Lesson 3.1.3.3

3.1.3.3 Metallic bonding Quiz: AQA Chemistry, Unit 1

20 questions

Written by Revision Ninja · Last reviewed 11 October 2026

In partnership with Revision Ninja

Lesson 3.1.3.3, Metallic bonding: 20 multiple choice questions for the AQA Chemistry (7405), Unit 1: Physical chemistry, written with Revision Ninja.

Host it live on the board and students join with a game code on their own devices, or revise alone with Free Play. The answers are revealed in the game.

Host this setFree Play

The 20 questions

  1. What is metallic bonding?

    • Sharing of electron pairs between two atoms, which are held together by a covalent bond
    • Electrostatic attraction between positive and negative ions that are held in a fixed lattice
    • Attraction between delocalised electrons and positive ions arranged in a lattice
    • Attraction between molecules through hydrogen atoms bonded to nitrogen, oxygen or fluorine

  2. Why are metals good conductors of electricity?

    • Delocalised electrons can move through the lattice
    • Metal lattices contain molecules with dipoles
    • Metal atoms share electrons in covalent bonds
    • Positive ions can move freely through the lattice

  3. What does the term 'sea of delocalised electrons' describe?

    • Electrons localised in covalent molecules
    • Lone pairs on oxygen in water
    • Delocalised electrons in a metallic lattice
    • Electrons in an ionic compound

  4. Why are metals malleable?

    • Covalent bonds break and reform easily between layers of atoms, so the structure bends without shattering or cracking
    • Ions are arranged in rigid molecules that bend under stress, so the whole structure can deform easily without breaking
    • Layers of positive ions can slide over each other while the delocalised electrons still hold the structure together
    • Metals contain no bonds at all, so the atoms are held only by weak forces that let them move freely past each other

  5. Which metal in Period 3 has the most delocalised electrons per atom?

    • Magnesium
    • Aluminium
    • Potassium
    • Sodium

  6. Metals have high melting points. Which best explains the strength of metallic bonding?

    • Strong attraction between delocalised electrons and the positive ions
    • Dative covalent bonds between metal atoms, which share electron pairs from one atom
    • Weak van der Waals forces between the atoms, which hold them only loosely together
    • Hydrogen bonds between layers of metal atoms, which hold the layers together firmly

  7. Which statement about a metallic lattice is true?

    • The lattice contains negative ions that are strongly attracted to the positive ions
    • Atoms form discrete molecules that are held together by intermolecular forces between them
    • Positive ions are arranged in a regular lattice surrounded by delocalised electrons
    • Electrons are localised between pairs of atoms, forming directional covalent bonds

  8. Compare sodium and magnesium. Which has the higher melting point, and why?

    • Sodium, because its ions are larger and attract electrons more strongly
    • Magnesium, because magnesium has covalent bonds
    • Magnesium, because it has more delocalised electrons per atom and stronger attraction
    • Sodium, because it has fewer delocalised electrons

  9. Which property of metals is best explained by the presence of delocalised electrons?

    • Fixed crystal shape in molten form
    • Solubility in water at room temperature
    • Brittleness when a force is applied
    • Electrical conductivity in the solid state

  10. Which ions and electrons are present in the metallic lattice of calcium?

    • Ca2- ions and bonding pairs of electrons
    • Ca2+ and Cl- ions
    • Ca+ ions and localised electrons
    • Ca2+ ions and delocalised electrons

  11. How many delocalised electrons does one atom of magnesium contribute to the metallic lattice?

    • 1
    • 0
    • 2
    • 3

  12. Which statement correctly compares metallic bonding with ionic bonding in sodium chloride?

    • Metallic bonding is non-directional, whereas ionic bonding is between specific oppositely charged ions in a fixed arrangement
    • Metallic bonding is a sharing of electron pairs, while ionic bonding has no electrons involved at all
    • Both involve discrete molecules that are held together by weak intermolecular forces in the lattice
    • Metallic bonds are always weaker than hydrogen bonds, so metals melt more easily than ionic solids

  13. Why do metals usually conduct heat well?

    • Positive ions have no charge, so they cannot carry energy from one part of the metal
    • Delocalised electrons carry kinetic energy through the lattice
    • Metals contain permanent dipoles that absorb heat and release it only at the surface
    • Metallic bonds are weak, so energy cannot pass between the atoms of the metal easily

  14. How many delocalised electrons does each atom of a Group 1 metal contribute?

    • Two
    • Three
    • None
    • One

  15. Why are alloys usually harder than pure metals?

    • Alloys contain no delocalised electrons, so the lattice is held together by fixed bonds
    • Atoms of different sizes disrupt the regular layers, so the layers slide less easily
    • Alloys contain a covalent network throughout, which locks every layer firmly in place
    • Alloys contain ionic bonds that are stronger than the metallic bonds in pure metals

  16. Why do Group 1 metals have lower melting points than Group 2 metals in the same period?

    • Group 1 atoms have more protons than Group 2 atoms, so the attraction on the delocalised electrons is weaker
    • Group 1 metals contain hydrogen bonds between layers, which are weaker than the metallic bonds of Group 2
    • Group 1 atoms contribute one delocalised electron each and form larger ions, so the metallic bonding is weaker
    • Group 1 metals contain covalent bonds that are much weaker than the metallic bonds of Group 2 metals

  17. Two metals have the same number of delocalised electrons per atom, but one has smaller positive ions. Which has stronger metallic bonding?

    • Neither, because metallic bonding is not electrostatic and so the ion size has no effect
    • The metal with smaller positive ions, because its ions are closer to the delocalised electrons
    • They are equal, because the number of delocalised electrons per atom is the same in both
    • The metal with larger positive ions, because the delocalised electrons are further away from them

  18. Why does the electrical conductivity of a metal decrease as its temperature increases?

    • Electrons leave the lattice at high temperature, which removes the carriers from the metal
    • The ions vibrate more, increasing collisions with the moving delocalised electrons
    • Delocalised electrons are used up as the temperature rises, so fewer carriers remain available
    • The ions become charged as the temperature rises, which reduces the number of free carriers

  19. How many moles of delocalised electrons are present in 1 mol of aluminium?

    • 27 mol
    • 1 mol
    • 3 mol
    • 2 mol

  20. Why do metals usually have high densities?

    • Metals contain hydrogen bonds between layers that pull the atoms closer together
    • Metals have many lone pairs of electrons that fill the spaces between atoms
    • Atoms are closely packed in a lattice, giving a large mass per unit volume
    • Metal atoms have covalent bonds that hold them tightly in place

All AQA Chemistry quizzes