Chemical Bonding, Lewis Structures, and Molecular Geometry: Study Notes for General Chemistry

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Chemical Bonding and Molecular Shape

Introduction to Chemical Bonding

Chemical bonding is fundamental to understanding the structure and properties of molecules. Atoms bond to lower their potential energy, resulting in stable compounds. The nature of these bonds determines molecular shape, stability, and reactivity.

Chemical bonds form when the potential energy of bonded atoms is less than that of separate atoms.
Interactions to consider: nucleus-nucleus repulsion, electron-electron repulsion, and nucleus-electron attraction.

Bonding Theories

Bonding theories explain how and why atoms attach to form molecules, predict molecular shapes, and help understand chemical and physical properties.

Lewis theory: Focuses on valence electrons to explain bonding and predict properties such as stability, shape, and polarity.

Types of Chemical Bonds

Bond Classification

Bonds are classified based on the types of atoms involved and the nature of electron sharing or transfer.

Covalent bonds: Electrons are shared between atoms.
Polar covalent bonds: Electrons are shared unequally, resulting in partial charges.
Ionic bonds: Electrons are transferred, creating ions.

The continuum of bond types: pure covalent, polar covalent, and ionic

Bond Polarity and Electronegativity

Bond polarity arises from differences in electronegativity, the ability of an atom to attract electrons in a bond.

Electronegativity increases across a period and decreases down a group.
Fluorine is the most electronegative element; francium is the least.
The greater the electronegativity difference, the more polar the bond.

Periodic table showing trends in electronegativity Bond polarity in HF molecule Electron density distribution in a polar bond

Electronegativity Difference and Bond Type

The type of bond formed depends on the difference in electronegativity between the atoms:

0: Pure covalent (nonpolar)
0.1–0.4: Nonpolar covalent
0.5–1.9: Polar covalent
≥2.0: Ionic

Percent Ionic Character

Percent ionic character measures how much a bond resembles a purely ionic bond, based on dipole moment.

Percent ionic character vs. electronegativity difference

Lewis Structures and Resonance

Writing Lewis Structures

Lewis structures represent the distribution of valence electrons in molecules and ions, helping predict shape and reactivity.

Central atom is usually the least electronegative (except hydrogen, which is always terminal).
Count total valence electrons, including charge for ions.
Distribute electrons to give octets (or duets for hydrogen).
Form double or triple bonds if necessary to complete octets.

Lewis structure for CO2 showing double bonds Lewis structure for NH3 Calculation of total electrons for NH4+ ion Lewis structure for NH4+ ion Bracketed Lewis structure for NH4+ ion with charge

Resonance Structures

Resonance occurs when multiple valid Lewis structures can be drawn for a molecule, differing only in electron placement. The actual molecule is a resonance hybrid, intermediate between these forms.

Resonance stabilizes molecules by delocalizing electrons.
Resonance structures must have the same connectivity and number of electrons.

Hybridization analogy and resonance hybrid structure Calculation of total electrons for NO3- ion Resonance structures for NO3- ion Three resonance structures for NO3- ion and space-filling model

Formal Charge

Formal charge is an electron bookkeeping method to distinguish between alternative Lewis structures. It is calculated as:

Formal charge = (valence electrons) - (nonbonding electrons) - 0.5 × (bonding electrons)
Sum of formal charges in a neutral molecule must be zero; in an ion, it equals the ion's charge.
Structures with smaller formal charges and negative charges on more electronegative atoms are preferred.

Exceptions to the Octet Rule

Odd-Electron Species (Radicals)

Some molecules have an odd number of electrons and cannot satisfy the octet rule. These are called free radicals and are typically unstable and reactive.

Incomplete Octets

Certain elements, such as hydrogen and boron, often have fewer than eight electrons around them.

Expanded Octets

Atoms in the third period and below can have more than eight electrons when they are the central atom, using empty d orbitals.

