Chemical Bonding and Molecular Structure: Study Notes part 1

Chemical Bonding and Molecular Structure

Fundamentals of Bonding

Chemical bonding is the interaction between atoms that leads to the formation of molecules and compounds. The nature of bonding determines the properties and behavior of substances.

• Bond Energy: The energy required to break a bond between two atoms in a molecule.

• Bond Length: The equilibrium distance between the nuclei of two bonded atoms.

• Electronegativity: A measure of an atom's ability to attract electrons in a chemical bond. Electronegativity increases across a period and decreases down a group.

• Electron Affinity: The energy change when an atom gains an electron.

Example: The bond in H2 is formed by the sharing of two electrons between two hydrogen atoms.

Ionic Bonding

Ionic bonding occurs when electrons are transferred from one atom to another, resulting in the formation of ions. The electrostatic attraction between oppositely charged ions forms an ionic compound.

• NaCl Structure: Sodium chloride consists of a lattice of Na+ and Cl- ions.

• Lattice Energy: The energy released when ions come together to form a solid lattice. It depends on the charges and sizes of the ions.

Compound	Lattice Energy (kJ/mol)
NaCl	786
MgO	3923
CaO	3414

Lewis Structures

Lewis structures are diagrams that show the bonding between atoms and the lone pairs of electrons in a molecule. They help predict molecular geometry and reactivity.

• Steps to Draw Lewis Structures:

  1. Count the total number of valence electrons.

  2. Arrange atoms and connect with single bonds.

  3. Place lone pairs on outer atoms.

  4. Assign remaining electrons to the central atom.

  5. Minimize formal charges.

• Formal Charge Formula: $\text{Formal charge} = \text{valence electrons} - \text{nonbonding electrons} - \frac{1}{2} \text{bonding electrons}$

Example: The Lewis structure of CO2 is O=C=O, with no formal charges on any atom.

Resonance Structures

Resonance structures are alternative Lewis structures for a molecule or ion that differ only in the arrangement of electrons. The true structure is a hybrid of all resonance forms.

• Example: The phosphate ion, PO43-, has multiple resonance structures with delocalized electrons.

Valence-Shell Electron-Pair Repulsion (VSEPR) Theory

VSEPR theory predicts the shapes of molecules based on the repulsion between electron pairs around a central atom. The geometry depends on the number of bonding and lone pairs.

Electron Pairs	Geometry	Example
2	Linear	CO2
3	Trigonal planar	BF3
4	Tetrahedral	CH4
5	Trigonal bipyramidal	PCl5
6	Octahedral	SF6

Properties of Covalent Bonds

Covalent bonds involve the sharing of electrons between atoms. Their properties include bond polarity, dipole moments, bond length, and bond energy.

• Dipole Moment: A measure of the separation of positive and negative charges in a molecule. $\mu = q \times r$

• Bond Length: Shorter bonds are generally stronger; bond length depends on atomic radii and bond order.

• Bond Energy: The energy required to break a bond; higher bond order and shorter bond length usually mean higher bond energy.

Bond Type	Bond Length (pm)	Bond Energy (kJ/mol)
H-H	74	436
C-H	109	413
C=C	134	614
C≡C	120	839

Summary of Molecular Shapes

The shape of a molecule is determined by the number of electron pairs around the central atom, including both bonding and lone pairs. VSEPR theory provides a systematic way to predict molecular geometry.

Sets of Electron Pairs	Geometry	Bond Angle	Example
2	Linear	180°	CO2
3	Trigonal planar	120°	BF3
4	Tetrahedral	109.5°	CH4
5	Trigonal bipyramidal	90°, 120°	PCl5
6	Octahedral	90°	SF6

Additional info: These notes cover the main concepts from Chapter 6: Chemical Bonding and Molecular Structure, including ionic and covalent bonding, Lewis structures, resonance, VSEPR theory, and properties of bonds. All tables and formulas are reconstructed from the provided textbook images.