Ionic bonding — AQA GCSE Chemistry
Verified courseTeaching note checked against the official specification.
Ionic bonding explained
Ionic bonding covers 7 specification points, including When a metal atom reacts with a non-metal atom… and Metal atoms lose electrons to become….
Read the full explanation
Open a point below for its explanation and revision guidance.
When a metal atom reacts with a non-metal atom electrons in the outer shell of the metal atom are transferred.
Ionic bonding occurs when a metal reacts with a non-metal, resulting in the complete transfer of outer shell electrons. Metal atoms have a small number of electrons in their highest energy level and lose them to achieve a stable, full outer shell, forming positively charged ions (cations). For example, a sodium atom loses its single outer electron to form a Na⁺ ion. These lost electrons are transferred directly to the outer shell of the non-metal atom, which gains them to complete its own outer shell, forming a negatively charged ion (anion), such as Cl⁻. This electron transfer creates the oppositely charged ions that form an ionic lattice.
Metal atoms lose electrons to become positively charged ions.
Metal atoms have relatively few electrons in their outer shell, so losing those electrons is energetically favourable and produces a positive ion. For example, a sodium atom has the electronic structure 2,8,1; losing its single outer electron gives Na⁺ with the structure 2,8, which is the same as neon. The number of electrons lost equals the group number for Groups 1, 2 and 3, so magnesium (2,8,2) loses two electrons to form Mg²⁺. The charge is written as a superscript after the symbol, and the size of the charge matches the number of electrons lost. This electron loss is the first half of ionic bonding and explains why metals form cations in compounds such as NaCl and MgO.
Non-metal atoms gain electrons to become negatively charged ions.
Non-metal atoms have outer shells that are closer to full, so gaining electrons is energetically favourable and produces a negative ion. For example, a chlorine atom has the electronic structure 2,8,7; gaining one electron gives Cl⁻ with the structure 2,8,8, the same as argon. The number of electrons gained is usually 8 minus the group number for Groups 6 and 7, so oxygen (2,6) gains two electrons to form O²⁻ and chlorine gains one to form Cl⁻. The charge is written as a superscript after the symbol, and the size of the negative charge equals the number of electrons gained. This electron gain is the second half of ionic bonding and explains the anions in compounds such as NaCl and MgO.
The ions produced by metals in Groups 1 and 2 and by non-metals in Groups 6 and 7 have the electronic structure of a noble gas (Group 0).
Atoms form ions by transferring electrons so that their outer shell becomes full, giving the stable electronic structure of a noble gas. A sodium atom (2,8,1) loses one electron to form Na⁺ (2,8), like neon. A magnesium atom (2,8,2) loses two electrons to form Mg²⁺ (2,8), also like neon. A chlorine atom (2,8,7) gains one electron to form Cl⁻ (2,8,8), like argon, and an oxygen atom (2,6) gains two electrons to form O²⁻ (2,8), like neon. In each case the ion has a full outer shell, which is why these ions are stable. For metals, the group number tells you how many electrons are lost; for non-metals, the number gained is 8 minus the group number. The resulting ion always matches the nearest 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 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.
A dot and cross diagram shows how electrons are transferred when an ionic compound forms. For sodium chloride, draw the sodium atom with its outer electron as a dot and the chlorine atom with seven outer electrons as crosses. Move the dot from sodium to chlorine so that chlorine now has eight outer electrons, shown as seven crosses and one dot. Then draw the ions in square brackets with the charges outside: Na⁺ and Cl⁻. For magnesium oxide, magnesium loses two electrons to oxygen, so two dots move to oxygen and the ions are Mg²⁺ and O²⁻. The diagram must show full outer shells on both ions, correct charges, and square brackets around each ion. This method works for Group 1 and 2 metals with Group 6 and 7 non-metals.
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.
Atoms gain or lose electrons to reach a stable noble gas arrangement. Metals in Group 1 lose one electron, giving ions with a 1⁺ charge, such as Na⁺. Metals in Group 2 lose two electrons, giving ions with a 2⁺ charge, such as Mg²⁺. Non-metals in Group 6 gain two electrons, giving ions with a 2⁻ charge, such as O²⁻. Non-metals in Group 7 gain one electron, giving ions with a 1⁻ charge, such as Cl⁻. The size of the charge equals the number of electrons gained or lost, and the sign shows whether electrons were lost (positive) or gained (negative). The group number therefore predicts the charge: Group 1 gives 1⁺, Group 2 gives 2⁺, Group 6 gives 2⁻ and Group 7 gives 1⁻.
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.
