Although ionic and covalent bonds are more common, metallic bonding describes a lattice of cations surrounded by a ‘sea’ of valence electrons🌊. The nucleus and core electrons of the metal stay in place, but the valence electrons are very mobile.
Electrons usually belong to a certain atom but in metals, they move so much that they don't belong to one, single atom.
This ‘sea’ of valence electrons contributes to metal properties:
Metals conduct electricity⚡ - because of delocalization of electrons in metals, electrons can move freely. This allows metals to conduct charge or, as we'll see in a future unit, be used in redox reactions.
Metals are malleable and ductile - Metals, because they are less rigid, can be bent and spun into wire🔌.
When comparing properties among the different solids, remember this chart:
|Type of Solid||Form of Unit Particles||Forces Between Particles||Properties||Examples|
|Molecular🧊||Atoms or Molecules||LDFs, dipole-dipole, hydrogen bonds||fairly soft, low melting point, bad conductor||Argon, methane, sucrose, dry ice|
|Covalent-Network💎||Atoms connected in a network of covalent bonds||Covalent bonds||Very hard, very high melting point, bad conductor||diamond, quartz|
|Ionic🧂||Positive and negative ions||Electrostatic attractions||Hard and brittle, high melting point, bad conductor||salts (NaCl)|
|Metallic✨||Atoms||Metallic bonds||varying hardness and melting points, good conductor, malleable, ductile||metals! Cu, Fe, Al|
Table Courtesy of unknown source
Right now, you should only be very familiar with the two bolded rows. The others are covered in unit 3 in more depth!
Metals can also bond with each other and become alloys. Alloys can be formed when two metals are in their liquid form being mixed together. When this mixture cools, the alloy is formed.
There are two types:
Image Courtesy of Chemistry LibreTexts
Alloys are harder and stronger than pure metals
because the added elements distorts the structure and properties. Alloys are also less malleable than pure metals.
Check your Understanding
*This question is similar to one discussed on the Advanced Placement YT*
It goes over content reviewed this key topic and the previous key topic.
A student ran an experiment to see if the following solids conduct electricity.
|Solids||Does it conduct electricity?|
(a) Explain the results the student saw.
Sample response (a)
This student found that a sample of iron conducted electricity since it is a metal. Metals have delocalized valence electrons, usually displayed by the sea of electrons diagram, allowing them to be good conductors of electricity.
The student found that the sample of FeCl2 didn't conduct electricity because it is an ionic solid. Ionic solids have a lattice structure. Therefore, electrons cannot move freely and the sample didn't conduct electricity.
(b) Is there anything that could have been different in this experiment to see the FeCl sample conduct electricity?
Recall: As long as there are mobile valence electrons, the sample will conduct electricity. There are two ways to do that. Either of the following responses are acceptable.
Melt the FeCl2 solid and then test it for conductivity. The liquid FeCl2 would conduct electricity since the ions would be mobile and able to flow.
Dissolve the FeCl2 solid into water. In an aqueous solution, the ions are able to flow and conduct electricity.