Chemical Bonding and Lewis Structures: General Chemistry Study Notes

Chemical Bonds

Introduction to Chemical Bonds

Chemical bonds are the forces of attraction that hold atoms together in compounds. The properties of compounds are typically different from those of their constituent elements. For example, hydrogen and oxygen are both gases, but they combine to form water, a liquid with distinct properties.

• Chemical bonds result in a more stable electron configuration for atoms, lowering their potential energy.

• Three main types of chemical bonds: ionic, covalent, and metallic.

Ionic Bonds

An ionic bond is formed by the transfer of valence electrons from a metal to a nonmetal, resulting in oppositely charged ions that are held together by electrostatic attraction.

• Metals lose electrons (low ionization energy) to form cations.

• Nonmetals gain electrons (high electron affinity) to form anions.

• Ionic compounds form a lattice structure, leading to high melting and boiling points.

• Examples: NaCl, MgO

Covalent Bonds

A covalent bond is formed when two atoms share valence electrons, typically occurring between nonmetals or between nonmetals and metalloids.

• Nonmetals have high ionization energy, so electrons are shared to achieve stability.

• Unequal sharing of electrons leads to polar covalent bonds; equal sharing leads to nonpolar covalent bonds.

• Examples: H2O, CO2, CH4

Electronegativity

Electronegativity is a measure of an atom's tendency to attract a bonding pair of electrons. It determines the type of bond formed between atoms.

• Bond type is determined by the difference in electronegativity:

Bond Type	Electronegativity Difference
Nonpolar covalent	< 0.5
Polar covalent	0.5 – 2.0
Ionic	> 2.0

• Examples: C–O (1.0), O–H (1.4)

Metallic Bonds

A metallic bond is formed by a "sea" of delocalized electrons shared among a lattice of metal cations. This bonding gives metals their characteristic properties.

• Metals have low ionization energy, allowing them to lose electrons easily.

• Valence electrons are shared by all metal ions, resulting in strong but flexible bonds.

• Properties: malleable, ductile, good electrical and thermal conductivity.

• Forms crystalline solids with continuous structure.

• Examples: Cu, Fe, Al

Summary Table: Bond Types

Bond Type	Example	Electron Behavior
Ionic	NaCl, MgO	Transferred from metal (cation) to nonmetal (anion)
Covalent	H2O, CO2	Shared between nonmetals or metalloids
Metallic	Cu, Fe	Delocalized, free-moving electrons

Lewis Bonding Theory

Introduction to Lewis Structures

Lewis Structures (Electron Dot Structures) are models used to represent valence electrons in atoms and molecules. They help in understanding molecular shape, polarity, and reactivity.

Lewis Structures for Atoms

• Represent the nucleus and core electrons with the element symbol.

• Valence electrons are shown as dots around the symbol.

• First two dots represent s orbital electrons; remaining dots represent p orbital electrons.

Lewis Structures for Ions

• Cations: Lost electrons; no dots shown; write the charge.

• Anions: Gained electrons; full octet (8 electrons); use brackets and write the charge.

Lewis Structures for Ionic Compounds

• Use Lewis symbols to show electron transfer from metal to nonmetal.

• Metal loses all valence electrons; nonmetal gains a full octet.

• Adjust the number of atoms to ensure even electron transfer.

Lewis Structures for Covalent Compounds

• Use Lewis symbols to show sharing of electrons between nonmetals.

• Bonding pairs: Shared electrons.

• Single bond: One pair (two electrons).

• Double bond: Two pairs (four electrons).

• Triple bond: Three pairs (six electrons).

• Lone pairs: Non-shared electrons on the central atom.

• Nonbonding pairs: Non-shared electrons on outside atoms.

Drawing Covalent Lewis Structures

1. Determine the total number of valence electrons for all atoms (adjust for ions).

2. Draw a skeletal structure, connecting atoms with single bonds.

3. Distribute electrons to terminal atoms to complete their octets.

4. Place remaining electrons on the central atom.

5. Rearrange to form multiple bonds if necessary to satisfy the octet rule.

Formal Charge

Definition and Calculation

Formal charge is the charge assigned to an atom in a molecule, assuming electrons in all chemical bonds are shared equally. It helps determine the most stable Lewis structure.

• Formula: $\text{Formal Charge} = \text{Valence Electrons} - \text{Nonbonding Electrons} - \frac{1}{2} \text{Bonding Electrons}$

• Structures with smaller formal charges are generally more stable.

• Best Lewis structures have formal charges close to zero.

• More electronegative element has the negative formal charge.

Octet Rule Exceptions

Types of Exceptions

• Incomplete octet: Molecules with less than 8 electrons (e.g., H (2 electrons), Be (4 electrons), B (6 electrons)).

• Expanded octet: Molecules with more than 8 electrons (hypervalent molecules), possible for atoms in the 3rd energy level and higher.

Resonance Structures

Definition and Drawing

Resonance structures occur when more than one valid Lewis structure exists for a molecule. The actual structure is a resonance hybrid, a blend of all possible structures.

• Delocalization of electrons stabilizes the molecule.

• Favorable structures have fewer and smaller formal charges.

• When drawing resonance structures:

  • Keep atom positions and connectivity the same.

  • Only move electron positions.

  • Total number of electrons and formal charges must remain constant.

  • Resonance often occurs with multiple bonds and/or lone pairs that allow electron delocalization.

  • Elements in the third period and beyond can form additional resonance structures due to expanded octets.

Knowledge Check & Review Questions

Sample Questions

• Match bond types to definitions (ionic, covalent, metallic).

• Draw Lewis structures for given molecules (e.g., CBr4, PF3, CO2).

• Calculate formal charges for atoms in a molecule.

• Draw resonance structures for polyatomic ions (e.g., CO32-, SO42-).

• Identify octet rule exceptions and describe metallic bonding.

Additional info: These notes are based on slides from OpenStax Chemistry: Atoms First 2e, covering foundational topics in chemical bonding and Lewis structures for General Chemistry students.