Lewis Structure of C2H4 (Ethene): Step-by-Step Guide

The Lewis structure C2H4 (ethene, also called ethylene) shows two carbon atoms joined by a double bond, with each carbon bonded to two hydrogen atoms. It uses 12 valence electrons, has no lone pairs, and gives each carbon a full octet and each hydrogen a full duet. This guide walks through each step so you can draw it yourself and understand why it looks the way it does.

How to draw the Lewis structure of C2H4 step by step

Step 1: Count the valence electrons

Carbon is in group 14 and has 4 valence electrons. Hydrogen has 1.

Atom	Number	Valence electrons each	Total
Carbon	2	4	8
Hydrogen	4	1	4
Total valence electrons			12

Every electron you draw must come from this total of 12. If you finish with more or fewer, something has gone wrong.

Step 2: Choose the central atoms

Hydrogen can form only one bond, so it is always a terminal atom and can never sit in the middle. That leaves the two carbon atoms as the skeleton. Place them next to each other and attach two hydrogens to each carbon.

Step 3: Draw single bonds and count what is used

Draw single bonds between the carbons and between each carbon and its hydrogens. That makes five bonds in total: one C-C bond and four C-H bonds. Each bond holds 2 electrons, so you have used 10 of your 12 electrons.

Step 4: Place the remaining electrons

There are 2 electrons left. If you place them on one carbon as a lone pair, that carbon would have six electrons around it and the other carbon only four, so neither would have a complete octet. Each carbon currently has only three bonds, giving it 6 electrons, so both are short of 8.

Step 5: Form the double bond

To give both carbons a full octet, share the last pair of electrons between them as a second bond. Now each carbon has:

• two single bonds to hydrogen (4 electrons),
• one double bond to the other carbon (4 electrons),
• a total of 8 electrons, so the octet rule is satisfied.

Each hydrogen shares one bond, so it has 2 electrons, which is a full outer shell for hydrogen. The final structure is H2C=CH2.

Step 6: Check using formal charge

Formal charge equals valence electrons minus nonbonding electrons minus half the bonding electrons.

Atom	Calculation	Formal charge
Each carbon	4 – 0 – 4	0
Each hydrogen	1 – 0 – 1	0

All formal charges are zero, which is the best sign that your Lewis structure is correct.

Shape and bonding of ethene

Molecular geometry

Using VSEPR theory, each carbon has three regions of electron density (two C-H bonds and the C=C double bond, which counts as one region). That gives a trigonal planar arrangement around each carbon with bond angles close to 120 degrees. In real ethene the H-C-H angle is slightly smaller, near 117 degrees, because the double bond region repels a little more strongly.

Hybridisation

Each carbon is sp2 hybridised. Its three sp2 orbitals form sigma bonds to two hydrogens and the other carbon, while the leftover p orbital overlaps sideways with the p orbital on the other carbon to form a pi bond. The double bond is therefore one sigma bond plus one pi bond, and the whole molecule lies in one plane.

Polarity

The C-H bonds are only slightly polar, and the symmetrical flat shape means any small dipoles cancel. Ethene is a nonpolar molecule. You can read more about the compound itself in the Ethylene article on Wikipedia.

Lewis structure of C2H4: bond summary

Feature	Value in C2H4
Total valence electrons	12
Sigma bonds	5 (four C-H and one C-C)
Pi bonds	1
Lone pairs	0
Geometry around each carbon	Trigonal planar
Hybridisation	sp2
Polarity	Nonpolar

Common mistakes to avoid

• Putting hydrogen in the centre. Hydrogen only makes one bond.
• Forgetting to count both carbons’ valence electrons.
• Leaving carbon with six electrons and calling the diagram finished.
• Adding lone pairs on carbon, which are not needed here.
• Writing a triple bond. That would be ethyne (C2H2), a different molecule.

Study tips

Practise by drawing a handful of related molecules, such as ethane (single bond), ethene (double bond) and ethyne (triple bond), and compare how the bond count changes the shape. Short regular practice sessions work better than a single long one, and a snack break helps: a simple idea is a fruit pot, like the ones in our guide on how to freeze fresh mango, ready in the freezer for study evenings.

Comparing ethane, ethene and ethyne

Putting C2H4 alongside its relatives makes the pattern clear. Ethane (C2H6) has a C-C single bond, ethene (C2H4) has a C=C double bond, and ethyne (C2H2) has a C≡C triple bond.

Molecule	Valence electrons	C-C bond	Hybridisation	Shape at carbon
Ethane, C2H6	14	Single	sp3	Tetrahedral
Ethene, C2H4	12	Double	sp2	Trigonal planar
Ethyne, C2H2	10	Triple	sp	Linear

As the number of bonds between the carbons rises, the C-C distance gets shorter and the bond gets stronger, though the double bond in ethene is less than twice as strong as a single bond because the pi component is weaker than the sigma bond. That is also why ethene is more reactive than ethane: the exposed pi electrons readily attack other molecules, which is the basis of addition reactions and the manufacture of polyethylene.

Why the Lewis structure matters

Drawing the structure is more than a classroom exercise. It tells you where electrons sit, predicts the shape, and helps explain reactivity, polarity and boiling point. Ethene is an important industrial chemical and is also a plant hormone that helps ripen fruit, which shows how a simple structure connects to real-world behaviour.

Frequently asked questions

Conclusion

The Lewis structure of C2H4 is easy once you follow the steps: count 12 electrons, join two carbons, add four hydrogens and share a second pair between the carbons to complete both octets. Check formal charges, and you have the correct H2C=CH2 diagram.