Ionic Bonding

This lesson covers ionic bonding — the transfer of electrons between metals and non-metals to form charged particles called ions — as required by the Edexcel GCSE Combined Science specification (1SC0). You need to understand how ionic bonds form, draw dot-and-cross diagrams, and explain why ions have charges.

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What Is Ionic Bonding?

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

The process involves the transfer of electrons:

1. A metal atom loses one or more electrons from its outer shell.
2. A non-metal atom gains those electrons into its outer shell.
3. Both atoms now have a full outer shell — a stable electron configuration like a noble gas.
4. The metal becomes a positive ion (cation) and the non-metal becomes a negative ion (anion).
5. The opposite charges attract each other strongly — this is the ionic bond.

Exam Tip: Always state that ionic bonding involves the transfer of electrons, not the sharing of electrons. Sharing is covalent bonding.

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Why Do Atoms Form Ions?

Atoms are most stable when they have a full outer electron shell. This is the same electron configuration as a noble gas (Group 0). Atoms achieve this by:

• Metals — losing electrons to empty their outer shell, revealing the full shell beneath.
• Non-metals — gaining electrons to fill their outer shell.

Atom	Group	Electrons Lost/Gained	Ion Formed	Noble Gas Configuration
Sodium (Na)	1	Loses 1 electron	Na⁺	Neon (2, 8)
Magnesium (Mg)	2	Loses 2 electrons	Mg²⁺	Neon (2, 8)
Aluminium (Al)	3	Loses 3 electrons	Al³⁺	Neon (2, 8)
Chlorine (Cl)	7	Gains 1 electron	Cl⁻	Argon (2, 8, 8)
Oxygen (O)	6	Gains 2 electrons	O²⁻	Neon (2, 8)

Exam Tip: The charge on an ion can be predicted from the group number. Group 1 metals form +1 ions, Group 2 form +2 ions, Group 6 non-metals form −2 ions, and Group 7 non-metals form −1 ions.

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Dot-and-Cross Diagrams for Ionic Bonding

A dot-and-cross diagram shows how electrons are transferred during ionic bonding. Dots represent electrons from one atom and crosses represent electrons from the other.

Example 1: Sodium Chloride (NaCl)

graph LR
    A["Na atom<br/>2, 8, 1"] -->|"Loses 1 electron"| B["Na⁺ ion<br/>2, 8"]
    C["Cl atom<br/>2, 8, 7"] -->|"Gains 1 electron"| D["Cl⁻ ion<br/>2, 8, 8"]
    A -->|"1 electron transferred"| C

    style A fill:#2980b9,color:#fff
    style B fill:#27ae60,color:#fff
    style C fill:#e67e22,color:#fff
    style D fill:#27ae60,color:#fff

• Sodium (2, 8, 1) loses its 1 outer electron → Na⁺ (2, 8).
• Chlorine (2, 8, 7) gains that electron → Cl⁻ (2, 8, 8).
• Both ions now have a full outer shell.
• The formula is NaCl because one sodium ion bonds with one chloride ion.

Example 2: Magnesium Oxide (MgO)

• Magnesium (2, 8, 2) loses 2 outer electrons → Mg²⁺ (2, 8).
• Oxygen (2, 6) gains those 2 electrons → O²⁻ (2, 8).
• The formula is MgO — the charges balance: 2+ and 2−.

Example 3: Magnesium Chloride (MgCl₂)

• Magnesium needs to lose 2 electrons but chlorine can only gain 1.
• Therefore, magnesium transfers one electron to each of two chlorine atoms.
• Mg²⁺ + 2Cl⁻ → MgCl₂.

Example 4: Sodium Oxide (Na₂O)

• Oxygen needs to gain 2 electrons but sodium can only lose 1.
• Two sodium atoms each transfer one electron to one oxygen atom.
• 2Na⁺ + O²⁻ → Na₂O.

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Drawing Dot-and-Cross Diagrams — Rules

When drawing dot-and-cross diagrams for ionic compounds in the exam:

1. Show only the outer shell electrons (unless the question asks for all shells).
2. Use dots for electrons from one atom and crosses for the other.
3. Draw square brackets around each ion.
4. Write the charge of each ion outside the brackets (e.g. [Na]⁺ or [Cl]⁻).
5. Show the transferred electron changing from a dot to a cross (or vice versa) so the examiner can see it came from the other atom.

