Chemical Bonding: Lewis Model, Molecular Geometry, and Electronegativity

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Chapter 10: Chemical Bonding

Bonding Theories: The Lewis Model

The Lewis model is a foundational theory in chemistry that predicts how atoms bond to form compounds and explains the shapes of molecules, which determine their physical and chemical properties. Developed by G.N. Lewis, this model uses dot structures to represent valence electrons and provides a simple method for predicting molecular stability and structure.

Bonding theories predict which combinations of atoms form compounds.
The Lewis model uses dots to represent valence electrons around element symbols.
Lewis structures help determine molecular stability and geometry.

Representing Valence Electrons with Dots

In the Lewis model, valence electrons of main-group elements are shown as dots around the element's symbol. The number of valence electrons equals the group number in the periodic table (except for helium).

Lewis structure: Dots represent valence electrons, placed around the element symbol.
Maximum of two dots per side; fill dots singly before pairing.
Atoms with eight valence electrons (octet) are particularly stable.
Helium is an exception, stable with two electrons (duet).

Electron configuration of oxygen showing 6 valence electrons Lewis structure of oxygen with 6 dots Lewis structures for Period 2 elements Lewis structure of phosphorus with 5 dots

Lewis Model of Chemical Bonding

Chemical bonds involve the sharing or transfer of electrons to achieve stable electron configurations. The two main types of bonds are ionic and covalent.

Ionic bond: Electrons are transferred from metal to nonmetal.
Covalent bond: Electrons are shared between nonmetals.
Octet rule: Atoms transfer or share electrons to obtain eight electrons in their outer shell (except H and He, which require two).

Lewis Structures for Ionic Compounds: Electrons Transferred

When metals bond with nonmetals, electrons are transferred, forming cations and anions. The attraction between these ions results in ionic compounds.

Electron dots are moved from metal to nonmetal in Lewis structures.
Anions are shown in brackets with their charge.

Lewis structures for potassium and chlorine Electron transfer from potassium to chlorine Lewis structure of KCl with charges

Lewis Structures for Covalent Compounds: Electrons Shared

Nonmetals form molecular compounds by sharing electrons, resulting in covalent bonds. Lewis structures show how atoms share electrons to achieve octets or duets.

Bonding pair electrons: Shared between two atoms.
Lone pair electrons: Not shared, belong to one atom.
Bonding pairs are often represented by dashes.

Lewis structures for hydrogen and oxygen Lewis structure for water (H2O) Shared electrons in water molecule Bonding pair and lone pair in water Bonding pairs represented by dashes in water

Double and Triple Bonds in Covalent Compounds

Some molecules require double or triple bonds to satisfy the octet rule. Double bonds involve four electrons, while triple bonds involve six electrons between two atoms.

Double bond: Two pairs of shared electrons (e.g., O2).
Triple bond: Three pairs of shared electrons (e.g., N2).
Double and triple bonds are shorter and stronger than single bonds.

Lewis structure for O2 before double bond Formation of double bond in O2 Octet achieved in O2 with double bond Formation of triple bond in N2

Steps to Write Lewis Structures for Covalent Compounds

Follow these steps to construct Lewis structures for covalent compounds:

Write the correct skeletal structure for the molecule.
Calculate the total number of valence electrons.
Distribute electrons to give octets (or duets for hydrogen).
Form double or triple bonds if necessary to satisfy octets.

Bonding electrons around central atom in CCl4 Lone pairs on terminal atoms in CCl4 Bonding electrons in CO2 Lone pairs moved to form double bonds in CO2

Exceptions to the Octet Rule

Some molecules do not follow the octet rule. Examples include molecules with odd numbers of electrons (free radicals), molecules with fewer than eight electrons (e.g., boron compounds), and molecules with expanded octets (e.g., SF6).

Lewis structures for BF3 and BH3 (boron compounds) Lewis structures for SF6 and PCl5 (expanded octets)

Resonance: Equivalent Lewis Structures

Some molecules, such as SO2, can be represented by more than one valid Lewis structure. These are called resonance structures, and the true structure is an average of all resonance forms.

