metallic bonding

The properties of metals

The general properties of metals can be summarised in the image below:

the properties of maetals

We can use the general properties of metals shown in the image above to help us figure out what the structure of a metal looks like. The table below tries to offer an explanation for some of the properties of metals and relate this property to a possible feature of the metal structure.

Metal property How this relates to its structure
Metals are good conductors of electricity and heat There must be free or delocalised electrons within the structure to conduct the electricity and heat.
Metals have high melting and boiling points. Metals must have a giant structure with lots of strong bonds.
Metals are malleable (can be hammered into shape) and ductile (can be pulled into wire). There are layers of particles that are able to slide over each other.
Metals are shiny. Free or delocalised electrons are able to reflect light at the surface.
Metals are hard and dense The particles within the metal structure are packed tightly together.

The structure of metals

These basic properties of metals allows us to suggest that metals consist of a giant structures of ion with free delocalised electrons. We also already know that metals are found on the left hand side of the periodic table in groups 1,2 and 3 and that they tend to lose electrons in their reactions to form positively charged ions. This leads us to the model shown below. The dark grey balls represent positively charged metal ions (atoms which have lost their outer shell electrons). These electrons are delocalised and free to move through the giant structure of ions. This might seem odd since we might expect a giant structure of ions, all with a positive charge to repel each other and so the structure would simply break down. However we need to think about the delocalised electrons within the structure.

metallic bonding

metallic bonding The negatively charged electrons are attracted to the positively charged metal ions and this prevents the metal ions from repelling each other. The electrons are attracted to the metal ions around them. This attraction of the negatively charged electrons to the neighbouring positively charged metal ions is called a metallic bond. A key feature of this bond is the fact that metallic bonds are not permanent but are constantly breaking and reforming as the electrons move freely around the structure.

The fact that these metallic bonds are not permanent but are constantly breaki and reforming allows the layers of ions to slide over each other. This is shown in the diagram below, which shows a pushing force being applied to the top two layers of ions in a metal structure. The metallic bonds in these layers immediately break and the layers slide along, but as soon as they stop moving the metallic bonds immediately reform. This is why metals are malleable (hammered into shape) and ductile (can be pulled into wires).

layers sliding in metallic bonding

Properties in detail

Metals generally have high melting points, high densities and are good electrical and thermal conductors, though these properties can vary significantly as we cross the periodic table. Let us examine some these typical properties of metals in a little more detail

Bond strength, conductivity and density of metals

The strength of the metallic bond depends mainly on two factors: The d-block metals not only release their electrons in their s-subshell but also the d-subshell electrons can be delocalised. This means that metal ions with higher charges and a smaller radius will be produced. This means stronger bonds within the metal structure which will lead to metals which are hard, dense and have high melting points. We can extend this arguement to trends down a group. Obviously as we descend a group in the periodic table the size of the metal atoms and ions will increase, this means that the attraction to the delocalised electrons will be weaker, so the strength of the metallic bond will be reduced. This means a decrease in the melting points for metals as we descend a group.
Metals are good electrical conductors due to the presence of the delocalised electrons. The more delocalised electrons there are the better an electrical and thermal conductor the metal will be. Aluminium for example is a better electrical conductor than magnesium because it has 3 valency electrons ([Ne]3s23p1) while magnesium ([Ne]3s2) only has 2 electrons in its valence shell



Key Points



Practice questions

Check your understanding - Questions on metallic bonding

Next