Ionic bonding
What you'll learn:
- Why atoms transfer electrons to form ions.
- How to predict the charge of an ion using its Group number on the periodic table.
- How to draw dot and cross diagrams to show the formation of ionic compounds.
Why atoms react
Most atoms do not have a full outer shell of electrons, which makes them unstable. To become stable, they react with other atoms to achieve the electronic structure of a noble gas (Group 0) — which means having a completely full outer shell. In ionic bonding, atoms achieve this by giving away or taking in electrons.
Forming ions
Ionic bonding always happens between a metal and a non-metal.
When these atoms react, the metal atom transfers its outer electrons to the non-metal atom.
- Metal atoms lose electrons to empty their outer shell. Because electrons are negatively charged, losing them leaves the metal with a positive charge.
- Non-metal atoms gain electrons to fill their outer shell. Gaining negative electrons gives the non-metal a negative charge.
Ion
An ion is an atom (or group of atoms) that has an electrical charge because it has lost or gained one or more electrons.
Because opposite electrical charges attract each other strongly, the positive metal ions and negative non-metal ions stick together. This strong electrostatic attraction is the ionic bond.
The electron swap
Think of it like a sticker collection. The metal has one or two spare stickers that don't fit in its album, and the non-metal has one or two blank spaces left to complete its set. The metal gives its spares to the non-metal, so both finish with perfectly completed collections.
Predicting charges from Group numbers
You don't need to memorise the charge of every ion. You can work it out by finding the element's Group number on the periodic table. The Group number tells you how many electrons are in the outer shell.
For the metals (Groups 1 and 2):
- Group 1 metals (like sodium and potassium) have 1 outer electron. They lose 1 electron to form a 1+1+1+ ion (e.g. Na+\text{Na}^+Na+).
- Group 2 metals (like magnesium and calcium) have 2 outer electrons. They lose 2 electrons to form a 2+2+2+ ion (e.g. Mg2+\text{Mg}^{2+}Mg2+).
For the non-metals (Groups 6 and 7):
- Group 7 non-metals (like fluorine and chlorine) have 7 outer electrons. They need 1 more electron to reach a full 8. They gain 1 electron to form a 1−1-1− ion (e.g. Cl−\text{Cl}^-Cl−).
- Group 6 non-metals (like oxygen and sulfur) have 6 outer electrons. They need 2 more electrons to reach a full 8. They gain 2 electrons to form a 2−2-2− ion (e.g. O2−\text{O}^{2-}O2−).
The Rule of Full Shells
Metals lose all their outer shell electrons so the charge equals their Group number. Non-metals gain electrons to make their total up to 8, so their charge is always 8 minus their Group number.
Dot and cross diagrams
We use dot and cross diagrams to show the transfer of electrons. We draw the electrons for one atom as dots (∙\bullet∙) and the electrons for the other atom as crosses (×\times×). This helps us keep track of where the electrons came from, even though all electrons are exactly the same in reality.
To draw a fully reacted ionic compound:
- Draw the symbol of the metal ion, surrounded by square brackets, with its positive charge outside. (You usually leave its outer shell totally empty, showing it has lost its outer electrons).
- Draw the symbol of the non-metal ion, surrounded by a drawn shell containing 8 electrons. Show the gained electrons using the opposite symbol (e.g. if the non-metal had crosses, the gained electrons must be dots).
- Put square brackets around the non-metal ion and write its negative charge outside.
Below is a diagram showing how a sodium atom transfers its single outer electron to a chlorine atom to form sodium chloride (NaCl\text{NaCl}NaCl):
Forgetting the brackets and charges
A very common mistake in the exam is drawing the correct number of dots and crosses, but forgetting to add the square brackets [ ] and the + or - charges. Without the charges, you have just drawn two atoms side-by-side, not an ionic compound!
Sometimes the charges don't match up 1-to-1, which means you need more than one atom of an element to make the transfer work.
Drawing a dot and cross diagram for magnesium fluoride
- Identify the outer electrons for each element using their Group number. Magnesium is in Group 2, so it has 2 outer electrons. Fluorine is in Group 7, so it has 7 outer electrons.
- Determine how many electrons must be transferred to give every atom a full outer shell. The magnesium atom needs to lose 2 electrons, but a single fluorine atom only has space to gain 1. Therefore, you need exactly two fluorine atoms to accept the two electrons from one magnesium atom.
- Draw the final ions. Draw one Mg\text{Mg}Mg symbol in square brackets with a 2+2+2+ charge and an empty outer shell.
- Draw two F\text{F}F symbols, each inside square brackets with a 1−1-1− charge. Give each fluorine a full outer shell containing 7 of its own electrons (crosses) and 1 gained electron (a dot) from the magnesium.
In the exam
- Read the question carefully to see if it asks for the atoms (before transfer) or the ions (after transfer). Often, you only need to draw the final bonded ions.
- Only draw the outer shell of electrons unless the question specifically tells you to draw all the inner shells.
- Always include the square brackets and the charges on the top right for every ion you draw.
Check yourself
- What happens to the outer shell electrons of a Group 2 metal when it reacts with a non-metal?
- If an atom from Group 6 reacts, what charge will its resulting ion have?
- Why do we use dots for the electrons of one atom and crosses for the other?
