Introduction to Covalent Bonding
Why is diamond so hard, yet graphite so soft , even though both are made entirely of carbon? The answer lies in how carbon atoms bond and arrange themselves. Understanding covalent bonds is the key to unlocking the diverse properties of materials all around us.
Covalent bond: A covalent bond is a chemical bond formed by the electrostatic attraction between a shared pair of electrons and the positively charged nuclei of the bonded atoms. Covalent bonds form between non-metal atoms , either of the same element (e.g., , ) or different non-metal elements (e.g., , ).
As two atoms approach each other:
- Each nucleus attracts the electrons of the other atom (attractive force).
- The two nuclei and the two electron clouds repel each other (repulsive force).
- At the optimal bond distance, attractive forces dominate , a stable covalent bond is established.
Think of two magnets being drawn together. The attraction between their opposite poles is analogous to the nucleus–electron attraction in a covalent bond. The bond length is the distance at which attraction and repulsion are perfectly balanced.
Covalent bonding is what connects non-metal atoms into molecules , from the simplest to complex biological macromolecules.
Single, Double, and Triple Bonds
Atoms can share one, two, or three pairs of electrons to form different types of covalent bonds.
Single bond: A single bond occurs when two atoms share one pair of electrons, represented by a single line (,).
Double bond: A double bond forms when two atoms share two pairs of electrons, represented by a double line (=).
Triple bond: A triple bond occurs when two atoms share three pairs of electrons, represented by a triple line (≡).
Examples:
- : (single bond)
- : (double bond)
- : (triple bond)
Bond Length and Bond Strength
Bond length: Bond length is the average distance between the nuclei of two bonded atoms.
Bond strength: Bond strength (bond enthalpy) is the energy required to break one mole of a given bond in the gaseous state.
As the number of shared electron pairs increases:
- Bond length decreases , more shared electrons pull the nuclei closer together.
- Bond strength increases , more electrons create a stronger electrostatic attraction.
| Bond Type | Example | Bond Enthalpy (kJ mol⁻¹) | Bond Length (pm) |
|---|---|---|---|
| C–C (single) | ethane | 346 | 154 |
| C=C (double) | ethene | 614 | 134 |
| C≡C (triple) | ethyne | 839 | 120 |
Values from IB Data Booklet, Section 11.
Shorter bonds are stronger bonds. The nuclei are held more tightly together by a greater number of shared electrons. This is a key relationship to remember for exam questions on bond properties.
When drawing double bonds, students often forget that each atom must still satisfy the octet rule. For example, in , each oxygen atom has four non-bonding electrons in addition to the shared double bond.