AQA GCSE Chemistry: Chemical Bonds and Ionic Bonding — Practice Questions

These are original questions written for Marlbridge, in the style and at the standard of the examination. They are not reproduced past-paper questions — examination boards hold copyright in their own papers. Use these alongside the official past papers available free from your board.

Related: Chemical Bonds and Ionic Bonding study guide | Chemical Bonds and Ionic Bonding revision notes

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Section A

1. Name the three types of strong chemical bond. [3]

2. State the specific named force that holds oppositely charged ions together in an ionic compound. [1]

Section B

3. A compound is formed between calcium and chlorine. Identify the type of bonding present, and justify your answer using the positions of the elements involved. [3]

4. State the two conventions that must be followed when drawing a correct dot and cross diagram for an ionic compound. [2]

5. A potassium atom (Group 1) reacts with a fluorine atom (Group 7). Explain, in terms of electron transfer, how the ionic bond forms, naming the electronic structure each resulting ion has. [5]

6. Work out the charge on the ion formed by a Group 2 metal and the charge on the ion formed by a Group 6 non-metal, explaining your reasoning for each. [4]

7. A student draws a dot and cross diagram for sodium chloride but shows chlorine gaining two electrons to form a 2− ion, and omits the square brackets and charges around each ion. Identify the two errors and explain why each loses marks. [4]

8. Explain, in terms of electron transfer, how the ionic bond forms in calcium oxide (calcium is in Group 2, oxygen is in Group 6). [5]

9. Ionic compounds such as sodium chloride have high melting points. Explain why, linking your answer to the bonding described in this sub-topic. [4]

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Answers

1. Ionic [1]; covalent [1]; metallic [1].

2. Electrostatic force(s) of attraction [1].

3. This is ionic bonding [1]. Calcium is a metal and chlorine is a non-metal [1], and ionic bonding occurs specifically in compounds formed from a metal combined with a non-metal [1].

4. Dots represent electrons from one atom and crosses represent electrons from the other atom [1]; only the outer (valence) shell electrons are shown, with each finished ion enclosed in square brackets with its charge [1].

5. Potassium (Group 1) has 1 electron in its outer shell and loses it, forming a K+ ion with the electronic structure of argon [1] [1]. Fluorine (Group 7) has 7 electrons in its outer shell and gains the 1 electron lost by potassium, forming an F− ion with the electronic structure of neon [1] [1]. The oppositely charged K+ and F− ions are then held together by strong electrostatic forces of attraction, forming the ionic bond [1].

6. A Group 2 metal loses 2 electrons, forming a 2+ ion [1] [1]. A Group 6 non-metal gains 2 electrons, forming a 2− ion [1] [1].

7. Error 1: chlorine (Group 7) has 7 outer electrons and needs to gain only 1 electron, forming a 1− ion with the electronic structure of argon, not a 2− ion [1] [1]. Error 2: omitting the square brackets and charges loses marks because a finished dot and cross diagram must show each ion as a distinct charged species, not simply atoms with electrons redistributed [1] [1].

8. Calcium (Group 2) has 2 electrons in its outer shell and loses both, forming a Ca2+ ion with the electronic structure of argon [1] [1]. Oxygen (Group 6) has 6 electrons in its outer shell and gains the 2 electrons lost by calcium, forming an O2− ion with the electronic structure of neon [1] [1]. The oppositely charged Ca2+ and O2− ions are held together by strong electrostatic forces of attraction, forming the ionic bond in calcium oxide [1].

9. The strong electrostatic forces of attraction between oppositely charged ions extend in all directions through a giant ionic lattice, not just between one pair of ions [1] [1]. Melting an ionic compound requires overcoming a very large number of these strong ionic bonds throughout the structure simultaneously [1], which needs a large amount of energy, giving ionic compounds like sodium chloride their characteristically high melting points [1].

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Where marks are usually lost

• Naming a bond type without justifying it from the elements involved (metal/non-metal for ionic, non-metal/non-metal for covalent, metal alone for metallic).
• Getting the number of electrons transferred wrong so the ion formed does not match a noble gas structure, or omitting the square brackets and charge around each finished ion.
• Applying the “charge equals group number” rule to elements outside Groups 1, 2, 6 and 7.
• Describing ionic bond formation as a single step (“electrons move”) rather than the required two-step sequence of electron transfer followed by electrostatic attraction.
• Explaining a property such as melting point without linking it back to the strength and number of ionic bonds in the lattice.

Approaching chemical bonding and ionic bonding questions

Always justify a bond-type identification from the elements’ positions in the periodic table rather than simply naming the bond, since full marks depend on showing the metal/non-metal (or non-metal/non-metal, or metal-alone) reasoning explicitly. For any electron-transfer explanation, work through the same three-part structure every time: state each element’s outer-shell electron count and group, state how many electrons are lost or gained and the resulting ion charge, then name the noble gas electronic structure each ion now matches, before finally stating that electrostatic attraction between the oppositely charged ions is what forms the bond — this sequence mirrors exactly how the specification itself builds the explanation, and applying it consistently to an unfamiliar element pair, not just to memorised examples like sodium chloride, is what the exam actually tests.