The ionic bond

This page covers a single Form 4 content standard: the ionic bond. It follows directly from the idea that atoms bond to become stable, and it shows the first route to that stability, transferring electrons from a metal to a non-metal. This is one of the most heavily examined standards in the chapter, so you should be able both to define the ionic bond precisely and to show its formation with electron-arrangement or dot-and-cross diagrams.

What an ionic bond is

An ionic bond is the strong electrostatic force of attraction between oppositely charged ions. That exact wording matters in the exam. The bond does not join two atoms directly; instead, electrons are transferred so that one particle becomes positive and another becomes negative, and the opposite charges pull on each other.

Ionic bonding happens between a metal and a non-metal:

• The metal atom loses its valence electrons and becomes a positive ion (cation).
• The non-metal atom gains those electrons and becomes a negative ion (anion).

Both ions end up with a full outer shell, a stable octet (or duplet for very small atoms) like a noble gas. The number of electrons the metal loses is exactly the number the non-metal gains.

How the electrons transfer

Consider sodium and chlorine. Sodium has the arrangement 2.8.1 and chlorine has 2.8.7. Sodium loses its single valence electron to chlorine:

• Sodium: 2.8.1 → Na⁺ (2.8), an octet like neon.
• Chlorine: 2.8.7 → Cl⁻ (2.8.8), an octet like argon.

The resulting Na⁺ and Cl⁻ ions attract each other by electrostatic force, that attraction is the ionic bond, and the compound formed is sodium chloride, NaCl. In a dot-and-cross diagram you draw only the outer electrons: the single dot (sodium’s electron) is shown moving into chlorine’s outer shell, which is drawn with crosses, and the whole structure is written inside square brackets with the charge, [Na]⁺ and [Cl]⁻.

Getting the formula right by balancing charges

The formula of an ionic compound is fixed by making the total positive charge equal the total negative charge, so the compound is electrically neutral. For example:

• Sodium (forms Na⁺) with chlorine (forms Cl⁻): one of each → NaCl.
• Magnesium (forms Mg²⁺) with chlorine (forms Cl⁻): one Mg²⁺ needs two Cl⁻ → MgCl₂.
• Aluminium (forms Al³⁺) with oxygen (forms O²⁻): balance 3 and 2 → Al₂O₃.

Worked example

Question. Magnesium (electron arrangement 2.8.2) reacts with chlorine (2.8.7) to form magnesium chloride. Explain how the ionic bond forms, state the charge and arrangement of each ion, and give the formula of the compound.

Step 1, Count the valence electrons. Magnesium has 2 valence electrons; each chlorine atom has 7.

Step 2, Decide the transfer. Magnesium loses its 2 valence electrons. Each chlorine atom can gain only 1 electron to complete its octet, so one magnesium atom must give its 2 electrons to two chlorine atoms.

Step 3, Form the ions. Magnesium becomes Mg²⁺ (2.8), an octet like neon. Each chlorine becomes Cl⁻ (2.8.8), an octet like argon.

Step 4, Describe the bond and formula. The Mg²⁺ ion and the two Cl⁻ ions attract one another by strong electrostatic forces, these are the ionic bonds. Balancing one 2+ charge against two 1− charges gives the formula MgCl₂.

Answer. Magnesium transfers 2 electrons, one to each of two chlorine atoms, forming Mg²⁺ (2.8) and two Cl⁻ (2.8.8). The oppositely charged ions attract by electrostatic force, and the compound is MgCl₂.

Practice question

Question. Aluminium has the electron arrangement 2.8.3 and oxygen has 2.6. Explain how aluminium and oxygen form an ionic compound, and deduce its formula.

Answer. Aluminium has 3 valence electrons, so each aluminium atom loses 3 to form Al³⁺ (2.8). Oxygen has 6 valence electrons, so each oxygen atom gains 2 to form O²⁻ (2.8). To balance the charges, the lowest whole-number ratio that makes total positive equal total negative is 2 Al³⁺ to 3 O²⁻ (giving 6+ and 6−). The oppositely charged ions attract by strong electrostatic forces, forming the ionic compound Al₂O₃.

Exam tip

Two things earn the marks. First, learn the definition word for word: an ionic bond is the strong electrostatic force of attraction between oppositely charged ions, not “the sharing of electrons” (that is covalent) and not simply “the transfer of electrons” (that is how it forms, not what it is). Second, when you draw a dot-and-cross diagram, show only the outer electrons, use dots for one atom and crosses for the other so the transfer is visible, and write each ion in square brackets with its charge. Always state the charge and the new electron arrangement of each ion, and get the ratio by balancing charges.

Where this fits

The ionic bond is the first of the two main bond types in the Chemical Bond chapter, and it builds directly on the stability standard before it. It contrasts with the covalent bond, where electrons are shared rather than transferred, and it explains the properties of ionic compounds you meet later in the chapter. Practise drawing the diagrams with the worked examples and test your formula deductions with the practice questions. In our online 1-to-1 SPM Chemistry lessons, taught in English, from RM50/hr, our teachers make sure your definition is exact and your dot-and-cross diagrams show the transfer clearly, which is exactly what full marks on this standard require across SPM Chemistry.

Quick recap

• An ionic bond is the strong electrostatic force of attraction between oppositely charged ions.
• It forms between a metal (loses electrons → cation) and a non-metal (gains electrons → anion).
• Each ion reaches a stable octet or duplet like a noble gas.
• Draw dot-and-cross with outer electrons only, ions in square brackets with charge.
• Find the formula by balancing total positive and negative charge.