Exam Tip: Always include square brackets and charges on your dot-and-cross diagrams. Missing brackets or charges will lose marks even if the electron arrangement is correct.

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Working Out Formulae of Ionic Compounds

To work out the formula, make sure the total positive charge equals the total negative charge so the compound is neutral overall.

Positive Ion	Negative Ion	Balancing	Formula
Na⁺	Cl⁻	1 × (+1) + 1 × (−1) = 0	NaCl
Mg²⁺	Cl⁻	1 × (+2) + 2 × (−1) = 0	MgCl₂
Na⁺	O²⁻	2 × (+1) + 1 × (−2) = 0	Na₂O
Ca²⁺	F⁻	1 × (+2) + 2 × (−1) = 0	CaF₂
Al³⁺	O²⁻	2 × (+3) + 3 × (−2) = 0	Al₂O₃

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Summary

• Ionic bonding is the strong electrostatic attraction between oppositely charged ions, formed by the transfer of electrons from a metal to a non-metal.
• 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).
• Dot-and-cross diagrams show the electron transfer — always include square brackets and charges.
• The formula of an ionic compound is found by balancing the charges so the compound is electrically neutral.

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Worked Examples — Building Formulae from Charges

Exam questions often give you two elements or two ions and ask you to deduce the formula of the resulting ionic compound. The trick is always the same: the total positive charge must cancel the total negative charge so the compound is electrically neutral. Learn the shortcut: swap and drop the charges as subscripts, then cancel any common factor.

Example 1 — Calcium chloride

• Calcium is in Group 2, so it forms Ca²⁺.
• Chlorine is in Group 7, so it forms Cl⁻.
• Swap the charges to get subscripts: Ca₁Cl₂ → CaCl₂.
• Check: 1 × (+2) + 2 × (−1) = 0. Correct.

Example 2 — Aluminium oxide

• Aluminium is in Group 3, so it forms Al³⁺.
• Oxygen is in Group 6, so it forms O²⁻.
• Swap the charges: Al₂O₃. Check: 2 × (+3) + 3 × (−2) = 0. Correct.

Example 3 — Magnesium nitride

• Magnesium forms Mg²⁺; nitrogen forms N³⁻ (Group 5, gains 3 electrons).
• Swap the charges: Mg₃N₂. Check: 3 × (+2) + 2 × (−3) = 0. Correct.

Example 4 — Potassium sulfide

• Potassium forms K⁺; sulfur forms S²⁻.
• Swap the charges: K₂S. Check: 2 × (+1) + 1 × (−2) = 0. Correct.

Exam Tip: After swapping, always cancel any common factor. For example, calcium oxide is CaO, not Ca₂O₂, because both subscripts share a common factor of 2. Leave the formula in its simplest whole-number ratio.

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Common Mistakes and How to Avoid Them

Mistake	What Happens	How to Fix
Writing "Na shares electrons with Cl"	Loses the mark for bonding type	State transfer of electrons for ionic bonding
Forgetting square brackets on ions	Examiner cannot see that the species is a charged ion	Always draw brackets and the charge, e.g. [Na]⁺
Drawing the transferred electron as a dot (same as the metal)	Examiner cannot see the electron came from the non-metal	Show the transferred electron as the opposite symbol inside the non-metal ion
Saying "atoms become stable" without specifying	Vague, no marks	Say atoms achieve a full outer shell, the noble gas configuration
Confusing +2 ions with Group 2 elements	Mix-ups over charge	Group number = number of outer electrons = charge of ion for Groups 1–2
Writing formulae with fractions	Wrong conventions	Always simplify to whole-number ratios

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Comparing Ionic Bonding with Covalent and Metallic

Feature	Ionic	Covalent	Metallic
Electrons	Transferred	Shared	Delocalised
Between	Metal + non-metal	Non-metal + non-metal	Metals only
Bond strength	Strong	Strong (within molecule)	Strong
Key phrase	Electrostatic attraction between ions	Shared pair of electrons	Sea of delocalised electrons

The common thread is electrostatic attraction — opposite charges attracting — but the particles involved differ. In ionic bonding the attracting particles are whole ions; in metallic bonding they are positive ions and delocalised electrons; in covalent bonding it is the shared electron pair attracted to both nuclei.

