Practise writing ionic equations for common reactions

Use the method above to practise writing ionic equations for the common aqueous reactions that are tested at the O Level.

Do check your answer to see if:

• The equation is balanced, whereby the number of atoms of each element remains the same before and after the reaction
• The charges are balanced, whereby the total charge on the left is equal to the total charge on the right
• State symbols are included
• Spectator ions are cancelled out