1. No Free Electrons: In covalent bonds, electrons are shared equally between atoms. These shared electrons are tightly bound to the atoms and are not free to move around the material.

2. Localized Electrons: Electrons in covalent bonds are localized in the space between the atoms. They are not free to move throughout the material, which is necessary for electrical conductivity.

3. Strong Bond Strength: Covalent bonds are generally strong, meaning it takes a lot of energy to break them and liberate the electrons. This makes it difficult for electrons to move freely.

4. Absence of Ions: Covalent compounds typically do not form ions, which are charged particles that can carry electric current.

In contrast, metals are good conductors because:

* They have free electrons called "delocalized electrons" that can move easily throughout the material.

* These electrons are not bound to any particular atom, making them highly mobile.

* The metallic bonds between metal atoms allow for this free movement of electrons.

Exceptions:

There are some exceptions to the general rule that covalent bonds are poor conductors.

* Graphite: A form of carbon with covalent bonds, but its structure allows for the movement of electrons within layers of carbon atoms. This makes graphite a relatively good conductor of electricity.

* Polymers: Some polymers can become conductive by doping with certain elements or through other modifications.

However, these are special cases, and generally, covalent bonds are considered poor conductors of electricity due to the localized nature of their electrons.