Ionic bonding — AQA GCSE Combined Science
Test yourself on Ionic bonding with AQA GCSE practice questions.
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Ionic bonding explained
Ionic bonding occurs when a metal atom reacts with a non-metal atom.
Read the full explanation
The metal atom transfers electrons from its outer shell to the non-metal atom. By losing electrons, the metal atom forms a positive ion; by gaining electrons, the non-metal atom forms a negative ion. For example, sodium (2,8,1) loses one electron to become Na⁺ with the electronic structure 2,8, matching neon. Chlorine (2,8,7) gains one electron to become Cl⁻ with the electronic structure 2,8,8, matching argon. Metals in Group 1 lose one electron to form 1⁺ ions; Group 2 metals lose two electrons to form 2⁺ ions. Non-metals in Group 6 gain two electrons to form 2⁻ ions; Group 7 non-metals gain one electron to form 1⁻ ions. In each case, the resulting ion has the same electronic structure as a noble gas.
The electron transfer during the formation of an ionic compound can be represented by a dot and cross diagram, eg for sodium chloride
A dot and cross diagram shows the electron transfer that forms an ionic compound. For sodium chloride, draw the sodium atom with its outer-shell electron as a dot and the chlorine atom with its outer-shell electrons as crosses. Show the transfer of the sodium outer-shell electron to the chlorine atom. After transfer, draw the sodium ion with no outer-shell electron and a 1⁺ charge, and the chloride ion with eight outer-shell electrons (the original seven crosses plus the transferred dot) and a 1⁻ charge. Use square brackets around each ion and include the charge outside the bracket. The diagram represents the electron transfer and the resulting ions, not the sharing of electrons.
Students should be able to draw dot and cross diagrams for ionic compounds formed by metals in Groups 1 and 2 with non-metals in Groups 6 and 7.
Ionic bonding joins a metal from Group 1 or 2 to a non-metal from Group 6 or 7. Draw the outer shells only, using dots for one atom's electrons and crosses for the other's. A Group 1 atom loses its single outer electron to become a 1⁺ ion; a Group 2 atom loses two to become 2⁺. A Group 6 atom gains two electrons to become 2⁻; a Group 7 atom gains one to become 1⁻. Show the transfer with an arrow, then draw each ion with a full outer shell and square brackets carrying its charge. For sodium oxide, two Na atoms each give one electron to one O atom, giving 2Na⁺ and O²⁻. For magnesium chloride, one Mg atom gives one electron to each of two Cl atoms, giving Mg²⁺ and 2Cl⁻.
The charge on the ions produced by metals in Groups 1 and 2 and by non-metals in Groups 6 and 7 relates to the group number of the element in the periodic table.
An element's group number predicts the charge on its ion. Metals in Group 1 lose their single outer electron and form 1⁺ ions, such as Na⁺ and K⁺. Metals in Group 2 lose two outer electrons and form 2⁺ ions, such as Mg²⁺ and Ca²⁺. Non-metals in Group 6 gain two electrons to complete their outer shell and form 2⁻ ions, such as O²⁻ and S²⁻. Non-metals in Group 7 gain one electron and form 1⁻ ions, such as Cl⁻ and Br⁻. The pattern arises because atoms gain or lose electrons to reach a full outer shell. For Group 1 and 2 metals, the positive charge equals the group number. For Group 6 and 7 non-metals, the negative charge equals eight minus the group number.
Students should be able to work out the charge on the ions of metals and non-metals from the group number of the element, limited to the metals in Groups 1 and 2, and non-metals in Groups 6 and 7.
Atoms in Groups 1 and 2 lose their outer-shell electrons to form positive ions, while atoms in Groups 6 and 7 gain electrons to complete their outer shell and form negative ions. The group number tells you how many outer electrons a metal loses, whereas for non-metals, 8 minus the group number gives the number of electrons gained. For example, sodium (Group 1) loses one electron to form Na⁺, magnesium (Group 2) loses two to form Mg²⁺, oxygen (Group 6) gains two to form O²⁻, and chlorine (Group 7) gains one to form Cl⁻. The size of the charge equals the number of electrons transferred, and the sign shows whether electrons were lost or gained. This lets you predict ion charges directly from the periodic table.