To work out an ion charge, first locate the element in the periodic table and note its group. A Group 1 metal loses one electron, so its ion has a 1⁺ charge, as in Na⁺. A Group 2 metal loses two electrons, so its ion has a 2⁺ charge, as in Mg²⁺. A Group 6 non-metal gains two electrons, so its ion has a 2⁻ charge, as in O²⁻. A Group 7 non-metal gains one electron, so its ion has a 1⁻ charge, as in Cl⁻. The number of electrons gained or lost equals the distance from the nearest noble gas group, and the sign shows whether electrons are lost or gained. This method applies only to the groups named in the statement.
Your focus
- Describe the transfer of outer shell electrons from metal atoms to non-metal atoms.
- Determine the charge of the ions formed when specific metals and non-metals react.
- Illustrate the electron transfer process using dot and cross diagrams.
Show all 21 objectives
- State that metal atoms lose electrons to form positive ions.
- Deduce the charge on a metal ion from its group number.
- Write the symbol and charge of a metal ion correctly using superscript notation.
- State that non-metal atoms gain electrons to form negative ions.
- Deduce the charge on a non-metal ion from its group number.
- Write the symbol and charge of a non-metal ion correctly using superscript notation.
- Describe how metal and non-metal atoms form ions with noble gas electronic structures.
- Deduce the electronic structure of an ion from the group number of the element.
- Compare the electronic structure of an ion with that of a noble gas.
- Draw dot and cross diagrams for ionic compounds formed from Group 1 and 2 metals with Group 6 and 7 non-metals.
- Show electron transfer and the resulting ions with correct charges and full outer shells.
- Interpret a dot and cross diagram to identify the ions present in an ionic compound.
- State the charge on ions formed by elements in Groups 1, 2, 6 and 7.
- Explain how the group number determines the number of electrons gained or lost.
- Write the symbol and charge for common ions such as Na⁺, Mg²⁺, O²⁻ and Cl⁻.
- Deduce the charge on an ion from the group number of its element.
- Apply the rule correctly to Groups 1, 2, 6 and 7 only.
- Write ion symbols with the correct superscript charge, such as Mg²⁺ and O²⁻.
Ionic bonding flashcards
Card 1 of 11
Metal atoms and ions
Ionic bonding exam tips
Marking Points
- State that metals lose electrons from their outer shell when reacting with non-metals.
- State that non-metals gain the electrons transferred from the metal atoms.
- Explain that the loss of electrons results in the formation of positive ions.
- Explain that the gain of electrons results in the formation of negative ions.
- Recognise that electron transfer allows both atoms to achieve a full outer electron shell.
- Metal atoms lose the electrons in their outermost shell first, because these are the least strongly attracted to the nucleus.
- The number of electrons lost by a metal atom is equal to its group number for Groups 1, 2 and 3.
- Loss of electrons produces a positive ion because the number of protons now exceeds the number of electrons.
- The charge on the ion is written as a superscript, for example Na⁺, Mg²⁺ and Al³⁺.
- The electronic structure of the ion formed is that of the nearest noble gas, for example Na⁺ is 2,8 like neon.
- Non-metal atoms gain electrons into their outermost shell until that shell is full.
- The number of electrons gained is 8 minus the group number for Groups 6 and 7.
- Gain of electrons produces a negative ion because the number of electrons now exceeds the number of protons.
- The charge on the ion is written as a superscript, for example O²⁻, F⁻ and Cl⁻.
- The electronic structure of the ion formed is that of the nearest noble gas, for example Cl⁻ is 2,8,8 like argon.
- Atoms lose or gain electrons so that their outer shell becomes full.
- Metal atoms in Groups 1 and 2 lose one or two electrons respectively to form positive ions.
- Non-metal atoms in Groups 6 and 7 gain two or one electrons respectively to form negative ions.
- The ion formed has the same electronic structure as the nearest noble gas in Group 0.
- Examples include Na⁺ and Mg²⁺ with the structure 2,8, and Cl⁻ with the structure 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 electrons from the metal atom to the non-metal atom.
- Draw square brackets around each ion and write the charge outside the bracket.
- Ensure each ion has a full outer shell of eight electrons, except for small ions such as Li⁺ which have two.
- Use the correct number of electrons transferred, for example two for MgO and one for NaCl.
- Group 1 metals lose one electron and form ions with a 1⁺ charge, for example Na⁺ and K⁺.
- Group 2 metals lose two electrons and form ions with a 2⁺ charge, for example Mg²⁺ and Ca²⁺.
- Group 6 non-metals gain two electrons and form ions with a 2⁻ charge, for example O²⁻ and S²⁻.
- Group 7 non-metals gain one electron and form ions with a 1⁻ charge, for example Cl⁻ and Br⁻.
- The charge is determined by the number of electrons gained or lost to reach a stable noble gas electron arrangement.
- The sign of the charge shows loss of electrons (positive) or gain of electrons (negative).
- Identify the group of the element from the periodic table before assigning a charge.
