Revision notes for Edexcel GCSE Chemistry Ionic bonding. Open the guide for explanations and worked examples. Written against the Edexcel GCSE Chemistry (1CH0) specification, so the content matches what's examinable rather than general Chemistry background.
Ionic bonding
Welcome to your study notes on ionic bonding! This is a core topic in your GCSE Chemistry journey. By understanding how and why atoms exchange electrons, you will unlock the secrets behind how many everyday materials behave—from the table salt you put on your food to the giant crystal minerals found deep in the Earth.
What you'll learn:
- How atoms gain or lose electrons to become stable cations and anions.
- How to represent ionic bonding using dot-and-cross diagrams.
- How to balance charges to write correct chemical formulas for ionic compounds.
- How ions organize themselves into a giant ionic lattice held together by strong electrostatic forces.
1. What is an Ion?
Before looking at how bonds form, we must first understand what an ion is.
In an atom, the number of positive protons in the nucleus is always equal to the number of negative electrons orbiting it. This means atoms have no overall electrical charge; they are neutral.
However, atoms are most stable when they have a full outer shell of electrons (usually 8 electrons, or 2 for the innermost shell). To achieve this stable state, atoms will often lose or gain electrons. When they do, the balance between protons and electrons is lost, and a charged particle is formed.
Definition
Ion
An ion is an atom or group of atoms with a positive or negative charge, caused by the loss or gain of one or more electrons.
- Cations are positively charged ions. They form when an atom (usually a metal) loses negative electrons. Because there are now more protons than electrons, the overall charge is positive.
- Anions are negatively charged ions. They form when an atom (usually a non-metal) gains negative electrons. Because there are now more electrons than protons, the overall charge is negative.
Analogy
Remembering Cations vs Anions
If you struggle to remember which is which, think of cats:
- Cations are pawsitive (positive).
- An anion sounds like "a negative ion" (A-Negative-Ion).
Calculating Protons, Neutrons, and Electrons in Ions
You need to be able to determine the subatomic particles in an ion using its mass number, atomic number, and charge. Let's look at how to do this:
- Proton number is always equal to the atomic number. This never changes when an atom becomes an ion.
- Neutron number is calculated by subtracting the atomic number from the mass number:
- Electron number depends on the charge:
- For a neutral atom, electrons = protons.
- For a positive ion (cation), subtract the charge value from the atomic number.
- For a negative ion (anion), add the charge value to the atomic number.
Example
Calculating particles in an oxide ion
Calculate the number of protons, neutrons, and electrons in an oxide ion, 816O2−^{16}_{8}\text{O}^{2-}816O2−.
- Find the proton number directly from the atomic number (the bottom number):
- Calculate the neutron number by subtracting the atomic number from the mass number (the top number):
- Determine the electron number by starting with the proton number and adjusting for the 2−2-2− charge (which means the ion has gained 2 negative electrons):
2. How Ionic Bonds Form
Ionic bonds form when a metal atom reacts with a non-metal atom.
The metal atom transfers its outer-shell electrons to the non-metal atom. This allows both atoms to achieve a stable, full outer electron shell.
Key Idea
The Ionic Transfer
- Metals lose electrons to form positive ions.
- Non-metals gain those electrons to form negative ions.
- The oppositely charged ions are then strongly attracted to each other by electrostatic forces. This attraction is the ionic bond.
Groups in the Periodic Table and Ion Charges
The group number of an element in the periodic table tells you how many electrons are in its outer shell. This makes it easy to predict what ion it will form:
- Group 1 metals have 1 outer electron. They lose 1 electron to form 1+1+1+ ions (e.g., Na+\text{Na}^+Na+).
- Group 2 metals have 2 outer electrons. They lose 2 electrons to form 2+2+2+ ions (e.g., Mg2+\text{Mg}^{2+}Mg2+).
- Group 6 non-metals have 6 outer electrons. They gain 2 electrons to form 2−2-2− ions (e.g., O2−\text{O}^{2-}O2−).
- Group 7 non-metals have 7 outer electrons. They gain 1 electron to form 1−1-1− ions (e.g., Cl−\text{Cl}^-Cl−).
Dot-and-Cross Diagrams
We use dot-and-cross diagrams to show how electrons are transferred during ionic bonding. One atom's electrons are drawn as dots (∙\bullet∙), and the other's are drawn as crosses (×\times×). This helps us keep track of where the electrons came from.
When drawing these in an exam:
- Show the electron transfer with an arrow from the metal's outer electron to the non-metal's outer shell.
- Draw the resulting ions inside square brackets with their charges written outside the brackets on the top right (e.g., [Na]+[\text{Na}]^+[Na]+ and [Cl]−[\text{Cl}]^-[Cl]−).
