Question

Considering that electricity is the flow of electrons from one place to another, propose an explanation for why metals are generally good conductors of electricity, whereas covalently bonded compounds are not.

Short answer

Metals are good conductors of electricity because they have 'free electrons' that can move and carry a charge. Covalently-bonded compounds do not generally conduct electricity because their electrons are tightly bound and therefore they lack 'free electrons' to carry a charge.

Step-by-step solution

Explain Electron Configuration in Metals

Metals typically have few electrons in their outermost shells, and they tend to lose these electrons easily. This creates what is referred to as a 'sea of free electrons' that are not associated with any particular atom and are free to move within the metal lattice structure. These free, or delocalized, electrons can move and carry a charge, which allows for the flow of electricity.

Explain Electron Configuration in Covalent Compounds

In contrast to metals, covalently-bonded compounds have tightly bound electrons. This is because they form by sharing electrons between atoms to fulfill their outer energy levels. These shared electrons are not free to move about the compound in the same way the 'free electrons' in metals are.

Correlation Between Electron Configuration and Conductivity

The fact that metals typically have 'free electrons' allows them to conduct electricity. The current, a flow of electrons, can pass through the metallic structure using these free electrons. However, covalently-bonded compounds with their tightly bound electrons restrict the movement of electricity because there are no 'free electrons' available to carry the electric charge.