- Group 1 metals form 1⁺ ions by losing one electron, for example Li⁺ and K⁺.
- Group 2 metals form 2⁺ ions by losing two electrons, for example Mg²⁺ and Ca²⁺.
- Group 6 non-metals form 2⁻ ions by gaining two electrons, for example O²⁻ and S²⁻.
- Group 7 non-metals form 1⁻ ions by gaining one electron, for example F⁻ and Br⁻.
- Use the noble gas electron arrangement to justify the number of electrons gained or lost.
Examiner Tips
- 💡Use dot and cross diagrams to clearly show the movement of electrons from the metal's outer shell to the non-metal's outer shell.
- 💡Always state the charge of the resulting ions, ensuring you use the correct superscript notation (e.g., Mg²⁺, O²⁻).
- 💡Remember that the number of electrons lost by the metal must equal the number of electrons gained by the non-metal, which may require multiple atoms (e.g., MgCl₂).
- 💡State clearly that electrons are lost from the outer shell, not from the nucleus.
- 💡Link the group number to the number of electrons lost before writing the ion symbol.
- 💡Check that the superscript charge matches the number of electrons removed, and that the sign is positive.
- 💡Count the electrons needed to complete the outer shell before writing the ion.
- 💡Use the group number to work out the charge quickly: Group 7 forms 1⁻ ions and Group 6 forms 2⁻ ions.
- 💡Check that the superscript charge is negative and matches the number of electrons gained.
- 💡Write the electron configuration of the atom first, then add or remove the outer electrons.
- 💡Compare the ion configuration with the nearest noble gas to check it is correct.
- 💡Use dots for one element and crosses for the other so the transfer is easy to see.
- 💡Check the charge on each ion matches the number of electrons lost or gained.
- 💡Include square brackets and charges for every ion in the final diagram.
- 💡Link each group to the number of electrons gained or lost before writing the charge.
- 💡Check that the total positive charge balances the total negative charge when writing a formula.
- 💡Practise writing ions for the first four elements of each named group so the pattern becomes automatic.
- 💡Write the electron loss or gain in words before writing the ion symbol.
- 💡For a formula, balance the total positive and negative charges to check the ion charges are consistent.
- 💡Learn one example from each group so you can recall the pattern quickly.
Common Mistakes
- Writing ionic charges incorrectly, such as Na+ or Cl-; correction: always use superscript signs for ionic charges, such as Na⁺ and Cl⁻.
- Stating that non-metals transfer electrons to metals; correction: electrons are always transferred from the metal atom to the non-metal atom.
- Describing the process as sharing electrons; correction: ionic bonding involves the complete transfer of electrons, not sharing.
- Writing the charge after the symbol on the baseline, such as Na1⁺, instead of using a superscript: the correct form is Na⁺.
- Saying that metal atoms gain electrons to become positive ions: metals lose electrons, and gaining electrons would make the ion negative.
- Forgetting to balance the charge with the number of electrons lost, for example writing Mg⁺ when magnesium loses two electrons and forms Mg²⁺.
- Writing the charge on the baseline, such as Cl1⁻, instead of using a superscript: the correct form is Cl⁻.
- Saying that non-metal atoms lose electrons to become negative ions: non-metals gain electrons, and losing electrons would make the ion positive.
- Miscounting the electrons gained, for example giving oxygen a single negative charge when it gains two electrons to form O²⁻.
- Thinking that the ion has the same number of electrons as the original atom: the number of electrons changes by the number lost or gained.
- Confusing the noble gas matched by an ion, for example saying Mg²⁺ has the structure of argon when it is 2,8 like neon.
- Writing the electron configuration of the atom instead of the ion, for example writing 2,8,7 for a chloride ion instead of the correct 2,8,8.
- Drawing the ions without square brackets or charges, which fails to show that ions have formed.
- Showing the wrong number of electrons transferred, for example moving one electron from magnesium instead of two.
- Leaving the outer shell of an ion incomplete, for example showing Cl⁻ with seven outer electrons instead of eight.
- Writing the charge after the symbol without a superscript sign, for example Na1⁺ instead of Na⁺; the charge belongs as a superscript sign after the symbol.
- Assuming all metal ions have a 1⁺ charge; Group 2 metals such as magnesium form 2⁺ ions because they lose two electrons.
- Giving Group 6 non-metals a 6⁻ charge by copying the group number directly; they gain two electrons, so the charge is 2⁻.
- Using the group number directly as the charge for Group 6 elements, giving 6⁻ instead of 2⁻; the correct charge is 2⁻ because two electrons are gained.
- Forgetting the sign on the charge, writing Mg² instead of Mg²⁺; the sign shows the direction of electron transfer.
- Applying the method to transition metals or Group 4 elements, which do not follow this simple pattern; the rule is limited to Groups 1, 2, 6 and 7.