Explain with examples single, double, and triple covalent bonds.

Covalent bonds occur when atoms share electrons to achieve a stable electron configuration. The number of shared electron pairs between atoms determines whether the bond is single, double, or triple. Let’s break these down:

  • Single Covalent Bond: A single covalent bond is formed when two atoms share one pair of electrons. This is the most basic type of covalent bond.

    • Example: The hydrogen molecule (H₂) is formed when two hydrogen atoms share one electron each. This sharing allows both hydrogen atoms to achieve the stable configuration of two electrons in their outer shell, fulfilling the duplet rule.

    • Structure: H : H

  • Double Covalent Bond: A double bond occurs when two atoms share two pairs of electrons. This bond is stronger and shorter than a single bond.

    • Example: The oxygen molecule (O₂) forms a double bond when two oxygen atoms share two pairs of electrons. Each oxygen atom originally has six valence electrons and needs two more to complete its octet. By sharing two pairs of electrons, both atoms achieve a stable configuration with 8 electrons in their outer shells.

    • Structure: O = O

  • Triple Covalent Bond: A triple bond occurs when two atoms share three pairs of electrons. This bond is even stronger and shorter than a double bond.

    • Example: The nitrogen molecule (N₂) forms a triple bond between two nitrogen atoms. Each nitrogen atom has five valence electrons, needing three more electrons to complete their octet. By sharing three pairs of electrons, both nitrogen atoms achieve a stable electron configuration.

    • Structure: N ≡ N

In general, as more electron pairs are shared, the bond strength increases, and the bond length decreases. Triple bonds are the shortest and strongest, while single bonds are the longest and weakest.

Related Questions:

  1. Explain how elements attain stability.
  2. Describe the ways in which bonds may be formed.
  3. Describe the formation of a covalent bond between two non-metallic elements.
  4. Find the number of valence electrons in the following atoms using the periodic table:
  5. Represent the formation of cations for the following metal atoms using electron dot structures:
  6. A sample of sulfur from a volcano was analyzed to give the following composition of isotopes (Atomic No. of S = 16):
  7. An atom of an element has an atomic number of 9 and a mass number of 19:
  8. Define ionic bond and give an example.
  9. State octet and duplet rules.
  10. Explain the formation of a covalent bond between two nitrogen atoms.
  11. How does Al form a cation?
  12. Draw an electron cross and dot structure for an H₂O molecule.
  13. Describe the importance of noble gas electronic configuration.
  14. The atomic number of an element is 23, and its mass number is 56.
  15. The atomic symbol of aluminum is written as Al. What information do you get from it?
  16. Write the valence shell electronic configuration of an element present in the 3rd period and Group IIA.
  17. Define halogens.
  18. Which atom has a higher shielding effect, Li or Na?
  19. Explain why Na has higher ionization energy than K.
  20. Arrange the elements in each of the following groups in order of increasing ionization energy: