Covalent bonding is another fundamental way atoms achieve stability, primarily through the sharing of electrons between two non-metal atoms.

Principle: Sharing of Electrons Between Two Non-metals

• Covalent bonds typically form between two non-metal atoms. Non-metals tend to gain electrons to complete their outermost shell (achieve an octet). However, when two non-metals react, neither atom is strong enough to completely pull electrons away from the other to form ions.
• Instead, they achieve stability by sharing one or more pairs of valence electrons. The shared electrons are simultaneously attracted to the nuclei of both bonding atoms, creating a strong attractive force that holds the atoms together.
• This sharing allows each atom involved in the bond to effectively "count" the shared electrons towards its own octet (or duet for hydrogen), thereby reaching a stable electron configuration similar to a noble gas.

Single, Double, and Triple Bonds (Simple Examples):

Atoms can share different numbers of electron pairs, leading to different types of covalent bonds:
- Single Covalent Bond:
- Formed when two atoms share one pair of electrons (2 electrons total).
- Represented by a single line (—) between the atomic symbols.
- Example: Hydrogen gas (H₂)
- Each Hydrogen (H) atom has 1 valence electron. To achieve the stable electron configuration of Helium (2 electrons), each needs 1 more electron.
- They achieve this by sharing their single valence electrons, forming one shared pair.
- H — H (Each H now effectively has 2 electrons: its own 1 + 1 from sharing).
- Double Covalent Bond:
- Formed when two atoms share two pairs of electrons (4 electrons total).
- Represented by two parallel lines (=) between the atomic symbols.
- Example: Oxygen gas (O₂)
- Each Oxygen (O) atom has 6 valence electrons. To achieve an octet (8 electrons), each needs 2 more electrons.
- They achieve this by sharing two pairs of electrons.
- O = O (Each O now effectively has 8 electrons: its own 6 + 2 from sharing each of the two pairs).
- Triple Covalent Bond:
- Formed when two atoms share three pairs of electrons (6 electrons total).
- Represented by three parallel lines (≡) between the atomic symbols.
- Example: Nitrogen gas (N₂)
- Each Nitrogen (N) atom has 5 valence electrons. To achieve an octet (8 electrons), each needs 3 more electrons.
- They achieve this by sharing three pairs of electrons.
- N ≡ N (Each N now effectively has 8 electrons: its own 5 + 3 from sharing each of the three pairs). Triple bonds are the strongest and shortest type of covalent bond between two given atoms.

Simple Molecular Compounds:

When different non-metal atoms bond covalently, they form molecules of a compound. These are often referred to as simple molecular compounds or simply molecules.
- Water (H₂O):
- Oxygen (O) has 6 valence electrons and needs 2 more.
- Each Hydrogen (H) has 1 valence electron and needs 1 more (to be like Helium).
- One oxygen atom shares 1 electron with the first hydrogen, and 1 electron with the second hydrogen. Each hydrogen shares its 1 electron with the oxygen.
- This forms two single covalent bonds: H—O—H. The oxygen effectively has 8 valence electrons (6 + 1 + 1), and each hydrogen effectively has 2 valence electrons.
- Carbon Dioxide (CO₂):
- Carbon (C) has 4 valence electrons and needs 4 more.
- Each Oxygen (O) has 6 valence electrons and needs 2 more.
- The carbon atom forms a double bond with one oxygen atom and another double bond with the second oxygen atom.
- O = C = O. The carbon effectively has 8 electrons (4 + 2 + 2), and each oxygen effectively has 8 electrons (6 + 2).
- Methane (CH₄):
- Carbon (C) has 4 valence electrons and needs 4 more.
- Each Hydrogen (H) has 1 valence electron and needs 1 more.
- The carbon atom forms four single covalent bonds, one with each of the four hydrogen atoms.
- H | H—C—H | H
- The carbon effectively has 8 electrons (4 + 1+1+1+1), and each hydrogen effectively has 2 electrons.

Properties of Simple Molecular Compounds:

The nature of covalent bonds and the formation of discrete molecules give simple molecular compounds distinct properties:
- Low Melting Points:
- Within each molecule (e.g., H₂O), the covalent bonds between atoms are very strong.
- However, the forces between separate molecules (called intermolecular forces) are much weaker than the actual covalent bonds within the molecules.
- When you melt or boil a molecular compound, you are not breaking the strong covalent bonds within the molecules; you are only overcoming the weak intermolecular forces between the molecules.
- Because these intermolecular forces are weak, little energy is required to separate the molecules, resulting in low melting and boiling points.
- This is why many simple molecular compounds are gases or liquids at room temperature (e.g., H₂O is liquid, CO₂ is gas, O₂ is gas).
- Often Gases or Liquids at Room Temperature:
- Due to their low melting and boiling points, many simple molecular compounds exist as gases (like oxygen, nitrogen, methane, carbon dioxide) or liquids (like water, ethanol) at typical room temperatures. Only larger, more complex molecules tend to be solids (like sugar, which is a network of simple molecules).
- Poor Conductors of Electricity:
- Simple molecular compounds generally do not conduct electricity in any state (solid, liquid, or gas).
- This is because all their valence electrons are localized in specific covalent bonds (either shared between atoms or existing as lone pairs on individual atoms). There are no free-moving ions or delocalized electrons available to carry an electrical charge.
- Example: Pure water (H₂O) is a very poor conductor of electricity, unlike saltwater (NaCl dissolved in water).