Bonding

Covalent Bonding - The Basics

Reference: Higher Level Chemistry,

Pages 176 - 177, 182 - 185 and 193 - 196

Essential Question

What determines the covalent nature and properties of a substance?

Covalent Bonding - Definitions

A covalent bond is defined as the simultaneous electrostatic attraction between a pair of electrons and positively charged nuclei.
Covalent bonds usually form between two non-metals.
In a covalent bond, electron pairs are shared to give each atom a stable octet of electrons.

Types of Covalent Bonds

1. Single Covalent Bonds

Example: Formation of the F-F bond in fluorine gas (F2) involves the sharing of one pair of electrons.

2. Double Covalent Bonds

Example: Formation of the O-O bond in oxygen gas (O2) involves the sharing of two pairs of electrons.

3. Triple Covalent Bonds

Example: Formation of the N-N bond in nitrogen gas (N2) involves the sharing of three pairs of electrons.

Comparing Single, Double, and Triple Covalent Bonds

Bond Type | Bond Length (nm) | Bond Enthalpy (kJ mol^{-1}) - single | 0.154 | 347 - double | 0.134 | 612 - triple | 0.120 | 838
Observations: - Bond strength increases with the number of shared electron pairs. - Bond length decreases as bonds become stronger.
Definition: Bond enthalpy is the energy required to break one mole of a given type of bond.

Types of Covalent Bonds - Electronegativity

Electronegativity (EN)

Definition: Electronegativity is a property of an atom relating to its ability to attract electrons.
Measured using the Pauling Scale: - 0: no or little tendency to attract electrons
- 4: strong tendency to attract electrons
Two Types of Covalent Bonds: Defined by the attraction of shared electrons.

Electronegativity Trends

Metals generally have lower EN, while non-metals have higher EN.
Group 18 elements exhibit no EN.
Notable Electronegativity Values:
- F = 4.0 - Br = 3.0 - Li = 1.0 - K = 0.8

Types of Covalent Bonds - Polar and Non-Polar

Non-Polar Covalent Bonds

Involve equal sharing of electrons between atoms of the same element (same EN).
No dipole moments are formed, as electron distribution is equal.

Polar Covalent Bonds

Involve unequal sharing of electrons where electrons are more attracted to one atom due to higher EN, creating dipoles (regions of partial charge).
Example: - Hydrogen (EN = 2.2) and Bromine (EN = 3.0) show unequal sharing resulting in a dipole.

Characteristics of Polar Covalent Bonds

Negative dipoles form around the atom with the higher EN.
Positive dipoles form around the atom with the lower EN.

Bond Dipoles and Strength

The strength of a polar covalent bond is related to the difference in electronegativity (∆EN) of the bonded atoms.
Examples of Polarities: - Most polar bond: ∆EN = 1.9 - Least polar bond: ∆EN = 0.0
The bond dipole's length represents strength.

The Bonding Continuum

Characteristics of bonds range from non-polar covalent to ionic.
Bond strength and character can be depicted graphically in a triangular bonding diagram.

Predictions of Bond Types Using Electronegativity

Use electronegativity differences to predict bond types for given elements: - Mg + O, H + Cl, C + S, - Cu + S, I + Br, N + N.

Example Calculations

Mg + O: ∆EN = 3.4-1.2=2.2 (ionic bond)
H + Cl: ∆EN = 3.2-2.2=1.0 (polar covalent bond)
C + S: ∆EN = 2.6-2.5=0.1 (non-polar covalent bond).

Lewis Structures for Covalent Compounds

Reference

Higher Level Chemistry, pp.179 - 182, 186 - 187, 227 - 236

Essential Objective

Goals: 1. Account for all valence electrons in the molecule. 2. Show distribution of valence electrons to achieve a complete valence shell (octet where applicable).

Steps for Drawing Lewis Structures

Count Total Valence Electrons: Adjust for any ions.
Draw Skeleton Structure: Identify central atom surrounded by outer atoms.
Draw Single Bonds: Place a pair of electrons between the central atom and each outer atom.
Complete Octets: Add remaining electrons to achieve stable octets.
Form Double or Triple Bonds (if necessary): If octets cannot be completed add double/triple bonds.

