* Octet Rule: Most atoms strive to have a full outer shell of electrons, resembling the stable electron configuration of noble gases. This is known as the octet rule (except for hydrogen and helium which aim for a duet of electrons).
* Electrostatic Attraction: Atoms that gain or lose electrons to achieve a full outer shell become charged particles called ions. Positively charged ions (cations) and negatively charged ions (anions) attract each other due to electrostatic forces, forming ionic bonds. This strong attraction holds them together in ionic compounds, a much more stable state than existing as individual atoms.
Here's a breakdown:
* Metals: Metals tend to lose electrons, becoming positively charged cations. They readily lose electrons to achieve a stable configuration, often forming ionic bonds with nonmetals.
* Nonmetals: Nonmetals tend to gain electrons, becoming negatively charged anions. They gain electrons to achieve a full outer shell, also forming ionic bonds with metals.
Examples:
* Sodium (Na): Sodium has one electron in its outer shell. It readily loses this electron to become a Na+ ion, achieving a stable configuration.
* Chlorine (Cl): Chlorine has seven electrons in its outer shell. It readily gains one electron to become a Cl- ion, achieving a stable configuration.
* Sodium chloride (NaCl): Sodium and chlorine react to form sodium chloride (table salt). The Na+ and Cl- ions are held together by strong electrostatic attraction in a crystal lattice structure.
Exceptions:
* Noble Gases: Noble gases already have a full outer shell of electrons, making them very stable and unreactive. They generally exist as single atoms.
* Covalent Bonding: Some atoms share electrons to achieve stability, forming covalent bonds. This is common among nonmetals.
In summary: The inherent instability of most atoms, driven by their desire to achieve a full outer shell of electrons, leads to them readily forming ions. These ions then bond together to form ionic compounds, creating a more stable state than existing as individual atoms.
