Welcome to your study notes on covalent bonding! This is a core topic in GCSE Chemistry that will help you understand how non-metal atoms join together to create everything from the water we drink to the oxygen we breathe.
What you'll learn:
- How covalent bonds are formed by sharing electrons.
- The scale and size of individual atoms and simple covalent molecules.
- How to construct and interpret dot-and-cross diagrams for key molecules like methane, water, and carbon dioxide.
1. What is a Covalent Bond?
In your earlier studies, you learned that atoms are most stable when they have a full outer shell of electrons (a stable octet, like the noble gases in Group 0).
When non-metal atoms react together, they cannot transfer electrons to each other because both atoms need to gain electrons to complete their outer shells. Instead, they share pairs of electrons.
Covalent bond
A covalent bond is the strong electrostatic attraction between a shared pair of electrons and the positively charged nuclei of the bonding atoms.
Because electrons are negatively charged and atomic nuclei are positively charged (due to protons), there is a powerful electrostatic force pulling them together. This shared pair of electrons acts like a glue holding the atoms together.
Molecules
Covalent bonding always results in the formation of molecules. A molecule is a small group of two or more atoms chemically joined together by covalent bonds.
2. Size and Scale of Atoms and Molecules
Before we start drawing these bonds, we need to appreciate how incredibly tiny atoms and simple molecules are. In your exam, you are expected to recall the typical order of magnitude (approximate size scale) of these particles.
- Typical Atom: An individual atom has a diameter of about 1×10−10 m1 \times 10^{-10}\text{ m}1×10−10 m (which is 0.1 nanometres, or 0.1 nm).
- Small Molecules: Small molecules made of a few covalently bonded atoms (like water, H₂O) have a typical size of about 1×10−10 m1 \times 10^{-10}\text{ m}1×10−10 m to 1×10−9 m1 \times 10^{-9}\text{ m}1×10−9 m (0.1 to 1 nm).
Sizing up standard form
Do not let negative indices confuse you! Remember that 1×10−10 m1 \times 10^{-10}\text{ m}1×10−10 m is smaller than 1×10−9 m1 \times 10^{-9}\text{ m}1×10−9 m. Each time the negative power increases by one (gets closer to zero), the size increases by a factor of 10.
Here is how you can use these scales to make comparison calculations in your exam.
Comparing atomic and molecular scales
An oxygen atom has an approximate diameter of 1.2×10−10 m1.2 \times 10^{-10}\text{ m}1.2×10−10 m. A complex protein molecule has a length of approximately 6.0×10−9 m6.0 \times 10^{-9}\text{ m}6.0×10−9 m. Calculate how many times larger the protein molecule is compared to the individual oxygen atom.
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Identify the values to compare: The diameter of the oxygen atom is d=1.2×10−10 md = 1.2 \times 10^{-10}\text{ m}d=1.2×10−10 m. The length of the protein molecule is L=6.0×10−9 mL = 6.0 \times 10^{-9}\text{ m}L=6.0×10−9 m.
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Set up the ratio equation: To find out how many times larger the protein is, divide the larger size by the smaller size:
- Calculate the final scale factor: Divide the coefficients:
Subtract the powers of 10:
−9−(−10)=1 -9 - (-10) = 1 −9−(−10)=1Combine these to get the final ratio:
5×101=50 5 \times 10^1 = 50 5×101=50The protein molecule is 50 times larger than the oxygen atom.
3. Dot-and-Cross Diagrams
We represent covalent bonds using dot-and-cross diagrams. In these diagrams:
- The atoms are drawn as overlapping circles representing their outer electron shells.
- Electrons from one atom are drawn as dots (∙\bullet∙).
- Electrons from the other atom are drawn as crosses (×\times×).
- The electrons sitting inside the overlapping region represent the shared covalent bond.
The Edexcel specification requires you to know how to draw and explain six specific molecules.
The 6 Key Molecules to Know:
- Hydrogen (H₂): Two hydrogen atoms share one pair of electrons, forming a single covalent bond. Each hydrogen atom now has 2 electrons in its outer shell (which is full for the first shell).
- Hydrogen chloride (HCl): A hydrogen atom shares its 1 electron with a chlorine atom (which has 7 outer electrons). This forms a single covalent bond.