Examples

Example 1: PF3

Valence e- = 5 + 3(7) = 26.
Skeleton: F - P - F - F
Bonds: F:P:F
Remaining pairs to complete octets.
Complete illustration shows all achieved octets.

Example 2: CH2O

Valence e- = 4 + 2(1) + 6 = 12.
Skeleton: O - C with H's.
Add bonds and complete octet by forming double bond if necessary.

Coordinate (Dative) Covalent Bonds

In a coordinate bond, both shared electrons come from one atom.
Example: NH4+ is formed by the dative bond with the NH3 molecule and H+ cation.
Notable structure includes the arrow combining the molecule to the positive ion.

Exceptions to the Octet Rule

Some stable structures do not follow typical octet rules.
Elements like Be and B form stable molecules with fewer than 8 electrons: - Example BH3: Has 6 valence electrons. - Example BeCl2: Forms stable with 4 valence electrons.

Expanded Octets

Elements in Group 3 (like P) can have more than 8 electrons: - Example PCl5: Can hold 10 electrons.

Resonance Structures

Definition: Three or more Lewis structures that describe electron delocalization.
Example: Ozone structure.

Properties of Compounds with Delocalized Electrons

Characteristics include stable bonds, consistent bonding angle, and change in bond length.

Covalent Network Structures

Reference

Higher Level Chemistry, pp.198 - 205

Comparison of Structures

Giant Covalent Lattice (like diamond): Each atom is bonded to others via strong covalent bonds and arranged in a regular lattice.
Molecular Structures (like water): Form due to weaker intermolecular forces.

Carbon Allotropes

Diamond: Hardest known substance, does not conduct electricity.
Graphite: Brittle, conducts electricity, consists of sheets of carbon atoms.
Graphene: Flexible and conductive, significantly stronger than steel.

Hybridization in Covalent Bonds

Reference

Higher Level Chemistry, pp.239 - 245

Overview

Covalent bonds are formed by overlapping atomic orbitals leading to the hybridization of different types of orbitals:

sp3: 4 EDs, tetrahedral (109.5º). - Occurs in CH4, K2H2.
sp2: 3 EDs, trigonal planar (120º). - Example: C2H4, O3.
sp: 2 EDs, linear (180º). - Example: C2H2, HCN.

Sigma and Pi Bonds

Overview

Sigma Bonds (σ): Formed by head-on overlap of orbitals, every single bond contains one sigma bond.
Pi Bonds (π): Formed by side-on overlap of p orbitals, found in double and triple bonds.

Summary of Bonding

Covalent bonds can exhibit different types based on the overlap and arrangement of orbitals leading to sigma and pi formation.

Metallic Bonding and Alloys

Reference

Higher Level Chemistry, pp.255 - 262 and 271 - 273

Overview of Metallic Bonds

Metallic bonding involves a sea of delocalized electrons, contributing to properties such as:
High thermal and electrical conductivity.
Ductility and malleability.
Lustrous appearance.

Ionic Bonding and Structure

Reference

Higher Level Chemistry, pp.149 - 168

Overview of Ionic Bonds

Ionic bonds form through the transfer of electrons from metals to non-metals, generating cations and anions.

Characteristics of Ionic Compounds

Form crystal lattice structures, highly soluble in water, have high melting and boiling points.
Exhibit strong electrostatic forces between charged ions.

Summary Table of Properties

Property | Ionic Compounds | Metallic Compounds | Molecular Compounds - Brittleness: High | High | Low - Conductivity: Molten/Aqueous | Yes | Yes | No - Solubility: Mostly polar solvents | Yes | No | Depends on polarity - Volatility: Low | Low | Higher

Lattice Enthalpy

Definition

Lattice enthalpy represents the energy change associated with the formation or breaking of an ionic lattice.
Lattice enthalpy provides insights into stability and characteristics of ionic compounds.

Conclusion

Key concepts of bonding, properties, and structural characteristics across covalent, ionic, and metallic bonds were discussed comprehensively. The interactions and types of bonding lead to diverse properties exhibited in various chemical species.