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Mini Flowchart — Deducing an Ionic Formula

graph TD
    A["Identify metal (cation)<br/>and non-metal (anion)"] --> B["Find group number<br/>of each"]
    B --> C["Work out charges:<br/>Groups 1,2,3 = +1,+2,+3<br/>Groups 5,6,7 = −3,−2,−1"]
    C --> D["Swap charges to<br/>become subscripts"]
    D --> E["Cancel any<br/>common factor"]
    E --> F["Check: total<br/>positive = total<br/>negative"]

    style A fill:#2c3e50,color:#fff
    style B fill:#2980b9,color:#fff
    style C fill:#27ae60,color:#fff
    style D fill:#e67e22,color:#fff
    style E fill:#c0392b,color:#fff
    style F fill:#1a1a2e,color:#fff

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Answering at different grade levels

Grade descriptors vary with how precisely you use the key terms: ionic bond, electrostatic attraction, cation, anion and giant ionic lattice.

• Grade 3–4 response: "Sodium gives an electron to chlorine. Sodium has +1 charge and chlorine has −1 charge. They stick together."
• Grade 5–6 response: "Sodium loses one outer electron to chlorine. This forms Na⁺ and Cl⁻. The ionic bond is the attraction between these oppositely charged ions. Both ions now have a full outer shell."
• Grade 7–9 response: "Sodium (2,8,1) transfers its single outer electron to chlorine (2,8,7). Sodium becomes the cation Na⁺ with the neon configuration (2,8), while chlorine becomes the anion Cl⁻ with the argon configuration (2,8,8). The ionic bond is the strong electrostatic attraction between oppositely charged ions, which extends throughout a giant ionic lattice rather than existing as a pair of particles."

The Grade 7–9 answer embeds electron configurations, names each ion type (cation/anion), and explicitly links the bond to the lattice — this is what earns the highest marks.

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Further Practice — Predicting Ion Charges from the Periodic Table

You should be able to predict the charge of any main-group ion without being given it. The rule is simple:

Group	Electrons in outer shell	Electrons lost or gained	Ion charge
1	1	Lose 1	+1
2	2	Lose 2	+2
3	3	Lose 3	+3
4	4	Rare (see note)	N/A
5	5	Gain 3	−3
6	6	Gain 2	−2
7	7	Gain 1	−1
0	8 (or 2 for He)	None	No ion

Note: Group 4 elements (like carbon) usually form covalent bonds rather than ions because transferring 4 electrons would leave a strongly charged ion that is energetically unfavourable. Silicon tends to bond covalently too.

Why do Group 0 elements not form ions?

Noble gases already have a full outer shell — they are at their most stable form. They have no energetic reason to lose or gain electrons, so they remain as individual, unbonded atoms at standard conditions.

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Deeper Look — Electrostatic Attraction in Context

The word electrostatic simply means "between stationary charges." When we say an ionic bond is electrostatic attraction, we mean:

• Na⁺ carries a +1 charge.
• Cl⁻ carries a −1 charge.
• Opposite charges attract with significant force at short range.
• Bigger charges and smaller distances both make the attraction stronger.

You will not be tested on Coulomb's formula at GCSE, but you should know the words "electrostatic attraction between oppositely charged ions" — this is the precise phrasing examiners want. The implications run through many exam questions: Mg²⁺O²⁻ has more attraction than Na⁺Cl⁻ because the charges are greater, which is why MgO melts at 2852 °C while NaCl melts at 801 °C.

Edexcel alignment: This content is aligned with Edexcel GCSE Combined Science (1SC0) Chemistry Topic 2 States of matter and mixtures / Topic 3 Bonding — specifically CC5 Ionic bonding, CC6 Covalent bonding, CC7 Types of substance, CC8 Metallic bonding. Assessed on Chemistry Paper 1.