- Water (H₂O): An oxygen atom (6 outer electrons) shares electrons with two separate hydrogen atoms, forming two single covalent bonds.
- Methane (CH₄): A carbon atom (4 outer electrons) shares electrons with four separate hydrogen atoms, forming four single covalent bonds.
- Oxygen (O₂): Two oxygen atoms (each with 6 outer electrons) must share two pairs of electrons to fill their outer shells. This is called a double covalent bond.
- Carbon dioxide (CO₂): A central carbon atom (4 outer electrons) shares two pairs of electrons with each of two oxygen atoms. This forms two double covalent bonds.
The diagram below displays the correct dot-and-cross structures for these six key molecules.
Sharing tools
Think of covalent bonding like two neighbours sharing garden tools. Neighbour C has a lawnmower but needs a hedge trimmer. Neighbour O has a hedge trimmer but needs a lawnmower. Instead of buying new ones, they keep both tools in a shared shed built on their boundary line. Both neighbours now have access to both tools, keeping them both happy and "bonded" together.
Let's walk through exactly how to construct one of the trickier double-bonded structures step-by-step.
Drawing the dot-and-cross diagram for carbon dioxide
Draw the dot-and-cross diagram for a molecule of carbon dioxide, CO₂, showing outer shell electrons only.
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Find the number of outer shell electrons for each element: Check your Periodic Table. Carbon is in Group 4, so it has 4 outer shell electrons. Oxygen is in Group 6, so each oxygen atom has 6 outer shell electrons.
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Work out how many bonds each atom must form: Carbon needs 4 more electrons to achieve a stable octet (8 electrons), so it must share all 4 of its outer electrons. Each oxygen atom needs 2 more electrons to reach 8, so each oxygen must share 2 of its outer electrons.
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Arrange the atoms and draw overlapping circles: Since carbon needs the most bonds, put it in the centre with an oxygen atom on either side. Draw overlapping circles to represent the shared outer shells: O—C—O.
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Distribute the shared electrons in the overlapping areas: Carbon shares 2 of its electrons (crosses) with the left oxygen, and 2 with the right oxygen. Each oxygen shares 2 of its electrons (dots) with the carbon. Place 2 dots and 2 crosses in each of the two overlapping sections. This represents the two double covalent bonds.
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Draw the remaining non-bonding outer electrons: Carbon has shared all 4 of its outer electrons, so it has none left outside the overlaps. Each oxygen started with 6 outer electrons and shared 2, leaving 4 non-bonding outer electrons on each oxygen. Draw these 4 dots on the outer, non-overlapping parts of each oxygen circle. Verify your final drawing by counting: the carbon circle has 8 electrons total, and each oxygen circle has 8 electrons total.
Forgetting non-bonding electrons
A very common exam mistake is forgetting to draw the non-bonding outer electrons (the ones not involved in the bond). For example, in a water molecule, you must draw the 4 non-bonding electrons on the oxygen atom, not just the shared pairs in the overlaps!
Outer shells only!
Unless a question specifically asks you to draw "all electron shells", only draw the outermost shell. Drawing inner shells takes up valuable exam time and increases the risk of making counting mistakes.
In the exam
- Count at the end: Always do a quick final count of the electrons in each atom's outer circle. Hydrogen circles must contain exactly 2 electrons. All other non-metal circles (carbon, nitrogen, oxygen, halogens) must contain exactly 8 electrons.
- Be distinct with dots and crosses: Make sure your dots (∙\bullet∙) and crosses (×\times×) are drawn clearly and are positioned unambiguously inside or outside the overlapping region. If the examiner cannot tell where an electron sits, they cannot award the mark.
- Use the exact definition: If asked to explain how a covalent bond is formed, always mention the "shared pair of electrons" and the "electrostatic attraction between these negative electrons and the positive nuclei."
Check yourself
- State the typical size range of a small covalent molecule in metres.
- How many pairs of electrons are shared between the atoms in a molecule of oxygen, O₂?
- Why does carbon form four single covalent bonds in methane, CH₄, while oxygen only forms two single covalent bonds in water, H₂O?
