Chemical bonds — AQA GCSE Combined Science
Test yourself on Chemical bonds with AQA GCSE practice questions.
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Chemical bonds explained
This statement defines the three strong chemical bonds.
Read the full explanation
Ionic bonding occurs between oppositely charged ions, usually a metal and a non-metal, where electrons are transferred to form positive and negative ions. For example, sodium loses one electron to form Na⁺ and chlorine gains one electron to form Cl⁻, giving NaCl. Covalent bonding occurs between non-metal atoms that share pairs of electrons, as in H₂ where each hydrogen shares one pair, or in O₂ where two pairs are shared. Metallic bonding occurs in metals, where positive metal ions are held in a lattice of delocalised electrons. The shared delocalised electrons explain why metals conduct electricity and heat. When describing a bond, state the particles involved and how electrons are involved.
Ionic bonding occurs in compounds formed from metals combined with non-metals.
Ionic bonding happens when a metal reacts with a non-metal. Metal atoms lose electrons to form positive ions, while non-metal atoms gain those electrons to form negative ions. The oppositely charged ions are held together by strong electrostatic attraction in a giant ionic lattice. For example, sodium (a Group 1 metal) transfers one electron to chlorine (a Group 7 non-metal), producing Na⁺ and Cl⁻ ions in sodium chloride. The compound is neutral overall because the total positive charge balances the total negative charge. This explains why ionic compounds have high melting and boiling points, conduct electricity when molten or dissolved, and often dissolve in water. The key idea is that ionic bonding is identified from the metal-plus-non-metal combination, not from the name alone.
Covalent bonding occurs in most non-metallic elements and in compounds of non-metals.
Covalent bonding occurs when non-metal atoms share pairs of electrons. This happens in most non-metallic elements, such as hydrogen (H₂), oxygen (O₂), nitrogen (N₂) and chlorine (Cl₂), and in compounds of non-metals, such as water (H₂O), carbon dioxide (CO₂) and methane (CH₄). Each shared pair of electrons is attracted to the nuclei of both atoms, holding them together. Atoms share electrons so that each atom achieves a full outer shell, often eight electrons, except hydrogen which needs two. The number of covalent bonds an atom forms depends on how many electrons it needs to complete its outer shell. For example, oxygen needs two more electrons, so it forms two covalent bonds, as in O₂ or H₂O. Covalent substances can be small molecules or giant covalent structures, and their properties depend on their structure and the strength of the forces between particles.
Metallic bonding occurs in metallic elements and alloys.
Metallic bonding explains why metals and alloys hold together and conduct. In a pure metal such as magnesium, each atom contributes its outer electrons to a delocalised sea that surrounds a regular lattice of positive ions. The attraction between the positive ions and the shared negative electrons is the metallic bond, acting in all directions rather than between specific pairs. Alloys such as steel or bronze contain atoms of different sizes, which disrupt the regular layers so they cannot slide as easily, making the alloy harder. This model accounts for the high melting points, malleability and electrical conductivity of metallic substances, and it applies to metallic elements and to alloys alike.
Students should be able to explain chemical bonding in terms of electrostatic forces and the transfer or sharing of electrons.
Chemical bonding can be explained by electrostatic forces and by the transfer or sharing of electrons. In ionic bonding, a metal atom transfers one or more electrons to a non-metal atom, forming positive and negative ions that attract each other in a giant lattice, as in sodium chloride. In covalent bonding, non-metal atoms share pairs of electrons so that both nuclei are attracted to the shared pair, as in a hydrogen molecule or water. In metallic bonding, positive ions are attracted to a sea of delocalised electrons. In every case the bond is the electrostatic attraction between opposite charges, and the type of bonding depends on whether electrons are transferred or shared.
Your focus
- Identify a bond as ionic, covalent or metallic from a given formula or particle diagram.
- Describe the role of electrons in each of the three strong chemical bonds.
- Represent ionic and covalent bonding using dot-and-cross diagrams or ion symbols.
Show all 15 objectives
- Identify from a formula or reaction whether a compound is likely to be ionic by checking for a metal combined with a non-metal.
- Draw or interpret a dot-and-cross diagram showing electron transfer and the resulting ions with correct charges.
- Explain at least one property of an ionic compound in terms of its giant lattice and electrostatic attraction.
- Identify covalent bonding in a given substance by recognising that all the elements present are non-metals.
- Draw or complete a dot-and-cross diagram for a simple covalent molecule such as H₂O or CH₄.
- Explain how the number of shared pairs relates to the outer-shell electron configuration of the atoms involved.
- Describe the arrangement of positive ions and delocalised electrons in a metal.
- Explain how metallic bonding holds a metallic element or alloy together.
- Relate the structure of an alloy to its greater hardness compared with a pure metal.
- Describe how electrostatic forces produce chemical bonds.
- Distinguish between electron transfer in ionic bonding and electron sharing in covalent bonding.
- Explain the bonding in a named ionic compound, covalent molecule or metal using the correct particles and charges.
Chemical bonds exam tips
Marking Points
- Name the three strong chemical bonds as ionic, covalent and metallic.
- Describe ionic bonding as electrostatic attraction between oppositely charged ions, formed by electron transfer.
- Describe covalent bonding as sharing pairs of electrons between atoms, often non-metal atoms.
- Describe metallic bonding as positive metal ions in a lattice with delocalised electrons.