Ball-and-stick model for SO2

Predicting the Shapes of Molecules: VSEPR Theory

The Lewis model, combined with Valence Shell Electron Pair Repulsion (VSEPR) theory, predicts molecular shapes based on the repulsion between electron groups around the central atom.

Electron geometry: Arrangement of electron groups.
Molecular geometry: Arrangement of atoms.
Bond angles depend on the number of electron groups.

Linear geometry of CO2 Lewis structure for H2CO (formaldehyde) Tetrahedral geometry of CH4 (ball-and-stick model) Space-filling model of CH4

Predicting Shapes with Lone Pairs: Molecular Geometry

Lone pairs on the central atom affect molecular geometry. For example, NH3 has a trigonal pyramidal shape, and H2O has a bent shape.

Lewis structure for NH3 Ball-and-stick model for NH3 Trigonal pyramidal structure for NH3 Lewis structure for H2O Ball-and-stick model for H2O Bent structure for H2O

Table: Electron and Molecular Geometries

This table summarizes the relationship between electron groups, bonding groups, lone pairs, electron geometry, bond angles, molecular geometry, and examples.

Electron Groups	Bonding Groups	Lone Pairs	Electron Geometry	Angle	Molecular Geometry	Example
2	2	0	Linear	180°	Linear	CO2
3	3	0	Trigonal planar	120°	Trigonal planar	H2CO
3	2	1	Trigonal planar	109.5°	Bent	SO2
4	4	0	Tetrahedral	109.5°	Tetrahedral	CH4
4	3	1	Tetrahedral	109.5°	Trigonal pyramidal	NH3
4	2	2	Tetrahedral	109.5°	Bent	H2O

Bond angles for molecular geometries

Representing Molecular Geometries on Paper

Chemists use different notations to represent three-dimensional molecular structures on two-dimensional paper, such as straight lines, hashed lines, and wedges.

Bond notation for 3D structures Major molecular geometries

Electronegativity and Polarity

Electronegativity is the ability of an atom to attract electrons in a covalent bond. Differences in electronegativity lead to polar covalent bonds and dipole moments, which affect molecular interactions and properties.

Electronegativity increases across a period and decreases down a group.
Fluorine is the most electronegative element (4.0).
Dipole moment: Separation of charge in a bond due to unequal sharing of electrons.
Polar covalent bonds: Electrons are shared unequally.
Pure covalent bonds: Electrons are shared equally (no dipole moment).
Ionic bonds: Electrons are transferred (large electronegativity difference).

Oil and water separation due to polarity Lewis structure for water showing polar bonds Dipole moment in O-H bond Electronegativity scale Pure covalent bond in Cl2 Ionic bond in NaCl Polar covalent bond in HCl Effect of electronegativity difference on bond type

Polar Bonds and Polar Molecules

The presence of polar bonds does not always result in a polar molecule. The overall molecular polarity depends on the geometry and whether dipole moments cancel.

Diatomic molecules with polar bonds are always polar.
In molecules with more than two atoms, polar bonds may cancel (nonpolar molecule) or add (polar molecule).

Dipole moments in CO2 Dipole moments in H2O Dipole moments cancel in linear and trigonal planar molecules Dipole moments in trigonal pyramidal and bent molecules

Everyday Chemistry: How Soap Works

Soap molecules have both polar and nonpolar ends, allowing them to interact with both water (polar) and grease/oil (nonpolar), enabling the removal of grease from surfaces.

Review

The Lewis model explains chemical bonding through electron transfer (ionic) or sharing (covalent).
Molecular shapes are predicted using VSEPR theory.
Electronegativity determines bond polarity and molecular interactions.
Polar molecules interact strongly with other polar molecules, while nonpolar molecules are excluded.

Highlight: Free Radicals

Free radicals are molecules with an odd number of valence electrons, making them chemically unstable and highly reactive. They are implicated in aging and disease due to their ability to damage cellular molecules.