Metallic Bonding & Melting Point Trends

Define metallic bonding

Lattice of positive ions surrounded by a 'sea' of delocalised electrons

Strength of metallic bonding

Strong electrostatic attraction between ions and delocalised electrons

Effect of ion charge on metallic bonding

Higher charge → stronger attraction, more delocalised electrons

Effect of ion size on metallic bonding

Larger ions → weaker attraction due to lower charge density

Metallic properties

Good conductors, malleable, high melting points, nearly always solid

Reason metals conduct

Delocalised electrons move freely and carry charge

Reason metals are malleable

Positive ion layers slide without breaking metallic bonding; electrons maintain lattice

Define macromolecular lattice

Giant covalent lattice of atoms covalently bonded

Diamond structure

Each carbon bonded to four others in a rigid 3D lattice

Properties of diamond

Hard, very high melting point

Graphite structure

Each carbon bonded to three others in flat layers; one delocalised electron per carbon atom

Properties of graphite

Conducts electricity, layers slide easily → good lubricant

Forces in graphite

Strong covalent bonds within layers, weak London forces between layers

Graphene structure

Single 2D sheets of hexagonal carbon rings, one atom thick

Properties of graphene

Strong, rigid, lightweight, conducts electricity via delocalised electrons

Melting points across Period 2

Li & Be → metallic bonding, increase with ion charge; B & C → giant covalent, very high; N, O, F, Ne → simple molecular, weak London forces, low

Melting points across Period 3

Na, Mg, Al → metallic bonding, increase with ion charge; Si → macromolecular, very high; P, S, Cl → simple molecular, weak London forces, low; Ar → noble gas, London forces, very low

Reason for low melting points in simple molecules

Weak London forces require little energy to overcome

Reason for high melting points in giant structures

Strong covalent or metallic bonds require lots of energy to break