A covalent compound is a compound formed when two or more non-metal atoms share electrons to achieve a stable electron configuration. Unlike ionic compounds — where electrons are transferred from one atom to another — covalent bonds involve the sharing of electron pairs. Common examples include water (H₂O), carbon dioxide (CO₂), methane (CH₄), ammonia (NH₃), and hydrogen chloride (HCl). Covalent compounds make up the majority of organic molecules and are found in every living organism.
A covalent compound is formed by the sharing of electron pairs between atoms (usually non-metals).
Two types: polar covalent (unequal sharing — H₂O, HCl) and non-polar covalent (equal sharing — Cl₂, CH₄).
Covalent bonds may be single (H₂), double (O₂), or triple (N₂) depending on the number of shared electron pairs.
Most covalent compounds have low melting/boiling points and are gases or liquids at room temperature.
Generally poor electrical conductors — no free ions (exception: graphite).
Solubility follows 'like dissolves like': polar in polar solvents, non-polar in non-polar solvents.
Network covalent solids (diamond, SiO₂) are exceptions with very high melting points and hardness.
All organic (carbon-based) compounds are covalent.
Covalent bonds form when two atoms — usually non-metals — each contribute one or more electrons to a shared pair. The shared pair of electrons is attracted to the nuclei of both atoms, holding them together.
Types of covalent bonds by number of shared pairs: • Single bond: one shared pair (e.g., H–H in H₂, H–Cl in HCl) • Double bond: two shared pairs (e.g., O=O in O₂, O=C=O in CO₂) • Triple bond: three shared pairs (e.g., N≡N in N₂, H–C≡N in HCN)
Why non-metals? Non-metals have high electronegativity and tend to share electrons rather than lose them. Metals typically lose electrons (forming ionic bonds), while non-metals share electrons (forming covalent bonds).
Octet rule: atoms form covalent bonds to achieve 8 electrons in their outer shell (2 for hydrogen).
Covalent compounds are broadly divided into two types based on how evenly electrons are shared:
Non-polar covalent compounds: • Electrons are shared equally between atoms of the same or similar electronegativity • No partial charges — the molecule has no dipole moment • Examples: Cl₂ (chlorine gas), O₂, N₂, CH₄ (methane), CCl₄ (carbon tetrachloride)
Polar covalent compounds: • Electrons are shared unequally — the more electronegative atom pulls electrons closer • Partial negative charge (δ−) on the more electronegative atom; partial positive (δ+) on the other • Examples: H₂O (water), HCl, NH₃ (ammonia), HF (hydrogen fluoride)
Note: CO₂ has polar bonds (C=O) but is non-polar overall because the dipole moments cancel out due to its linear shape.
Network covalent solids (exceptions): • Diamond: carbon atoms bonded in a giant covalent network — extremely hard, very high melting point • Graphite: layered covalent structure — conducts electricity (an exception among covalent compounds) • Silicon dioxide (SiO₂/quartz): network solid with very high melting point
Physical state at room temperature: • Most are gases (CO₂, CH₄, HCl, NH₃), liquids (H₂O, ethanol), or soft solids (sugar, wax) • Network covalent solids (diamond) are exceptions with very high melting points
Melting and boiling points: • Generally low — molecules are held together by weak intermolecular forces, not strong ionic bonds • Examples: ethanol boils at 78°C, methane at −162°C • Exceptions: diamond (melts above 3500°C) and other network covalent solids
Electrical conductivity: • Generally poor conductors of electricity — no free ions or electrons to carry charge • Exception: graphite (has delocalised electrons in layers)
Solubility: • Non-polar covalent compounds dissolve in non-polar solvents (like benzene, hexane) • Polar covalent compounds (like HCl, ethanol) dissolve in polar solvents (like water) • Rule: like dissolves like
Hardness: • Usually soft and brittle (molecular solids) • Exception: diamond is the hardest natural substance
Molecule | Formula | Bond type Water | H₂O | Polar covalent Carbon dioxide | CO₂ | Non-polar (polar bonds that cancel) Methane | CH₄ | Non-polar covalent Ammonia | NH₃ | Polar covalent Hydrogen chloride | HCl | Polar covalent Oxygen | O₂ | Non-polar covalent Nitrogen | N₂ | Non-polar covalent (triple bond) Sugar (sucrose) | C₁₂H₂₂O₁₁ | Polar covalent Ethanol | C₂H₅OH | Polar covalent Glucose | C₆H₁₂O₆ | Polar covalent
Covalent compounds dominate biochemistry — proteins, DNA, carbohydrates, fats are all built on covalent bonds.
A covalent compound is a compound formed when two or more non-metal atoms share electron pairs to form covalent bonds. The shared electrons hold the atoms together. Examples include water (H₂O), carbon dioxide (CO₂), ammonia (NH₃), and methane (CH₄). Covalent compounds generally have low melting points and are poor conductors of electricity.
In covalent compounds, electrons are shared between atoms — usually non-metals. In ionic compounds, electrons are transferred from one atom to another (usually metal to non-metal), creating oppositely charged ions that attract each other. Covalent compounds generally have lower melting points, are poor conductors, and exist as molecules. Ionic compounds are hard crystalline solids with high melting points and conduct electricity when dissolved in water or melted.
Three common covalent compounds are: (1) Water (H₂O) — polar covalent, formed by oxygen and two hydrogen atoms sharing electrons; (2) Carbon dioxide (CO₂) — formed by one carbon and two oxygen atoms with double bonds; (3) Methane (CH₄) — non-polar covalent, formed by carbon and four hydrogen atoms. All are formed by non-metal atoms sharing electrons.
Properties of covalent compounds: (1) Low melting and boiling points (except network covalent solids like diamond). (2) Exist as gases, liquids, or soft solids at room temperature. (3) Poor conductors of electricity (no free ions). (4) Non-polar covalent compounds dissolve in non-polar solvents; polar covalent compounds dissolve in polar solvents. (5) Usually not hard (except diamond). (6) Molecules — not made of ions.
In polar covalent compounds, electrons are shared unequally because one atom is more electronegative than the other, creating partial charges (δ+ and δ−). Examples: H₂O, HCl, NH₃. In non-polar covalent compounds, electrons are shared equally — no partial charges. Examples: Cl₂, CH₄, N₂. Polar covalent compounds tend to dissolve in water; non-polar ones dissolve in organic solvents.
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