Your focus
- Describe how electrons are transferred when a metal atom reacts with a non-metal atom.
- Explain how metal atoms form positive ions and non-metal atoms form negative ions.
- Deduce the charge on ions formed by elements in Groups 1, 2, 6 and 7 and relate this to noble gas electronic structures.
Show all 15 objectives
- Draw a dot and cross diagram to show the electron transfer in sodium chloride.
- Represent the ions formed in an ionic compound with correct charges and electron arrangements.
- Interpret a dot and cross diagram to identify the electron transfer and the resulting ions.
- Draw dot and cross diagrams for compounds formed from Group 1 or 2 metals and Group 6 or 7 non-metals.
- Use dots and crosses to show which atom supplied each transferred electron.
- Show the resulting ions with full outer shells and correct charges.
- State the charge on ions formed by elements in Groups 1, 2, 6 and 7.
- Explain how the group number relates to the charge on the ion formed.
- Predict the charge on an ion from the element's position in the periodic table.
- State the charge on ions formed by Group 1, Group 2, Group 6 and Group 7 elements.
- Explain how the group number determines the number of electrons lost or gained.
- Write correct ion symbols with superscript charges for these elements.
Ionic bonding exam tips
Marking Points
- Electrons in the outer shell of the metal atom are transferred to the non-metal atom during reaction.
- Metal atoms lose electrons and become positively charged ions.
- Non-metal atoms gain electrons and become negatively charged ions.
- Group 1 metals lose one electron to form 1⁺ ions; Group 2 metals lose two electrons to form 2⁺ ions.
- Group 6 non-metals gain two electrons to form 2⁻ ions; Group 7 non-metals gain one electron to form 1⁻ ions.
- The ions formed have the electronic structure of a noble gas, such as Na⁺ (2,8) and Cl⁻ (2,8,8).
- Draw the outer-shell electrons of each atom using dots for one atom and crosses for the other.
- Show the transfer of one electron from the sodium atom to the chlorine atom.
- Draw the sodium ion with no outer-shell electrons and a 1⁺ charge.
- Draw the chloride ion with eight outer-shell electrons and a 1⁻ charge.
- Use square brackets around each ion and place the charge outside the bracket.
- Ensure the diagram represents electron transfer, not electron sharing.
- Outer-shell electrons only are shown; inner shells may be omitted or drawn as a labelled circle.
- Dots represent electrons from one atom and crosses those from the other, so the transfer is traceable.
- Group 1 and Group 2 metals lose one or two electrons respectively, forming 1⁺ and 2⁺ ions.
- Group 6 and Group 7 non-metals gain two or one electrons respectively, forming 2⁻ and 1⁻ ions.
- The number of metal and non-metal atoms drawn gives overall neutrality, for example two Na atoms per O atom in sodium oxide.
- Each ion is drawn with a complete outer shell and its charge shown, commonly in square brackets.
- An arrow or equivalent shows electron transfer from metal to non-metal.
- The diagram is labelled with the compound name or formula and the ion charges.
- Group 1 metals form ions with a 1⁺ charge by losing one outer electron.
- Group 2 metals form ions with a 2⁺ charge by losing two outer electrons.
- Group 6 non-metals form ions with a 2⁻ charge by gaining two electrons.
- Group 7 non-metals form ions with a 1⁻ charge by gaining one electron.
- The charge follows from the number of electrons transferred to reach a full outer shell.
- The group number can be used to predict the charge for these groups without recalling each ion individually.
- Ion charges are written with the number before the sign, for example 2⁺ and 2⁻, and with the sign for single charges, for example Na⁺ and Cl⁻.
- Group 1 metals lose one outer electron and form ions with a 1+ charge, such as Na⁺ and K⁺.
