Structure and bonding Notes

Chemical Bonding: How Electrons Hold Substances Together

At the heart of chemistry is a simple idea:

Chemical bonds form when atoms transfer or share electrons to achieve more stable electron configurations.

There are three main types of bonding you need to understand at this level:

1. Ionic bonding – electrons are transferred
2. Covalent bonding – electrons are shared
3. Metallic bonding – electrons are delocalised (a “sea of electrons”)

These bonding types, plus the way particles are arranged in space, explain why substances have such different properties (like melting point, conductivity, hardness, solubility, etc.).

Ionic Bonding – Electron Transfer and Giant Lattices

How Ionic Bonds Form

1. The metal loses one or more electrons → forms a positive ion (cation).
2. The non-metal gains those electrons → forms a negative ion (anion).
3. Oppositely charged ions are held together by strong electrostatic attraction.

Structure: Giant Ionic Lattice

1. Ionic compounds (like NaCl, MgO) don’t form single “molecules”.
2. Instead they form a giant ionic lattice:
   1. Every Na⁺ is surrounded by six Cl⁻.
   2. Every Cl⁻ is surrounded by six Na⁺.
   3. Positive and negative ions alternate in a regular 3D structure.
3. This maximizes attraction and minimizes repulsion.

How Structure Explains Ionic Properties

1. Very high melting and boiling points
   1. Strong electrostatic forces in the lattice require a lot of energy to break.
2. Electrical conductivity
   1. Solid: does not conduct – ions are fixed in place.
   2. Molten (liquid) or aqueous (dissolved in water): does conduct – ions are free to move and carry charge.
3. Brittle
   1. If layers are forced to slide, ions of the same charge may be pushed next to each other → repel → lattice shatters.
4. Often soluble in water
   1. Water is polar and can surround and separate ions (hydration), pulling them away from the lattice.

Covalent Bonding – Sharing Electrons and Simple Molecules

How Covalent Bonds Form

Each shared pair of electrons = one covalent bond.

Simple Covalent Molecules and Intermolecular Forces

1. Covalent substances like H₂, O₂, HCl, CH₄ are often simple molecules:
   1. Inside each molecule: strong covalent bonds.
   2. Between molecules: only weak intermolecular forces.
2. This explains their typical properties:
   1. Low melting and boiling points
      1. Only weak forces between molecules need breaking, not the strong covalent bonds.
   2. Poor conductors of electricity
      1. No free ions or delocalised electrons to carry charge.
   3. Often gases or liquids at room temperature, or low-melting solids (e.g. iodine).

Polarity and Electronegativity

1. In some covalent bonds, electrons are not shared equally.
   1. Electronegativity = how strongly an atom attracts the shared electrons.
   2. If one atom is more electronegative, the bond becomes polar.
2. We show this with partial charges:
   1. δ⁻ (delta minus) = slightly negative
   2. δ⁺ (delta plus) = slightly positive
3. Dipole = separation of charge within a bond or molecule (δ⁺…δ⁻).

Example

Hydrochloric acid HCl

• Cl is more electronegative than H.
• The shared pair is pulled closer to Cl → Cl is δ⁻, H is δ⁺.
• The bond is polar covalent and the molecule is polar.

Why Polarity Matters

1. Affects solubility:
   1. “Like dissolves like” – polar substances dissolve in polar solvents (e.g. HCl in water).
   2. Non-polar substances dissolve in non-polar solvents (e.g. hydrocarbons in hexane).
2. Affects boiling points (polar molecules usually have stronger intermolecular forces than similar-sized non-polar molecules).

Metallic Bonding – Lattices of Ions and a Sea of Electrons

How Metallic Bonds Form

1. In metallic bonding, metal atoms:
   1. Lose their outer shell electrons and become positive ions.
   2. These electrons become delocalized – free to move throughout the structure.
   3. The metal is now a lattice of positive ions in a “sea” of delocalized electrons.
2. There is a strong electrostatic attraction between:
   1. The positive metal ions and
   2. The negative electron sea.
3. This is the metallic bond.

Structure: Giant Metallic Lattice

1. Regular 3D lattice of positive ions
2. Surrounded by a “cloud” of delocalized electrons
3. All ions are arranged in layers that can slide over each other.

How Structure Explains Metallic Properties

1. Good electrical and thermal conductors
   1. Delocalized electrons can move and carry charge and thermal energy.
2. Malleable and ductile
   1. Layers of ions can slide while the electron sea still holds the structure together → metals can be bent, hammered, drawn into wires.
3. Generally high melting and boiling points
   1. Strong attraction between ions and electrons requires lots of energy to overcome (especially in metals with high charge and many delocalized electrons).
4. Alloys
   1. Mixing metals (or metals + other elements) distorts the lattice → harder, stronger, less malleable (useful in construction, tools, etc.).

Comparing Structures: Ionic, Covalent and Metallic

How Bonding and Structure Explain Key Properties

1. Melting and boiling point
   1. Giant ionic: high – many strong ionic bonds to break.
   2. Simple covalent: low – only weak intermolecular forces to overcome.
   3. Metallic: usually high – strong attraction between ions and delocalised electrons.
2. Electrical conductivity
   1. Ionic solids: do not conduct – ions fixed in place.
   2. Ionic molten/aqueous: conduct – ions free to move and carry charge.
   3. Simple covalent: do not conduct – no free ions or electrons (exceptions in advanced courses like graphite).
   4. Metals: conduct well – delocalized electrons freely move through the lattice.
3. Solubility
   1. Ionic compounds: often soluble in polar solvents like water (ions are stabilized by ion–dipole interactions).
   2. Simple covalent:
      1. Polar molecules: often soluble in water.
      2. Non-polar molecules: soluble in non-polar solvents (e.g. hexane).
   3. Metals: generally insoluble in both water and organic solvents (though they may react with acids, etc.).
4. Mechanical properties
   1. Ionic: brittle – shifting layers can bring like charges together → strong repulsion → crystal shatters.
   2. Metals: malleable & ductile – layers of ions can slide without breaking the metallic bond.
   3. Simple covalent: often soft or gas/liquid – molecules move relatively easily because intermolecular forces are weak.