- The non-metal ion should show its original electrons (crosses) plus the electron(s) it gained from the metal (dots).
Common Mistake
Missing the square brackets or charges
A very common mistake is drawing the ions at the end of the reaction but forgetting to draw the square brackets or writing the charges. Examiners will look for both! Always put ions in square brackets with the charge clearly written outside at the top right.
3. Naming Ionic Compounds
The names of ionic compounds tell you which ions are present. You need to understand the difference between two common endings: -ide and -ate.
Compounds ending in "-ide"
If an ionic compound contains only two elements (a metal and a non-metal), the name of the non-metal changes to end in -ide.
- Sodium + Chlorine →\rightarrow→ Sodium chloride
- Magnesium + Oxygen →\rightarrow→ Magnesium oxide
Compounds ending in "-ate"
If the compound contains a metal, a non-metal, and oxygen, the name of the non-metal part changes to end in -ate. These compounds contain polyatomic ions (groups of atoms with an overall charge that contain oxygen).
- Copper + Sulfur + Oxygen →\rightarrow→ Copper sulfate (CuSO4\text{CuSO}_4CuSO4)
- Sodium + Nitrogen + Oxygen →\rightarrow→ Sodium nitrate (NaNO3\text{NaNO}_3NaNO3)
4. Deducing the Formulae of Ionic Compounds
Because compounds are electrically neutral, the total positive charge of the cations must exactly equal the total negative charge of the anions.
To write the formula of an ionic compound, you must balance the charges. You need to know the formulas and charges of these key ions:
| Ion Name | Chemical Formula | Ion Name | Chemical Formula |
|---|---|---|---|
| Oxide | O2−\text{O}^{2-}O2− | Fluoride / Chloride | F−/Cl−\text{F}^- / \text{Cl}^-F−/Cl− |
| Hydroxide | OH−\text{OH}^-OH− | Nitrate | NO3−\text{NO}_3^-NO3− |
| Sulfate | SO42−\text{SO}_4^{2-}SO42− | Carbonate | CO32−\text{CO}_3^{2-}CO32− |
Tip
Shorthand for polyatomic ions
When you have more than one of a polyatomic ion (like NO3−\text{NO}_3^-NO3− or OH−\text{OH}^-OH−) in a chemical formula, you must put the whole ion in brackets before adding a subscript number. For example, two nitrate ions are written as (NO3)2(\text{NO}_3)_2(NO3)2.
Example
Deducing the formula of Calcium Nitrate
Deduce the chemical formula for calcium nitrate.
- Write down the formulas and charges of the individual ions:
- Identify the balance of charges. The calcium ion has a 2+2+2+ charge, while each nitrate ion has only a 1−1-1− charge. To make the compound neutral, you need two nitrate ions to balance one calcium ion:
- Write the formula using brackets for the nitrate ion because you need more than one of it:
5. Giant Ionic Lattices
Ionic compounds do not exist as single, isolated pairs of molecules. Instead, millions of positive and negative ions cluster together in a highly ordered, three-dimensional structure called a giant ionic lattice.
Key Features of a Giant Ionic Lattice:
- Regular arrangement: The ions are packed together in a highly ordered, repeating grid alternating between positive and negative ions.
- Electrostatic forces: Strong forces of electrostatic attraction hold the oppositely-charged ions together.
- Omnidirectional bonding: These electrostatic attractions (the ionic bonds) act in all directions throughout the entire giant structure.
Because these ionic bonds are extremely strong, it takes a massive amount of thermal energy to break them apart. This explains why ionic compounds like sodium chloride have extremely high melting and boiling points!
Exam technique
In the exam
- Show every shell or just the outer? Read exam questions carefully. Some questions ask you to draw only the outer shell electrons in a dot-and-cross diagram, whereas others ask for all electron shells.
- Use different symbols. Always use dots (∙\bullet∙) for one atom's electrons and crosses (×\times×) for the other's to clearly show the transfer.
- Double check polyatomic brackets. If you need more than one polyatomic ion, always use brackets. Write Mg(OH)2\text{Mg(OH)}_2Mg(OH)2, not MgOH2\text{MgOH}_2MgOH2 (which would incorrectly mean 1 magnesium, 1 oxygen, and 2 hydrogens).
Self review
Check yourself
- State the formula and charge of a sulfate ion.
- Explain, in terms of subatomic particles, why a magnesium atom forms a Mg2+\text{Mg}^{2+}Mg2+ ion rather than a Mg2−\text{Mg}^{2-}Mg2− ion.
- Describe the structure of a giant ionic lattice, including the forces that hold it together.
Recap questions
Test yourself with 5 quick questions on this guide. Answer them all correctly to complete it.
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