Ionic Bonding

This lesson covers the formation of ionic bonds as required by the Edexcel GCSE Chemistry specification (1CH0), Topic 1: Key Concepts in Chemistry — Structure, Bonding and Properties. You need to understand how ionic bonds form through the transfer of electrons and how to represent ionic compounds using dot-and-cross diagrams.

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What Is Ionic Bonding?

Ionic bonding is the strong electrostatic attraction between oppositely charged ions. It occurs between metals and non-metals.

• Metals are on the left of the periodic table — they have relatively few electrons in their outer shell and tend to lose electrons.
• Non-metals are on the right of the periodic table — they need only a few electrons to complete their outer shell and tend to gain electrons.

When a metal atom reacts with a non-metal atom:

1. The metal atom loses one or more electrons from its outer shell, forming a positive ion (cation).
2. The non-metal atom gains one or more electrons into its outer shell, forming a negative ion (anion).
3. Both ions now have a full outer shell — the same electronic structure as a noble gas.

Exam Tip: Always state that the bond is due to the "strong electrostatic attraction between oppositely charged ions." This is the definition the examiner is looking for.

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Why Do Atoms Form Ions?

Atoms are most stable when they have a full outer shell of electrons (like the noble gases). Atoms with incomplete outer shells will lose, gain or share electrons to achieve this stable arrangement.

• Group 1 metals (e.g. Na, K) lose 1 electron → form ions with a +1 charge (Na⁺, K⁺).
• Group 2 metals (e.g. Mg, Ca) lose 2 electrons → form ions with a +2 charge (Mg²⁺, Ca²⁺).
• Group 6 non-metals (e.g. O, S) gain 2 electrons → form ions with a -2 charge (O²⁻, S²⁻).
• Group 7 non-metals (e.g. Cl, Br) gain 1 electron → form ions with a -1 charge (Cl⁻, Br⁻).

Exam Tip: The charge on the ion is determined by the group number. Group 1 → +1, Group 2 → +2, Group 6 → -2, Group 7 → -1.

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Dot-and-Cross Diagrams for Ionic Compounds

Dot-and-cross diagrams show what happens to the outer-shell electrons when ionic bonds form. Electrons from the metal atom are shown as dots (•) and electrons from the non-metal atom are shown as crosses (×) — or vice versa. This allows you to track which atom each electron originally belonged to.

Example 1: Sodium Chloride (NaCl)

• Sodium (Na): Electronic structure 2, 8, 1 — has 1 electron in its outer shell.
• Chlorine (Cl): Electronic structure 2, 8, 7 — has 7 electrons in its outer shell.

What happens:

1. Sodium loses 1 electron → Na⁺ (electronic structure 2, 8).
2. Chlorine gains 1 electron → Cl⁻ (electronic structure 2, 8, 8).

Reading the diagram: Na has transferred its single outer electron (a dot •) onto Cl's outer shell, which now holds 8 electrons in total — 7 of its own (drawn as crosses ×) and 1 from sodium. Both ions now have full outer shells, matching the electronic configurations of the noble gases neon and argon respectively. The square brackets and charges (+ and −) indicate that each species is an ion, not a neutral atom.

The dot-and-cross diagram shows:

• Na⁺ with an empty outer shell (all electrons transferred), drawn in square brackets with a + charge.
• Cl⁻ with 8 electrons in its outer shell (7 crosses + 1 dot from sodium), drawn in square brackets with a - charge.

The formula is NaCl — one sodium ion to one chloride ion.

Example 2: Magnesium Oxide (MgO)

• Magnesium (Mg): Electronic structure 2, 8, 2 — has 2 electrons in its outer shell.
• Oxygen (O): Electronic structure 2, 6 — has 6 electrons in its outer shell.

What happens:

1. Magnesium loses 2 electrons → Mg²⁺ (electronic structure 2, 8).
2. Oxygen gains 2 electrons → O²⁻ (electronic structure 2, 8).

