Ethylene is a hydrocarbon which has the formula C₂H₄ or H₂C=CH₂. It is a colourless, flammable gas with a faint “sweet and musky” odour when pure. It is the simplest alkene. Ethylene is widely used in the chemical industry, and its worldwide production exceeds that of any other organic compound
Lewis structure of C_2H_4, also known as ethene, is H_2_C=CH_2. This structure involves a double bond between the two carbon atoms, and each carbon atom is also bonded to two hydrogen atoms.
Here’s a detailed explanation of the Lewis structure for C₂H₄:
Determine the Total Number of Valence Electrons:
- Carbon (C): Each carbon atom has 4 valence electrons. Since there are two carbon atoms, this gives us 4×2=8 electrons.
- Hydrogen (H): Each hydrogen atom has 1 valence electron. There are four hydrogen atoms, so 1×4=4 electrons.
- Total Valence Electrons: 8 (from C)+4 (from H)=12 electrons
Arrange the Atoms:
- In ethene, the two carbon atoms are bonded to each other, and each carbon is also bonded to two hydrogen atoms. The molecule has a planar geometry with the hydrogen atoms spread out to minimize repulsion.
Form Bonds:
- Each bond represents a pair of electrons.
- Single Bonds: Start by forming single bonds between the atoms:
- C-C (one bond = 2 electrons)
- C-H (four bonds = 4 x 2 = 8 electrons)
- This uses up 2+8=10 electrons.
Place Remaining Electrons:
- We have 12−10=2 electrons left. These electrons will go into forming a double bond between the two carbon atoms because carbon needs to fulfill the octet rule (8 electrons around each atom).
Draw the Structure
- Here’s how the Lewis structure looks:
- Each line represents a bond (2 electrons). The double line between the carbons represents a double bond (4 electrons, 2 from each carbon).
Check the Octet Rule:
- Each carbon atom now has:
- 2 electrons from the single bonds with hydrogen (2 H atoms x 1 bond each)
- 4 electrons from the double bond with the other carbon.
- This gives each carbon atom 8 electrons in total (2 from H + 6 from C=C), satisfying the octet rule.
Formal Charges:
- In this structure, formal charges are ideally zero for all atoms if we consider each bond equally shared (which isn’t the case in reality due to differences in electronegativity, but for Lewis structures, this simplifies things):
Additional Notes:
- Geometry: Ethene is planar with bond angles around 120 degrees due to the sp² hybridization of carbon atoms.
- Pi Bond: The double bond contains one sigma bond and one pi bond. The pi electrons are above and below the plane, which makes ethene less stable than molecules with only single bonds, leading to its reactivity, particularly in addition reactions.
This structure explains why ethene can act as a monomer in polymerization reactions, as the pi bond can break to form new bonds with other monomers.