Basic Concepts of Chemical Bonding

Chemical Bonds

Chemical bonds are the forces that hold atoms together in compounds. There are three primary types of chemical bonds:

• Ionic Bonds: Electrostatic attraction between ions, typically formed between metals and nonmetals.

• Covalent Bonds: Sharing of electrons between atoms, usually between nonmetals.

• Metallic Bonds: Free electrons hold metal atoms together, resulting in properties such as conductivity and malleability.

Lewis Symbols and the Octet Rule

G. N. Lewis introduced a method to represent valence electrons using dots around the element symbol. The octet rule states that atoms tend to gain, lose, or share electrons to achieve eight valence electrons, similar to the configuration of noble gases.

Ionic Bonding

Ionic bonding occurs between metals and nonmetals (excluding group 8A elements) and involves the transfer of electrons. This process is highly exothermic.

• One element (metal) gives up electrons (low ionization energy).

• Another element (nonmetal) gains electrons (high electron affinity).

• Arrows in diagrams indicate electron transfer.

Properties of Ionic Substances

• Brittle

• High melting points

• Crystalline structure

• Cleave along smooth lines

Energetics of Ionic Bonding

The formation of ionic compounds involves several energy changes:

• Energy is required to convert elements to atoms (endothermic).

• Energy is required to create cations (endothermic).

• Energy is released to form anions (exothermic).

• Formation of solid releases a large amount of energy (exothermic).

The Born–Haber Cycle is used to calculate the overall energy change.

Lattice Energy

Lattice energy is a measure of the stability gained by arranging oppositely charged ions in an ionic solid. It is defined as the energy required to completely separate one mole of a solid ionic compound into its gaseous ions.

Trends in lattice energy:

• Increases with increasing charge on the ions

• Increases with decreasing size of the ions

Electron Configuration of Ions

• Main group metals lose electrons to achieve the configuration of the previous noble gas.

• Nonmetals gain electrons to achieve the configuration of the nearest noble gas.

• Transition metals lose valence electrons first, then d-electrons as needed for the ion charge.

Covalent Bonding

Covalent bonds involve the sharing of electrons between atoms, primarily nonmetals. Several electrostatic interactions occur:

• Attractions between electrons and nuclei

• Repulsions between electrons

• Repulsions between nuclei

For a bond to form, attractions must outweigh repulsions.

Lewis Structures

Lewis structures visually represent the sharing of electrons in covalent bonds. The goal is to give each atom the same number of electrons as the nearest noble gas.

• Lone pairs: Unshared pairs of electrons on one atom

• Bonding pairs: Shared electrons between two atoms, shown as two dots or a line

Number of Bonds for Nonmetals

The group number indicates the number of valence electrons. The number of bonds needed to complete the octet equals the number of electrons required to reach eight.

Multiple Bonds

• Single bond: One pair of electrons shared

• Double bond: Two pairs of electrons shared

• Triple bond: Three pairs of electrons shared

Polarity of Bonds and Electronegativity

Electrons in covalent bonds are not always shared equally. Bond polarity measures the equality of electron sharing:

• Nonpolar covalent bond: Electrons shared equally

• Polar covalent bond: One atom attracts electrons more strongly

Electronegativity is the ability of an atom to attract electrons in a molecule. It increases from left to right across a period and from bottom to top within a group.

The greater the difference in electronegativity, the more polar the bond. If the difference is greater than 2.0, the bond is often ionic.

Dipoles

A dipole forms when two equal but opposite charges are separated by a distance. The dipole moment is calculated as:

where is the dipole moment, is the charge, and is the distance. It is measured in debyes (D).

Comparing Ionic and Covalent Bonding

• Ionic: Complete electron transfer (metal + nonmetal)

• Covalent: Electron pair sharing (two nonmetals)

• Use electronegativity difference and physical properties to distinguish

Drawing Lewis Structures

Steps for drawing Lewis structures:

1. Sum valence electrons from all atoms, adjusting for charge

2. Write atom symbols, connect with single bonds

3. Complete octets around atoms bonded to the central atom

4. Place remaining electrons on the central atom

5. If needed, use multiple bonds to complete octets

6. Assign formal charges; the best structure minimizes formal charges and places negative charges on the most electronegative atoms

Resonance Structures

Resonance structures are used when a single Lewis structure cannot accurately depict a molecule. Electrons are delocalized and shared among multiple atoms. Examples include ozone and benzene.

Exceptions to the Octet Rule

Some ions and molecules do not follow the octet rule:

• Odd number of electrons (e.g., nitrogen oxide)

• Fewer than eight electrons (e.g., boron compounds)

• More than eight electrons (expanded octet, e.g., phosphate ion)

Strengths and Lengths of Single and Multiple Bonds

Bond enthalpy is the energy required to break a bond. Multiple bonds are stronger and shorter than single bonds. Bond enthalpy values are averages and may vary between compounds.

• As the number of bonds increases, bond length decreases

• Bond breaking is endothermic (positive enthalpy)

Summary Table: Types of Chemical Bonds

Summary Table: Exceptions to the Octet Rule