- Group 2 metals lose two outer electrons and form ions with a 2+ charge, such as Mg²⁺ and Ca²⁺.
- Group 6 non-metals gain two electrons and form ions with a 2- charge, such as O²⁻ and S²⁻.
- Group 7 non-metals gain one electron and form ions with a 1- charge, such as Cl⁻ and Br⁻.
- The number of electrons lost or gained equals the size of the charge, and the sign shows loss (positive) or gain (negative).
- Ion charges can be deduced from group number alone for these groups, without writing full electron configurations.
Examiner Tips
- 💡Link the group number of the metal to the number of electrons lost and the charge on the positive ion.
- 💡Link the group number of the non-metal to the number of electrons gained and the charge on the negative ion.
- 💡Check that the electron configuration of each ion matches that of a noble gas, such as 2,8 for Na⁺ and 2,8,8 for Cl⁻.
- 💡Use dots for one atom and crosses for the other so the transferred electron is easy to track.
- 💡Check that the total number of outer-shell electrons shown after transfer equals eight for the negative ion and zero for the positive ion.
- 💡Include the charge on each ion outside the square brackets to show the ionic compound formed.
- 💡Count electrons before and after transfer to check each ion has a full outer shell.
- 💡Check the total positive charge equals the total negative charge before finishing the diagram.
- 💡Label charges clearly, for example Na⁺ and Cl⁻, rather than leaving the reader to infer them.
- 💡Link each charge to the number of electrons lost or gained, not just to the group number.
- 💡Check that the charge sign matches whether the element is a metal or a non-metal.
- 💡Use the group number to predict charges for unfamiliar elements in Groups 1, 2, 6 and 7.
- 💡Link each group number to a fixed electron change: Group 1 loses 1, Group 2 loses 2, Group 6 gains 2, Group 7 gains 1.
- 💡Check the sign first, then the size of the charge, then write it as a superscript on the symbol.
- 💡Practise with unfamiliar elements in the same groups, such as potassium, calcium, sulfur and bromine, to show the pattern is general.
Common Mistakes
- Writing that metal atoms gain electrons to become positive ions; correct this by stating that metals lose electrons and become positive ions.
- Stating that non-metal atoms lose electrons to become negative ions; correct this by stating that non-metals gain electrons and become negative ions.
- Giving ion charges without the correct sign or number, such as writing Na²⁺ for sodium; correct this by using the group number to determine the charge, so Group 1 gives 1⁺ and Group 2 gives 2⁺.
- Drawing the ions without square brackets or charges; correct this by enclosing each ion in square brackets and writing the charge outside, such as [Na]⁺ and [Cl]⁻.
- Showing the transferred electron as a shared pair between the atoms; correct this by showing the electron fully transferred to the non-metal atom.
- Giving the chloride ion seven outer-shell electrons instead of eight; correct this by adding the transferred electron to the seven original outer-shell electrons.
- Drawing full inner shells for every atom, which obscures the outer-shell transfer; show only the outer shell or a labelled inner circle.
- Giving a Group 2 metal a 1⁺ charge; magnesium loses two electrons, so its ion is Mg²⁺.
- Writing oxide as O²⁻ but then drawing only one sodium ion; two Na⁺ ions are needed to balance one O²⁻.
- Writing the sign before the number for ion charges, such as +2 instead of 2⁺; the correct format is number then sign.
- Assuming all metals form positive ions with the same charge; Group 1 and Group 2 metals differ because they lose different numbers of electrons.
- Giving Group 6 non-metals a 6⁻ charge by copying the group number; they gain two electrons, so the charge is 2⁻.
- Writing the charge after the symbol as a plain number, for example Na1; the correct notation places the charge as a superscript, Na⁺.
- Giving Group 6 non-metals a 6- charge by copying the group number; the correct charge is 2- because they gain two electrons (8 minus 6).
- Assuming all metals form 1+ ions; Group 2 metals such as magnesium form 2+ ions because they lose two electrons.