Lewis structures – also called Lewis dot formulas, Lewis dot structures, electron dot structures, or Lewis electron dot structures – are diagrams that show the bonding between atoms of a molecule, as well as the lone pairs of electrons that may exist in the molecule.
CO2
- Lewis Structure: O=C=O with each bond being a double bond.
- Molecular Geometry: Linear.
- Hybridization: Carbon is sp hybridized, which allows for the formation of two σ bonds and two π bonds with the oxygen atoms.
CO2 Lewis Structure
Steps to Draw
Total Valence Electrons
- Carbon has 4 valence electrons.
- Each Oxygen has 6 valence electrons, and since there are two, this gives us 12.
- Total = 4 (C) + 12 (O) = 16 valence electrons.
Central Atom
- Carbon is the central atom since it is less electronegative than oxygen (and carbon can form more bonds, making it ideal as the central atom).
Bonding
- Place carbon in the center with oxygen atoms on either side.
- Connect each oxygen to carbon with two lines (each line representing a pair of shared electrons for a double bond), since oxygen prefers to have a double bond to complete its octet.
Electron Distribution
- Each double bond uses up 4 electrons (2 from each bond).
- This accounts for 8 electrons in bonds.
- The remaining 8 electrons are placed as lone pairs on the oxygen atoms to complete their octets (each oxygen needs 4 more electrons, or two pairs, to reach 8).
Check Octets
- Carbon shares 4 electrons with each oxygen, giving it an octet through the double bonds.
- Each oxygen has 2 lone pairs (4 electrons) and shares 4 electrons with carbon, also completing an octet.
Formal Charges
- Carbon: 4 (valence e-) – 4 (from bonds) = 0
- Oxygen: 6 (valence e-) – 4 (lone pairs) – 4 (from bonds) = 0
Molecular Geometry of CO2
The molecular geometry of CO2, or carbon dioxide, is linear. When you visualize CO2, think of carbon in the middle with an oxygen atom at each end, all in a straight line.
Molecular Geometry Determination:
VSEPR Theory: The Valence Shell Electron Pair Repulsion (VSEPR) theory is used to predict the shape of molecules based on the electron pairs around the central atom.
Electron Domains:
- Carbon, as the central atom in CO2, forms two double bonds with oxygen atoms.
- Double bonds count as one electron domain for the purpose of determining geometry.
Electron Domain Geometry:
- With two electron domains (the double bonds), the most stable arrangement to minimize repulsion is linear.
Bond Angles:
- The bond angle in a linear molecule is 180°.
Absence of Lone Pairs:
- Carbon has no lone pairs of electrons in CO2 since all its valence electrons are involved in bonding with oxygen. This lack of lone pairs means there’s no additional repulsion to alter the linear shape.
Implications of Linear Geometry
- Symmetry: The linear structure makes CO2 a symmetrical molecule which contributes to it being nonpolar, despite having polar bonds. This is because the dipole moments of the C=O bonds are equal and opposite, cancelling each other out.
- Physical Properties: The linear shape affects the molecule’s behavior in terms of its interactions with other molecules, its phase at different temperatures, and its ability to dissolve in various solvents.
Hybridization
This hybridization allows carbon to form the necessary bonds to create CO2 with its characteristic linear structure and double bonds. The sp hybridization explains the linear molecular geometry of CO2, where the nuclei of the atoms are in a straight line, and it also accounts for the molecule’s ability to form strong, stable double bonds with oxygen.
The hybridization of carbon in CO2 is an excellent example of sp hybridization. Here’s how it works:
Steps to Determine Hybridization:
Electronic Configuration:
- Carbon’s ground state electron configuration is 1s22s22p2.
Excitation:
- One electron from the 2s orbital is excited to the 2p orbital, resulting in four unpaired electrons: 1s22s12p3.
Hybridization:
- The 2s orbital and one of the 2p orbitals mix to form two sp hybrid orbitals. This leaves carbon with:
- Two sp hybrid orbitals, which are linear in shape.
- Two unhybridized 2p orbitals, which remain at right angles to each other and to the sp orbitals.
Bond Formation:
- Each sp hybrid orbital overlaps with a p orbital from an oxygen atom to form a sigma (σ) bond.
- The remaining two p orbitals on carbon each form a pi (π) bond with the p orbitals of the oxygen atoms. Since each oxygen atom also has one p orbital available for π bonding, this results in two π bonds, making the CO2 bonds double bonds.
Visual Representation
- Sp Hybrid Orbitals: They lie along the same axis, giving the molecule its linear shape, with an angle of 180° between them.
- Pi Bonds: These are formed above and below, as well as in front and behind the plane of the sigma bonds due to the lateral overlap of p orbitals.
Summary
Hybridization Type: sp
- 2 sigma (σ) bonds: Formed by the overlap of sp hybrid orbitals of carbon with p orbitals of oxygen.
- 2 pi (π) bonds: Formed by the overlap of unhybridized p orbitals of carbon with those of oxygen.