Covalent Bonds and Their Properties

Why Do Atoms Bond?

  • Noble gases have stable electron configurations with full outer energy levels, leading to lower potential energy and less tendency to form compounds.
  • Other elements bond to achieve stability.

Gaining Stability

  • The stability of atoms is linked to energy; lower energy states are more stable.
  • Metals and nonmetals can transfer electrons to form ions that achieve stable noble-gas configurations.

What is a Covalent Bond?

  • Covalent bonds are formed when atoms share electrons to achieve stable electron configurations.
  • Key questions:
    • How is sharing achieved?
    • Are there different methods of electron sharing?
    • How do properties of covalent compounds differ from ionic compounds?

Shared Electrons

  • Atoms in nonionic compounds share electrons; this sharing creates a covalent bond.
  • Molecules are formed when two or more atoms bond covalently.
  • Shared electrons count as part of both atoms' outer energy levels.
  • Covalent bonds generally occur between elements close on the periodic table, primarily nonmetals.

Covalent Bond Formation

  • Diatomic molecules, such as H₂, N₂, O₂, F₂, Cl₂, Br₂, and I₂, form when two atoms of the same element share electrons.

Covalent Bond Formation (cont.)

  • Example: Fluorine (F) with the electron configuration of 1s²2s²2p⁵ needs one electron to complete its octet.
  • As fluorine atoms approach, their nuclei attract each other until they share electrons and form F₂, with one shared pair and three lone pairs.

Single Covalent Bonds

  • A single covalent bond forms when only one pair of electrons is shared.
  • The shared electron pair is known as the bonding pair.
  • Example: Hydrogen molecule (H₂), where the shared pair gives each hydrogen a full outer level.
  • Electron dot (Lewis) diagrams can depict valence electrons in molecules.

Group 17 and Single Bonds

  • Halogens (group 17) need one additional electron to form single covalent bonds.
  • They can bond with other nonmetals, including identical atoms (e.g., F₂, Cl₂).

Group 16 and Single Bonds

  • Group 16 elements can share two electrons and form two covalent bonds.
  • Example: Water (H₂O)—each H shares one electron with O, resulting in stable configurations across all atoms.

Group 15 and Single Bonds

  • Group 15 elements can form three covalent bonds.
  • Example: Ammonia (NH₃)—Nitrogen shares three electrons with three hydrogen atoms, leaving one lone pair.

Group 14 and Single Bonds

  • Group 14 elements can form four covalent bonds.
  • Example: Methane (CH₄)—Carbon needs four additional electrons provided by four hydrogen atoms, forming four bonds.

The Sigma Bond

  • Single covalent bonds are referred to as sigma bonds, represented by the Greek letter 𝛔.
  • A sigma bond forms when shared electrons are localized between the two nuclei.
  • Present in molecules like H₂O, NH₃, and CH₄.

Multiple Covalent Bonds

  • Multiple bonds involve sharing more than one pair of electrons.
  • Double and triple covalent bonds are examples of multiple bonds formed by elements like carbon, nitrogen, oxygen, and sulfur.
  • The number of covalent bonds typically equals the number of valence electrons needed to achieve an octet.

Double Bonds

  • A double bond is formed by sharing two pairs of electrons.
  • Example: Oxygen (O₂)—each O atom shares two electrons to complete their valence shell.

Triple Bonds

  • A triple bond occurs when three pairs of electrons are shared.
  • Example: Diatomic nitrogen (N₂)—each N atom shares three pairs of electrons forming a triple bond.

The Pi Bond

  • A multiple covalent bond includes one sigma bond and one or more pi bonds.
  • A pi bond forms through the overlap of parallel orbitals, occupying regions above and below the bonding axis.

The Strength of Covalent Bonds

  • Covalent bonds vary in strength; bond strength is influenced by:
    • Bond length
    • Type of bond
    • Energy
  • Nuclei attract shared electrons while repelling each other and electron pairs.

Bond Length

  • Bond strength is related to bond length, the distance between bonded nuclei at maximum attraction.
  • As more electron pairs bond, bond length decreases.
  • Shorter bond lengths generally indicate stronger bonds.

Bonds and Energy

  • Energy changes occur during bond formation and breaking. Bond formation releases energy; bond breaking requires energy input (bond-dissociation energy).
  • Always a positive value.

Bonds and Energy (cont.)

  • Total energy change in a chemical reaction depends on energy of bonds broken vs formed.
  • Endothermic reaction: More energy is needed to break reactant bonds than released in product formation.
  • Exothermic reaction: More energy is released during product formation than required to break reactant bonds.