BackChemical Bonding: Lewis Structures, Molecular Shapes, and Bond Theories
Study Guide - Smart Notes
Tailored notes based on your materials, expanded with key definitions, examples, and context.
Chapter 5: Chemical Bonding I – Lewis Structures and Molecular Shapes
Bond Polarity and Electronegativity
Bond polarity arises from differences in electronegativity between bonded atoms, leading to unequal sharing of electrons.
Electronegativity: The ability of an atom to attract shared electrons in a chemical bond.
Trend in Electronegativity:
Increases across a period (left to right) on the periodic table.
Decreases down a group (top to bottom).
Fluorine is the most electronegative element.
Bond Polarity Determination:
If the difference in electronegativity () between two atoms is:
0: Nonpolar covalent bond
0 < < 2.0: Polar covalent bond
> 2.0: Ionic bond
Example: In HCl, Cl is more electronegative than H, so the bond is polar with a partial negative charge on Cl.
Lewis Structures
Lewis structures represent the arrangement of valence electrons among atoms in a molecule.
Steps to Draw Lewis Structures:
Count total valence electrons for all atoms.
Arrange atoms (central atom is usually least electronegative, except H).
Connect atoms with single bonds.
Distribute remaining electrons as lone pairs to complete octets (or duets for H).
Form double or triple bonds if necessary to satisfy octet rule.
Example: For CO2, carbon is central, each oxygen forms a double bond with carbon.
Resonance Structures
Some molecules cannot be represented by a single Lewis structure; instead, multiple resonance structures are drawn.
Definition: Resonance structures are two or more valid Lewis structures for the same molecule that differ only in the placement of electrons.
Example: Ozone (O3) has two resonance structures with the double bond in different positions.
Formal Charge
Formal charge helps determine the most stable Lewis structure.
Formula:
Lowest formal charges (closest to zero) are preferred.
Example: In the nitrate ion (NO3-), formal charges are distributed to minimize charge separation.
Exceptions to the Octet Rule
Some molecules have atoms with less or more than eight valence electrons.
Incomplete Octet: Elements like Be, B can have fewer than 8 electrons (e.g., BF3).
Expanded Octet: Elements in period 3 or higher (e.g., P, S, Cl) can have more than 8 electrons (e.g., SF6).
Odd-Electron Molecules: Molecules with an odd number of electrons (e.g., NO).
Bond Strength and Bond Length
The strength and length of a bond depend on the number of shared electron pairs.
Single Bond: Longest and weakest
Double Bond: Intermediate length and strength
Triple Bond: Shortest and strongest
Example: C≡C (triple bond) is shorter and stronger than C=C (double bond) or C–C (single bond).
VSEPR Theory: Shapes and Bond Angles
Valence Shell Electron Pair Repulsion (VSEPR) theory predicts molecular shapes based on electron pair repulsion.
Electron Domains: Regions of electron density (bonds and lone pairs) around a central atom.
Common Shapes and Bond Angles:
Electron Domains
Geometry
Bond Angle
2
Linear
180°
3
Trigonal planar
120°
4
Tetrahedral
109.5°
5
Trigonal bipyramidal
90°, 120°
6
Octahedral
90°
Lone Pairs: Lone pairs occupy more space and decrease bond angles.
Example: In NH3, the presence of a lone pair reduces the bond angle from 109.5° to about 107°.
Predicting Molecular Shape and Polarity
The shape and distribution of polar bonds determine the overall polarity of a molecule.
Nonpolar Molecule: Symmetrical shape with bond dipoles canceling (e.g., CO2).
Polar Molecule: Asymmetrical shape or uneven distribution of polar bonds (e.g., H2O).
Example: CH4 is nonpolar; NH3 is polar.
Chapter 6: Chemical Bonding II – Valence Bond and Molecular Orbital Theory
Hybridization of Atoms in Molecules
Hybridization describes the mixing of atomic orbitals to form new, equivalent hybrid orbitals for bonding.
Types of Hybridization:
sp: Linear geometry (2 electron domains)
sp2: Trigonal planar geometry (3 electron domains)
sp3: Tetrahedral geometry (4 electron domains)
sp3d: Trigonal bipyramidal geometry (5 electron domains)
sp3d2: Octahedral geometry (6 electron domains)
Example: In methane (CH4), carbon is sp3 hybridized.
Sigma (σ) and Pi (π) Bonds
Sigma and pi bonds are types of covalent bonds formed by different orbital overlaps.
Sigma (σ) Bond: Formed by head-on overlap of orbitals; all single bonds are sigma bonds.
Pi (π) Bond: Formed by side-to-side overlap of p orbitals; present in double and triple bonds (in addition to a sigma bond).
Example: In ethene (C2H4), the C=C bond consists of one sigma and one pi bond.
Molecular Orbital (MO) Theory for Homonuclear Diatomic Molecules/Ions
MO theory describes the combination of atomic orbitals to form molecular orbitals that are delocalized over the entire molecule.
Bonding and Antibonding Orbitals: Constructive overlap forms bonding orbitals (lower energy); destructive overlap forms antibonding orbitals (higher energy).
MO Diagrams: Show the relative energies and filling of molecular orbitals for diatomic molecules (e.g., O2, N2).
Example: O2 has two unpaired electrons in π* orbitals, explaining its paramagnetism.
Bond Order and Magnetic Behavior
Bond order indicates the strength and stability of a bond; magnetic behavior depends on unpaired electrons.
Bond Order Formula:
Bond Order Interpretation:
Higher bond order = stronger, shorter bond
Bond order of 0 = molecule is unstable
Magnetic Behavior:
Paramagnetic: Molecules with unpaired electrons (attracted to magnetic field)
Diamagnetic: Molecules with all electrons paired (repelled by magnetic field)
Example: N2 has a bond order of 3 and is diamagnetic; O2 has a bond order of 2 and is paramagnetic.
Additional info: For MO diagrams, refer to provided figures in your course materials for specific orbital energy ordering in different diatomic molecules.
