Notes

How to Write Ionic Equations

Peer into the underwater drama of aqueous reactions with 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

Ions of sodium chloride dissociate in water

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, NO3All 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, OHAll are insoluble, except:
Sodium hydroxide, NaOH
Potassium hydroxide, KOH
*The solubility rule for halides only apply to chlorides, bromides and iodides, but not fluorides.

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
Write the ionic equation for the acid-metal reaction between zinc and sulfuric acid to form zinc sulfate salt and hydrogen gas.

Chemical equation: Zn(s) + H2SO4(aq) ⟶ ZnSO4(aq) + H2(g)

Ionic equation: Zn(s) + 2H+(aq) ⟶ Zn2+(aq) + H2(g)

Each formula unit of H2SO4 dissociates to give 2 hydrogen ions, H+, and 1 sulfate ion, SO42-.

Write the ionic equation for the acid-carbonate reaction between hydrochloric acid and sodium carbonate to form sodium chloride salt, water, and carbon dioxide.

Chemical equation: 2HCl(aq) + Na2CO3(aq) ⟶ 2NaCl(aq) + H2O(l) + CO2(g)

Ionic equation: 2H+(aq) + CO32-(aq) ⟶ H2O(l) + CO2(g)

Sodium carbonate and sodium chloride are soluble sodium-containing ionic compounds. They dissolve in water and their state symbol is (aq).

Write the ionic equation for the acid-carbonate reaction between hydrochloric acid and insoluble magnesium carbonate to form magnesium chloride salt, water, and carbon dioxide.

Chemical equation: 2HCl(aq) + MgCO3(s) ⟶ MgCl2(aq) + H2O(l) + CO2(g)

Ionic equation: 2H+(aq) + MgCO3(s) ⟶ Mg2+(aq) + H2O(l) + CO2(g)

Magnesium carbonate is insoluble. Therefore, it does not dissociate and we write it as MgCO3(s).

Write the ionic equation for the neutralisation reaction between sodium hydroxide and sulfuric acid to form sodium sulfate salt and water.

Chemical equation: H2SO4(aq) + 2NaOH(aq) ⟶ Na2SO4(aq) + 2H2O

Ionic equation: H+(aq) + OH(aq) ⟶ H2O(l)

You might have gotten 2H+(aq) + 2OH(aq) ⟶ 2H2O(l). This is technically correct but not simplified to the lowest ratio. Simplify by dividing the coefficients by 2 to get H+(aq) + OH(aq) ⟶ H2O(l).

Write the ionic equation for the precipitation reaction between calcium chloride solution and sodium carbonate solution.

Chemical equation: CaCl2(aq) + Na2CO3(aq) ⟶ CaCO3(s) + 2NaCl(aq)

Ionic equation: Ca2+(aq) + CO32-(aq) ⟶ CaCO3(s)

All carbonates are insoluble except sodium carbonate, potassium carbonate, and ammonium carbonate. In other words, the calcium carbonate produced is insoluble. It forms as a solid precipitate that does not dissolve or dissociate. We write it as CaCO3(s).

Write the ionic equation for the redox reaction between copper and nitric acid to form copper nitrate, nitrogen dioxide, and water.

Chemical equation: Cu(s) + 4HNO3(aq) ⟶ Cu(NO3)2(aq) + 2NO2(g) + 2H2O(l)

Ionic equation: Cu(s) + 4H+(aq) + 2NO3(aq) ⟶ Cu2+(aq) + 2NO2(g) + 2H2O(l)

This is a tricky redox reaction. For every 4 nitrate ions involved, 2 nitrate ions take part in the reaction by oxidising copper to copper(II) ion. The other 2 nitrate ions are instead the spectator ions. Also, unlike other acid-metal reactions, the oxidising agent is not hydrogen ions but the nitrate ions.