The dot-and-cross diagram shows:

• Mg²⁺ drawn in square brackets with a 2+ charge.
• O²⁻ with 8 electrons in its outer shell (6 crosses + 2 dots from magnesium), drawn in square brackets with a 2- charge.

The formula is MgO — one magnesium ion to one oxide ion.

Example 3: Calcium Chloride (CaCl₂)

• Calcium (Ca): Electronic structure 2, 8, 8, 2 — has 2 electrons in its outer shell.
• Chlorine (Cl): Electronic structure 2, 8, 7 — each chlorine has 7 electrons in its outer shell.

What happens:

1. Calcium loses 2 electrons → Ca²⁺ (electronic structure 2, 8, 8).
2. Each chlorine atom gains 1 electron → 2 × Cl⁻ (electronic structure 2, 8, 8 each).

Because calcium loses 2 electrons and each chlorine only needs 1, two chlorine atoms are needed for every one calcium atom.

The formula is CaCl₂.

Exam Tip: When drawing dot-and-cross diagrams, always put square brackets around each ion and write the charge outside the bracket. Don't forget the charge!

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The Giant Ionic Lattice

Ionic compounds do not exist as individual pairs of ions. Instead, the ions are arranged in a regular, repeating three-dimensional pattern called a giant ionic lattice.

• Each positive ion is surrounded by negative ions, and each negative ion is surrounded by positive ions.
• The strong electrostatic forces act in all directions between the oppositely charged ions.
• This arrangement maximises the attractive forces between oppositely charged ions and minimises the repulsive forces between ions of the same charge.

For example, in sodium chloride:

• Each Na⁺ ion is surrounded by 6 Cl⁻ ions.
• Each Cl⁻ ion is surrounded by 6 Na⁺ ions.
• The structure extends in all three dimensions.

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Formulae of Ionic Compounds

The formula of an ionic compound is determined by the need for the total positive charge to balance the total negative charge, so the overall compound is electrically neutral.

Metal Ion	Non-metal Ion	Formula	Reasoning
Na⁺	Cl⁻	NaCl	+1 and -1 balance
Mg²⁺	O²⁻	MgO	+2 and -2 balance
Ca²⁺	Cl⁻	CaCl₂	+2 needs 2 × (-1) to balance
Na⁺	O²⁻	Na₂O	2 × (+1) needed to balance -2
Mg²⁺	Cl⁻	MgCl₂	+2 needs 2 × (-1) to balance

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Worked Example: Lithium Fluoride (LiF)

Question: Describe the formation of lithium fluoride (LiF) from its elements using a dot-and-cross diagram.

Answer:

• Lithium (Li) has the electronic structure 2, 1 — it has 1 electron in its outer shell.
• Fluorine (F) has the electronic structure 2, 7 — it has 7 electrons in its outer shell.
• Lithium loses 1 electron → Li⁺ (electronic structure 2).
• Fluorine gains 1 electron → F⁻ (electronic structure 2, 8).
• The dot-and-cross diagram shows Li⁺ in square brackets with a + charge, and F⁻ with 8 electrons (7 crosses + 1 dot) in square brackets with a - charge.
• The ions are held together by strong electrostatic attraction in a giant ionic lattice.
• The formula is LiF.

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Key Points

• Ionic bonding occurs between metals and non-metals.
• Metals lose electrons to form positive ions (cations).
• Non-metals gain electrons to form negative ions (anions).
• Both ions achieve a full outer shell (noble gas configuration).
• An ionic bond is the strong electrostatic attraction between oppositely charged ions.
• Ions are arranged in a giant ionic lattice — a regular, repeating 3D structure.
• The formula of an ionic compound reflects the ratio of ions needed for overall electrical neutrality.

Exam Tip: A very common 6-mark question asks you to "Describe, with the aid of a diagram, the bonding in sodium chloride." You should: (1) state electrons are transferred, (2) draw a dot-and-cross diagram with correct charges, (3) state ions are held by strong electrostatic attraction, and (4) mention the giant ionic lattice.

