Metallic Bonding

Think of metals as having a "sea of electrons" that's free to move around. When metal atoms bond together, they form giant structures with regularly arranged atoms, but here's the clever bit - the electrons from their outer shells become delocalised, meaning they're no longer tied to specific atoms.

This sharing of delocalised electrons creates strong metallic bonds. The strength comes from the attraction between the positively charged metal ions and the negatively charged delocalised electrons moving freely through the structure.

These free electrons are the reason metals have their unique properties. They conduct electricity because the delocalised electrons carry electrical charge, and they conduct heat as these electrons transfer energy efficiently through the closely packed structure.

Quick tip: Remember that metallic bonding only happens between metal atoms - it's what gives metals their "metallic" properties!

Metals are also malleable and ductile (you can hammer and stretch them) because the layers of metal ions can slide over each other when force is applied. Plus, they have high melting and boiling points since you need loads of energy to break those strong metallic bonds.

Metallic Bonding Trends

Here's something that'll help you predict properties across the periodic table. As you move across any period, metals get higher melting and boiling points.

This happens because there are more delocalised electrons available as you go across a period. More delocalised electrons means stronger attraction between the ions and free electrons, which creates stronger bonds overall.

Exam tip: This trend question comes up regularly - just remember more electrons equals stronger bonds!

It's a straightforward pattern once you get it, and it explains why some metals are much harder to melt than others.

Ionic Bonding

Ionic bonding happens when metals meet non-metals, and it's basically an electron transfer situation. Metal atoms lose electrons to become positive ions, whilst non-metal atoms gain those electrons to become negative ions.

The key idea is that both atoms are trying to achieve full outer shells. Metals happily give up electrons (they've got few in their outer shell anyway), and non-metals readily accept them (they're nearly full already).

You'll need to draw dot and cross diagrams to show this electron transfer. Take sodium chloride - sodium (2,8,1) gives its outer electron to chlorine (2,8,7), creating Na⁺ and Cl⁻ ions.

Memory trick: Metals GIVE electrons (they're generous!), non-metals TAKE electrons (they're greedy!).

For compounds like lithium oxide, you need two lithium atoms to give up one electron each to fill oxygen's outer shell, since oxygen needs two electrons to complete its shell.

Covalent Bonding

Covalent bonding is completely different from ionic bonding - instead of transferring electrons, atoms share electrons to achieve full outer shells. This only happens between non-metal atoms.

Single bonds involve sharing one pair of electrons. Water (H₂O) is a perfect example - the oxygen atom shares electrons with two hydrogen atoms, giving everyone a full outer shell.

You'll also encounter double bonds where two pairs of electrons are shared. Oxygen gas (O₂) and carbon dioxide (CO₂) are classic examples of this.

Quick check: If it's metal + non-metal = ionic bonding. If it's non-metal + non-metal = covalent bonding!

The dot and cross diagrams for covalent bonding show the shared electron pairs between atoms, rather than the complete electron transfer you see in ionic bonding.

Giant Ionic and Simple Molecular Structures

Giant ionic structures form when ionic compounds arrange themselves into regular patterns called ionic lattices. The oppositely charged ions attract each other in a repeating 3D arrangement that maximises attraction and minimises repulsion.

These structures have high melting and boiling points because of the strong electrostatic forces between ions. They only conduct electricity when molten or dissolved - that's when the ions are free to move and carry charge.

Simple molecular structures are completely different beasts. These are small molecules like hydrogen, water, and carbon dioxide held together by strong covalent bonds within each molecule, but only weak forces between different molecules.

Key insight: It's the weak forces BETWEEN molecules that determine the properties, not the strong bonds WITHIN molecules!

This explains why simple molecular compounds have low melting and boiling points - you're only breaking the weak intermolecular forces, not the strong covalent bonds. They also don't conduct electricity since there are no free electrons or ions.

Giant Covalent Structures

Giant covalent structures are like simple molecules on steroids - thousands of atoms held together by covalent bonds in massive lattices. The sheer number of bonds makes these structures incredibly strong.

Diamond has each carbon atom bonded to four others in a 3D network. It doesn't conduct electricity (no delocalised electrons) but it's brilliant for drill bits and glass cutting because of its extreme hardness.

Graphite is more interesting - each carbon bonds to only three others, creating layers. It conducts electricity because of delocalised electrons between the layers, and the layers can slide over each other, making it perfect for pencils and lubricants.

Exam favourite: You'll definitely get asked about diamond vs graphite properties - remember the bonding pattern determines everything!

The key difference between diamond and graphite shows how the same element can have completely different properties based on how the atoms are arranged and bonded.