Bond Energies and Bond Lengths

Bond Energies

Bond energy is the energy required to break one mole of a bond in the gaseous state. Stronger bonds require more energy to break.

More shared electrons = stronger bond
Shorter bonds = stronger bond

Bond Lengths

Bond length is the distance between the nuclei of bonded atoms. Triple bonds are shorter than double bonds, which are shorter than single bonds.

VSEPR Theory and Molecular Geometry

Valence Shell Electron Pair Repulsion (VSEPR) Theory

VSEPR theory predicts molecular shapes based on the repulsion between electron groups around a central atom. Electron groups include lone pairs, single, double, and triple bonds.

Electron groups arrange themselves to minimize repulsion.
Five basic electron group geometries: linear, trigonal planar, tetrahedral, trigonal bipyramidal, octahedral.

VSEPR theory: electron group repulsion

Electron Group Geometry Examples

Linear: Two electron groups, bond angle 180°.
Trigonal planar: Three electron groups, bond angle 120°.
Tetrahedral: Four electron groups, bond angle 109.5°.
Trigonal bipyramidal: Five electron groups, bond angles 90°, 120°, 180°.
Octahedral: Six electron groups, bond angle 90°.

Linear electron geometry Linear electron geometry example Trigonal planar electron geometry Tetrahedral electron geometry Trigonal bipyramidal electron geometry Octahedral electron geometry

Effect of Lone Pairs on Molecular Geometry

Lone pairs occupy more space than bonding pairs, causing bond angles to decrease from ideal values. This leads to shapes such as bent and trigonal pyramidal.

Trigonal pyramidal geometry Bent geometry from tetrahedral electron geometry

Derivatives of Trigonal Bipyramidal and Octahedral Geometries

Lone pairs in these geometries lead to shapes such as seesaw, T-shaped, linear, square pyramidal, and square planar.

Seesaw geometry T-shaped geometry Linear geometry from trigonal bipyramidal Square pyramidal geometry Square planar geometry

Summary Table: Electron and Molecular Geometries

Electron Groups	Electron Geometry	Molecular Geometry	Bond Angle
2	Linear	Linear	180°
3	Trigonal planar	Trigonal planar / Bent	120° / <120°
4	Tetrahedral	Tetrahedral / Trigonal pyramidal / Bent	109.5° / <109.5°
5	Trigonal bipyramidal	Trigonal bipyramidal / Seesaw / T-shaped / Linear	90°, 120°, 180°
6	Octahedral	Octahedral / Square pyramidal / Square planar	90°

Table of electron and molecular geometries

Molecular Polarity

Determining Molecular Polarity

A molecule is polar if it contains polar bonds and the bond dipoles do not cancel due to molecular geometry.

Draw the Lewis structure and determine geometry.
Identify polar bonds (electronegativity difference).
Superimpose vectors for each polar bond.
Sum vectors: if they cancel, molecule is nonpolar; if not, molecule is polar.

Polar bond in HCl Nonpolar CO2 molecule Polar H2O molecule

Vector Addition and Dipole Moments

Dipole moments are vector quantities. Their addition determines overall molecular polarity.

Identical polar bonds in opposite directions cancel (nonpolar).
Non-identical or angled polar bonds may not cancel (polar).

Dipole moments cancel in linear geometry Dipole moments do not cancel in bent geometry Dipole moments cancel in trigonal planar geometry Dipole moments cancel in tetrahedral geometry

Practice: Is a Molecule Polar?

Example: NH3 (ammonia) has three polar N–H bonds and a trigonal pyramidal geometry, resulting in a net dipole moment (polar molecule).

Lewis structure for NH3 Dipole moment vectors in NH3 Net dipole moment in NH3

Summary

Bonding theories and Lewis structures are essential for predicting molecular shape and properties.
VSEPR theory explains molecular geometry based on electron group repulsion.
Molecular polarity depends on both bond polarity and molecular geometry.