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Practice Questions

1. Describe what happens to the electrons when sodium reacts with chlorine to form sodium chloride.
2. Draw a dot-and-cross diagram for magnesium oxide (MgO).
3. Explain why the formula of calcium chloride is CaCl₂ and not CaCl.
4. What is meant by the term "ionic bond"?
5. Describe the structure of an ionic compound.
6. Draw a dot-and-cross diagram for lithium fluoride (LiF) and state the charges on each ion.

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Worked Example: Predicting Formulae from Charges

A common exam skill is to predict the formula of an ionic compound from the charges on the ions. Use the swap-and-drop rule: the numerical charge on the cation becomes the subscript on the anion, and the numerical charge on the anion becomes the subscript on the cation (ignoring signs).

Cation	Anion	Charges	Swap subscripts	Simplified formula
Al³⁺	O²⁻	3, 2	Al₂O₃	Al₂O₃
Ca²⁺	N³⁻	2, 3	Ca₃N₂	Ca₃N₂
K⁺	S²⁻	1, 2	K₂S₁	K₂S
Mg²⁺	O²⁻	2, 2	Mg₂O₂	MgO (divide by 2)

Common Mistake: Students often forget to simplify. Mg²⁺ with O²⁻ gives Mg₂O₂ using the swap rule, but since both subscripts share a common factor of 2, the formula simplifies to MgO. Always check whether subscripts can be divided by a common factor.

Polyatomic Ions

Some ions contain more than one atom joined by covalent bonds, yet carry an overall charge. These are called polyatomic ions. You are not required to memorise many of them at GCSE, but you should recognise a few common ones:

Polyatomic ion	Formula	Charge
Hydroxide	OH⁻	-1
Nitrate	NO₃⁻	-1
Carbonate	CO₃²⁻	-2
Sulfate	SO₄²⁻	-2
Ammonium	NH₄⁺	+1

When writing formulae containing polyatomic ions, use brackets if there is more than one of the polyatomic ion, e.g. Ca(OH)₂, Al₂(SO₄)₃.

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Balanced Equations for Ion Formation

You may be asked to write half-equations showing electron transfer. Electrons are written as e⁻:

• Sodium losing an electron: Na → Na⁺ + e⁻
• Chlorine atom gaining an electron: Cl + e⁻ → Cl⁻ (or ½ Cl₂ + e⁻ → Cl⁻)
• Magnesium losing two electrons: Mg → Mg²⁺ + 2e⁻
• Oxygen gaining two electrons: O + 2e⁻ → O²⁻

Exam Tip: In half-equations, electrons appear on the right-hand side for oxidation (loss of electrons, done by the metal) and on the left-hand side for reduction (gain of electrons, done by the non-metal). The mnemonic OIL RIG — "Oxidation Is Loss, Reduction Is Gain" — is your friend.

Overall Equation

The overall reaction between sodium and chlorine is:

2Na(s) + Cl₂(g) → 2NaCl(s)

Note that chlorine exists as a diatomic molecule (Cl₂), so two sodium atoms are needed to react with one Cl₂ molecule to produce two NaCl formula units.

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Visualising Electron Transfer

flowchart LR
    A["Na atom<br/>2, 8, 1"] -->|"loses 1e⁻"| B["Na⁺ ion<br/>2, 8"]
    C["Cl atom<br/>2, 8, 7"] -->|"gains 1e⁻"| D["Cl⁻ ion<br/>2, 8, 8"]
    B --> E["Giant ionic lattice<br/>NaCl"]
    D --> E

The arrow from Na to Cl represents the transfer of one electron. Notice that electrons are neither created nor destroyed — they simply move from the metal atom to the non-metal atom.

