Covalent Bonding: Sharing of Electrons Explained
Covalent bonding occurs when two atoms share one or more pairs of electrons, resulting in a stable molecule. This bond typically forms between non-metallic
TL;DR: Covalent bonding occurs when two atoms share one or more pairs of electrons, resulting in a stable molecule. This bond typically forms between non-met…
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Covalent Bonding: Sharing of Electrons: Covalent bonding occurs when two atoms share one or more pairs of electrons, resulting in a stable molecule. This bond typically forms between non-metallic
Covalent Bonding: Sharing of Electrons Explained Simply
Covalent bonding occurs when two atoms share one or more pairs of electrons, resulting in a stable molecule. This bond typically forms between non-metallic This is a key Chemistry concept for Class 10-12.
Characteristics of Covalent Bonds
In covalent bonds, electrons are shared between atoms rather than transferred. A single covalent bond involves the sharing of one pair of electrons (two electrons total). A double bond shares two electron pairs (four electrons), and a triple bond shares three electron pairs (six electrons). Covalent bonds form between atoms with similar electronegativities where neither atom can completely remove electrons from the other.
Types of Covalent Bonds
Polar covalent bonds form when atoms have different electronegativities; the shared electrons spend more time around the more electronegative atom, creating partial charges. Non-polar covalent bonds form between identical atoms or atoms with very similar electronegativities; electrons are shared equally. Coordinate covalent bonds (dative bonds) occur when both electrons in the bond come from the same atom; once formed, they are identical to regular covalent bonds.
Covalent Compounds and Molecules
Covalent compounds exist as molecules held together by covalent bonds. Molecular compounds have low melting and boiling points because intermolecular forces (van der Waals forces, hydrogen bonds) are weak. They are usually liquids or gases at room temperature and insoluble in water but soluble in non-polar solvents. They do not conduct electricity in any state because they lack mobile charged particles.
Lewis Structures and Bonding
Lewis structures (dot structures) represent valence electrons as dots around atomic symbols. Shared electron pairs between atoms are shown as lines or pairs of dots. Each atom tries to achieve an octet (8 valence electrons) or duet for hydrogen (2 electrons). Nitrogen (N2) forms a triple bond N≡N. Oxygen (O2) forms a double bond O=O. Hydrogen (H2) forms a single bond H-H.
Electronegativity and Polarity
Electronegativity is the ability of an atom to attract shared electrons in a bond. When electronegativity difference is 0-0.4, the bond is non-polar covalent. When difference is 0.4-1.7, the bond is polar covalent. The dipole moment measures the polarity of a bond; more electronegative atoms pull electron density toward themselves. In water (H2O), oxygen is more electronegative than hydrogen, making it a polar molecule with a bent shape.
Factors Affecting Covalent Bond Strength
Bond strength depends on bond order; triple bonds are stronger than double bonds, which are stronger than single bonds. Shorter bonds are generally stronger because atoms are closer and overlap is better. Bond enthalpy is the energy required to break a bond; strong bonds require more energy. N≡N has exceptionally high bond enthalpy (946 kJ/mol) making nitrogen extremely stable and unreactive at room temperature.
Worked Examples
Example 1. Draw the Lewis structure of Carbon Dioxide (CO2).
Solution: Carbon has 4 valence electrons, each oxygen has 6. C is central: O=C=O. Each O forms a double bond with C. C has 8 electrons around it (octet), each O has 8. Total valence electrons used: 4 + 6 + 6 = 16.
Example 2. Explain why N2 is more stable than O2.
Solution: N2 has a triple bond (N≡N) with bond enthalpy of 946 kJ/mol. O2 has a double bond (O=O) with bond enthalpy of 498 kJ/mol. The stronger triple bond makes N2 significantly more stable and less reactive.
Example 3. Is HCl polar or non-polar? Explain.
Solution: HCl is polar covalent. Electronegativity of Cl (3.16) is much higher than H (2.20); difference is 0.96. The shared electrons are pulled closer to Cl, creating a dipole with partial negative charge on Cl and partial positive charge on H.
Where You See This in Real Life
- Water (H2O) is a polar covalent compound essential for all life; its bent geometry and polarity allow it to dissolve many substances.
- Oxygen gas (O2) sustains aerobic respiration through covalent bonding; its reactivity is due to the double bond between atoms.
- Diamond is a covalent network structure where each carbon atom shares four electrons with four other carbons in an extremely strong 3D lattice.
- Plastics and polymers are long chains of covalently bonded carbon atoms; their properties depend on the types of covalent bonds and side groups.
Common Mistakes Students Make
- Confusing covalent and ionic bonding; covalent involves sharing, ionic involves transfer of electrons.
- Thinking all covalent bonds are non-polar; many covalent bonds are polar due to electronegativity differences.
- Writing incorrect Lewis structures by not accounting for all valence electrons or forgetting lone pairs.
- Assuming covalent compounds always dissolve in water; non-polar covalent compounds are water-insoluble but dissolve in non-polar solvents.
Frequently Asked Questions
What is the difference between a single, double, and triple bond?
A single bond is one shared pair of electrons, a double bond is two shared pairs, and a triple bond is three shared pairs. More shared pairs mean stronger bonds and shorter bond lengths.
Why do covalent compounds have low melting points?
Covalent compounds consist of discrete molecules held together by weak intermolecular forces (van der Waals forces or hydrogen bonds). These forces are much weaker than the covalent bonds within molecules, so molecules separate easily when heated.
Can metals form covalent bonds?
Metals typically do not form covalent bonds with each other. However, some metalloids and metals can form covalent bonds with non-metals. For example, silicon dioxide is a covalent network of Si-O bonds.