Chemical Bonding: quick revision, 18 bites

Why Atoms Bond: The Octet Rule

Atoms crave eight valence electrons.

Atoms combine to achieve a **noble gas configuration**—typically 8 valence electrons (or 2 for hydrogen). This drive is the **octet rule**, the foundation of chemical bonding. It explains why sodium loses one electron and chlorine gains one to form NaCl, and why carbon shares four electrons in CH₄.

Takeaway. The octet rule is the starting point for predicting how atoms bond.

Ionic vs Covalent Bonds

Transfer or share? That is the question.

Takeaway. Ionic bonds involve electron transfer; covalent bonds involve sharing.

Lattice Enthalpy

Energy released when ions form a crystal.

Takeaway. Lattice enthalpy measures the strength of an ionic bond.

Lewis Structures: Representing Valence Electrons

Dots and lines show bonding.

A **Lewis structure** uses dots for valence electrons and lines for shared pairs. For example, in water, oxygen shares one pair with each hydrogen, leaving two lone pairs on oxygen. Drawing Lewis structures helps predict molecular shape and reactivity.

Takeaway. Lewis structures are a visual shorthand for electron distribution.

Formal Charge Calculation

Which resonance structure is best?

Takeaway. Formal charge helps choose the most plausible Lewis structure.

VSEPR Theory: Predicting Molecular Shape

Electron pairs repel, so shapes form.

The **Valence Shell Electron Pair Repulsion (VSEPR)** theory states that electron pairs around a central atom repel each other and arrange to minimize repulsion. This determines the molecular geometry. For example, with four bonding pairs and no lone pairs, the shape is **tetrahedral** (e.g., CH₄).

Takeaway. VSEPR theory uses electron-pair repulsion to predict molecular shapes.

Common Molecular Shapes

From linear to seesaw.

Takeaway. Memorize these common shapes to quickly identify molecular geometry.

Valence Bond Theory: Orbital Overlap

Bonds form when orbitals overlap.

The **Valence Bond (VB) theory** explains that a covalent bond forms when atomic orbitals of two atoms overlap, and the electrons are shared in the region of overlap. The greater the overlap, the stronger the bond. For example, the s-orbital of hydrogen overlaps with the p-orbital of chlorine to form the H–Cl bond.

Takeaway. VB theory links bond formation to atomic orbital overlap.

Hybridization: Mixing Orbitals

Orbitals mix to form new shapes.

**Hybridization** is the mixing of atomic orbitals to form new, equivalent hybrid orbitals that point in specific directions. This explains molecular geometry. For example, carbon in methane undergoes **sp³ hybridization**, forming four equivalent orbitals pointing to the corners of a tetrahedron.

Takeaway. Hybridization explains the directional nature of covalent bonds.

Hybridization and Geometry

Quick reference for common cases.

Takeaway. Hybridization type directly determines molecular geometry.

Molecular Orbital Theory: Bonding and Antibonding

Electrons occupy molecular orbitals.

**Molecular Orbital (MO) theory** describes electrons as occupying orbitals that span the entire molecule. When atomic orbitals combine, they form **bonding** (lower energy) and **antibonding** (higher energy) molecular orbitals. The number of electrons in bonding vs antibonding orbitals determines bond order and stabi…

Takeaway. MO theory explains bond order and magnetic behavior.

Bond Parameters: Length, Angle, and Energy

Bonds have measurable characteristics.

**Bond length** is the distance between bonded nuclei; **bond angle** is the angle between two bonds; **bond enthalpy** is the energy required to break one mole of bonds. These parameters influence molecular properties. For example, shorter bonds are generally stronger.

Takeaway. Bond parameters quantify the strength and geometry of bonds.

Polarity of Bonds and Molecules

Unequal sharing creates poles.

When atoms with different **electronegativities** bond, the electron pair is shared unequally, creating a **polar bond**. The molecule's overall polarity depends on the vector sum of bond dipoles. For example, CO₂ has polar C=O bonds, but the linear shape cancels the dipoles, making the molecule nonpolar.

Takeaway. Molecular polarity depends on both bond polarity and geometry.

Confusing Geometry with Electron-Pair Geometry

Lone pairs change the shape.

A common error is to describe the shape of NH₃ as tetrahedral. While the electron-pair geometry is tetrahedral (four electron pairs), the **molecular geometry** (considering only atoms) is **trigonal pyramidal**. Similarly, H₂O is bent, not tetrahedral.

Takeaway. Lone pairs affect molecular shape, not just electron-pair arrangement.

Predicting Geometry: A Step-by-Step Approach

A systematic method for exam success.

Takeaway. Follow these steps to reliably predict molecular shapes.

Remembering Hybridization

A simple mnemonic.

Use the phrase **“SP, SP², SP³”** to remember the order: **S**ingle bonds → **P**lanar? No—**S**ingle bonds → **sp³**, **D**ouble bonds → **sp²**, **T**riple bonds → **sp**. Think: “Single, Double, Triple” → “sp³, sp², sp”.

Takeaway. Associate bond order with hybridization type.

Water's Unique Properties

Bonding explains why ice floats.

The bent shape and polarity of water lead to **hydrogen bonding**, which gives water its high boiling point, surface tension, and the fact that ice is less dense than liquid water. This is a direct consequence of the molecular structure covered in this chapter.

Takeaway. Molecular structure dictates macroscopic properties.

Why Salt Dissolves in Water

Ionic bonds break, hydration occurs.

When NaCl dissolves in water, the polar water molecules surround and stabilize the Na⁺ and Cl⁻ ions, overcoming the ionic bond's lattice enthalpy. This is why salt dissolves easily—a real-world application of ionic bonding and polarity.

Takeaway. Polarity and lattice enthalpy explain solubility.