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Common Mistakes Callout

Watch Out! Students routinely lose marks for the following errors:

• Describing ionic bonding as "sharing electrons" — this is covalent bonding, not ionic.
• Forgetting the charges on ions in a dot-and-cross diagram (always include + or - outside the bracket).
• Writing "ionic bonds are strong" without specifying what makes them strong (it is the electrostatic attraction between oppositely charged ions).
• Drawing individual molecules of NaCl — sodium chloride does not exist as molecules; it is a giant ionic lattice.
• Forgetting to state that attraction acts in all directions throughout the lattice.

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Extended Worked Example: Aluminium Oxide

Question: Aluminium reacts with oxygen to form aluminium oxide. Describe the bonding in aluminium oxide and write the balanced equation for its formation.

Answer:

• Aluminium has electronic structure 2, 8, 3 — it has 3 electrons in its outer shell.
• Oxygen has electronic structure 2, 6 — it needs 2 electrons to fill its outer shell.
• Each aluminium atom loses 3 electrons to form Al³⁺ (electronic structure 2, 8).
• Each oxygen atom gains 2 electrons to form O²⁻ (electronic structure 2, 8).
• To balance charges: 2 × Al³⁺ gives +6, and 3 × O²⁻ gives -6. Therefore the formula is Al₂O₃.
• Balanced equation: 4Al(s) + 3O₂(g) → 2Al₂O₃(s).
• Al³⁺ and O²⁻ ions are held in a giant ionic lattice by strong electrostatic attraction acting in all directions. This explains why aluminium oxide has an extremely high melting point (around 2072 °C) — a very large amount of energy is required to overcome the attractions between 3+ and 2- ions.

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Answering at different grade levels

The same question about ionic bonding in sodium chloride can be answered at very different levels of precision. Examiners reward specific, technical vocabulary.

Grade 3-4 response: "Sodium gives an electron to chlorine. The sodium becomes positive and the chlorine becomes negative. They stick together because opposite charges attract."

Why this scores low: It is vague. It does not use the term ionic bond, does not mention the giant ionic lattice, does not explain why the atoms transfer electrons, and does not refer to electrostatic attraction or full outer shells.

Grade 5-6 response: "Sodium loses 1 electron from its outer shell to form a Na⁺ ion with a full outer shell. Chlorine gains that electron to form a Cl⁻ ion, also with a full outer shell. The positive Na⁺ and negative Cl⁻ ions are attracted to each other and form an ionic bond. The ions form a regular pattern called a giant ionic lattice."

Why this scores mid-band: It mentions full outer shells, names the ions correctly and uses the terms "ionic bond" and "giant ionic lattice". However, it does not explicitly say electrostatic attraction, does not say the attraction acts in all directions, and does not mention that the noble-gas electron configuration is the reason for electron transfer.

Grade 7-9 response: "Atoms react to achieve a full outer shell, the same electron configuration as the nearest noble gas. Sodium (2, 8, 1) loses its single outer-shell electron to form a Na⁺ cation with the configuration 2, 8 (like neon). Chlorine (2, 8, 7) gains that electron to form a Cl⁻ anion with the configuration 2, 8, 8 (like argon). The resulting oppositely charged ions are held together by strong electrostatic attraction — this is an ionic bond. In the solid, billions of Na⁺ and Cl⁻ ions are arranged in a regular, repeating 3D pattern called a giant ionic lattice, with each Na⁺ surrounded by 6 Cl⁻ and vice versa. The electrostatic attraction acts in all directions between oppositely charged ions, which is why very high temperatures are needed to separate them."

Why this scores top-band: Uses precise terms ("cation", "anion", "electrostatic attraction", "giant ionic lattice"), explains the underlying reason for electron transfer (noble-gas configuration), describes the 3D nature and coordination number of the lattice, and explicitly states that the attraction acts in all directions.

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Edexcel alignment: This content is aligned with Edexcel GCSE Chemistry (1CH0) specification Topic 1 Key concepts in chemistry — specifically 1.5 Ionic bonding and 1.7 Types of substance. Assessed on Paper 1.