Having covered some weak intramolecular forces in my posts on hydrogen bonds and van der Waals forces, I ventured into the world of the strong forces last month with ionic bonds. This month I'll be looking at metallic bonding, the forces that hold together the atoms of all pure metals. There are a lot of metals in the periodic table, so for the sake of simplicity, I'll be breaking this post into two sections. First, a sort of GCSE explanation of metallic bonding in the alkali-metals, i.e things like sodium, potassium and other things that don't really pop into people's head when they think the word 'metal'. Then I'll take a look at the transition metals, just the pure metallic form for now because the many exciting properties of transition metals are a subject for several separate blog posts!

1) Alkali Metals

The periodic table, with the alkali metals circled in red.

The alkali metals have very few electrons (either one or two) in their outermost shell. As discussed in the post on ionic bonds, one way for them to achieve a full outer shell is to form an ion - loose that outer shell electron to form a positive ion and then bond with a negative ion in a giant lattice. Another way is for all the atoms of one element (i.e all Na atoms or Mg atoms) to form a lattice just composed of atoms, with the outer electons floating around freely in the spaces between. This is the essence of metallic bonding, which is sometimes described as a lattice of positive ions in a sea of electrons (even though they are not strictly speaking actual 'ions' - metallic elements will be written as Na metal rather than Na+ ions)

A stylised diagram of metallic bonding, from wikimedia commons, credit link below.

A stylised diagram of metallic bonding, from wikimedia commons, credit link below.

For the alkali metals, there are only one or two electrons from each atom actually participating in this delocalised sea of electrons, which explains why these metals are so soft and can be cut with a knife. It also explains their generally volatile nature, they simply aren't held together very well, and metals like potassium and caesium will catch fire if you get them wet.

2) Transition metals

Transition metals are the block described as "a little complicated" - circled here in red.

The general idea of metallic bonding is the same for the transition metals as it is for any other metals - a delocalised sea of electrons surrounding a lattice of positive ions. However in order to explain the properties of the transition metals I have to go back a bit and explain why they are complicated and this involves orbitals and a certain amount of quantum. I introduced in the ionic bond post the idea that each shell of electrons has eight electrons in it, and while this is true(ish) for the first three shells, it gets a bit complicated in the fourth.

Without getting too involved for now (I might write a post on orbital shapes later ...), the transition metals have a lot more than eight electrons in their outer shell - between the third and the fourth shell extra electron orbitals are introduced. Not only does this allow more electrons to be present, but these orbitals are all at very similar distances from the nucleus and therefore can all participate in bonding. It's probably easiest just to think of transition metals as surrounded by a blur of many easy-going electrons, as well as some empty orbitals that are happy to accept a few extra electrons.

This means that rather than having one free electron for each atom in the lattice, some transition metals will have around five or six. This is why copper (for example) is so much stronger than magnesium. The metallic bond also explains why metals can conduct electricity; free-flowing electrons are what electricity is. The lattice is rigid, but once melted a little it can be bent and stretched and shaped in a way that ionic compounds like salt cannot, which is why copper can be stretched into wires and aluminium can be bent into cars. It's a property that we take for granted, but it only happens in metals.

Although metalic bonds are fascinating, they aren't all that biochemically important. Metals are trace elements in organic lifeforms, but they tend to be bonded to other things, or in ion form rather than as the solid metal. Which is why next month I'll be covering covalent bonding - the forces which hold together all organic life on earth.


Credit link for metallic bond image.