- Use examples such as NaCl, H₂O or Mg to illustrate each bonding type.
- Identify that ionic bonding occurs between a metal and a non-metal, for example sodium and chlorine in sodium chloride.
- Describe electron transfer: metal atoms lose electrons to form positive ions and non-metal atoms gain electrons to form negative ions.
- State that the ions formed have full outer shells and are held in a giant ionic lattice by strong electrostatic attraction.
- Explain that the overall compound is neutral because the total positive charge equals the total negative charge.
- Link the structure to a property, such as high melting point or conduction when molten or aqueous, to show understanding of the bonding model.
- Identify covalent bonding as sharing pairs of electrons between non-metal atoms.
- Give examples of non-metallic elements that are covalently bonded, such as H₂, O₂, N₂ or Cl₂.
- Give examples of non-metal compounds that are covalently bonded, such as H₂O, CO₂ or CH₄.
- Draw or interpret dot-and-cross diagrams showing shared electron pairs and full outer shells.
- Relate the number of covalent bonds formed by an atom to the number of electrons needed to complete its outer shell.
- Metallic bonding is the electrostatic attraction between positive metal ions and delocalised electrons.
- The structure is a giant lattice of positive ions in a sea of delocalised electrons.
- The bonding acts in all directions and is not directional between particular pairs of atoms.
- Outer-shell electrons are delocalised, so they are free to move through the lattice.
- Alloys contain atoms of different sizes, which distort the regular layers and make sliding harder.
- The model explains properties such as conductivity, malleability and high melting points.
- Chemical bonds arise from electrostatic attraction between oppositely charged particles.
- Ionic bonding involves transfer of electrons from a metal atom to a non-metal atom.
- Ionic bonding produces positive and negative ions held in a giant lattice.
- Covalent bonding involves sharing pairs of electrons between non-metal atoms.
- In covalent bonding, the shared pair is attracted to both nuclei.
- Metallic bonding involves attraction between positive ions and delocalised electrons.
- The type of bonding present depends on whether electrons are transferred or shared.
Examiner Tips
- 💡For each bond type, state the particles involved and how electrons are involved.
- 💡Use a dot-and-cross diagram to support a written description of covalent bonding.
- 💡When comparing bonds, use the terms 'transfer', 'share' and 'delocalised' accurately.
- 💡Underline the metal and non-metal in the question before deciding whether ionic bonding is present.
- 💡When drawing dot-and-cross diagrams for ionic compounds, show the transfer of electrons and include the charge on each ion.
- 💡Use the phrase 'strong electrostatic attraction between oppositely charged ions' when explaining properties of ionic compounds.
- 💡Check that every element in the substance is a non-metal before deciding that covalent bonding is present.
- 💡In dot-and-cross diagrams, use different symbols for electrons from different atoms and show the shared pair in the overlap region.
- 💡When explaining properties of simple molecular substances, refer to weak intermolecular forces rather than breaking covalent bonds.
- 💡Use the phrase 'electrostatic attraction between positive ions and delocalised electrons' when explaining metallic bonding.
- 💡Link each property you describe back to the model, for example free electrons explain conductivity.
- 💡When comparing an alloy with a pure metal, refer to the sizes of the atoms and the effect on sliding layers.
- 💡Name the particles involved and state whether electrons are transferred or shared before explaining the attraction.
- 💡Use a dot-and-cross diagram to support your explanation of electron transfer or sharing.
- 💡For each example, identify the electrostatic attraction that holds the particles together.
Common Mistakes
- Saying covalent bonds share single electrons rather than pairs of electrons. Correction: covalent bonds involve shared pairs of electrons.
- Describing ionic bonding as sharing electrons. Correction: ionic bonding involves transfer of electrons to form oppositely charged ions.
- Stating that metallic bonding involves shared pairs of electrons between specific atoms. Correction: metallic bonding involves delocalised electrons shared across the whole metal lattice.
- Writing that ionic bonds form between two non-metals; correct this by checking that one element is a metal and the other is a non-metal.
- Saying that atoms share electrons in ionic bonding; correct this by describing transfer of electrons to form ions.
- Drawing ions without charge labels or with incorrect charges; correct this by showing Na⁺ and Cl⁻, and balancing charges such as Mg²⁺ with two Cl⁻ ions.
- Stating that covalent bonds involve transfer of electrons; correct this by saying electrons are shared between non-metal atoms.
- Drawing dot-and-cross diagrams with unpaired electrons in the outer shell of a bonded atom; correct this by ensuring each atom has a full outer shell where possible.
- Confusing covalent compounds with ionic compounds; correct this by checking that all elements involved are non-metals.
- Saying metallic bonding is between positive and negative ions: correct this by stating it is between positive ions and delocalised electrons.
- Describing delocalised electrons as belonging to one particular atom: correct this by saying they are shared across the whole lattice.
- Claiming alloys are weaker because different-sized atoms break the bonds: correct this by explaining that different-sized atoms disrupt the regular layers so they cannot slide, making the alloy harder.
- Saying covalent bonds involve transfer of electrons: correct this by stating that non-metal atoms share pairs of electrons.
- Describing ionic bonding as sharing electrons: correct this by explaining that electrons are transferred to form ions.
- Stating that bonds form because atoms want full shells without mentioning electrostatic attraction: correct this by linking bond formation to attraction between opposite charges.