Chemical Bonding

Valence Electrons and Bonding

Valence electrons are the electrons in the outermost shell of an atom and are primarily involved in chemical bonding. These electrons occupy the highest energy level and determine the chemical reactivity and bonding capacity of the atom.

• Valence electrons: Electrons in the outermost shell, involved in bonding.

• Valency: The number of electrons an atom can share, lose, or gain to achieve a stable configuration.

• Example: Hydrogen has 1 valence electron (configuration: ), so its valency is 1.

Ionic Bonds

Ionic bonds are formed between metals and non-metals through the transfer of electrons, resulting in the formation of positively and negatively charged ions. This process leads to filled valence orbitals and stable electronic configurations.

• Formation: Occurs between a positively charged ion (cation, e.g., or ) and a negatively charged ion (anion, e.g., or ).

• Mechanism: Atoms bond ionically by giving up or accepting electrons.

• Example: Sodium chloride (): Sodium () electron configuration: Sodium loses one electron to become ; chlorine gains one electron to become .

Covalent Bonds

Definition and Formation

Covalent bonds involve the sharing of electrons between two atoms, typically non-metals. The shared electrons allow each atom to achieve a stable electronic configuration.

• Valency: Number of electrons an atom shares to complete its octet.

• Examples:

  • Hydrogen (): 1 electron (), valency = 1

  • Oxygen (): 8 electrons (), valency = 2

  • Nitrogen (): 7 electrons (), valency = 3

  • Carbon (): 6 electrons (), valency = 4

Types of Covalent Bonds

• Single bonds: Sharing of one pair of electrons (e.g., in ).

• Double bonds: Sharing of two pairs of electrons (e.g., in ethene).

• Triple bonds: Sharing of three pairs of electrons (e.g., in acetylene).

Polarity and Electronegativity

Polar and Nonpolar Covalent Bonds

The polarity of a covalent bond depends on the difference in electronegativity between the bonded atoms. Unequal sharing of electrons leads to polar covalent bonds, while equal sharing results in nonpolar covalent bonds.

• Electronegativity: Measure of the attraction of an atom for bonding electrons in molecules compared to that of the other atom.

• Periodic Trend: Electronegativity increases across a period and decreases down a group. Fluorine () is the most electronegative element.

Bond Classification by Electronegativity Difference

• Dipole moment: Polar bonds create dipoles, as in .

• Examples:

  • Acetonitrile (): bond is polar ( vs ).

  • and bonds are nonpolar; , , and -halogen bonds are polar.

Formal Charge

Definition and Calculation

Formal charge (FC) is a bookkeeping tool used to determine the charge distribution within a molecule. It helps in identifying the most stable Lewis structure.

• Formula:

• Examples:

  • Hydronium ion ():

  • Ammonium ion ():

  • Methane ():

  • Methyl anion ():

  • Acetonitrile (): , , ; overall FC = 0

Three-Dimensional Structure of Molecules

VSEPR Theory and Molecular Geometry

The three-dimensional structure of molecules is determined by the number of bonds and lone pairs around the central atom, known as the steric number (SN). The Valence Shell Electron Pair Repulsion (VSEPR) theory predicts molecular shapes based on minimizing electron pair repulsion.

• Steric number: SN = Number of bonds + number of lone pairs around an atom.

• Examples:

  • Methane (): SN = 4, tetrahedral geometry.

  • Ethene (): SN = 3, trigonal planar geometry.

  • Acetylene (): SN = 2, linear geometry.

Types of Covalent Bonds: Sigma and Pi Bonds

Sigma () Bonds

Sigma bonds are formed by the head-on overlap of atomic orbitals, typically s-orbitals. All single bonds are sigma bonds.

• Formation: Overlap of two s-orbitals.

• Example: bond in .

Pi () Bonds

Pi bonds are formed by the side-to-side overlap of two parallel p-orbitals. Double and triple bonds contain one sigma bond and one or two pi bonds, respectively.

• Formation: Overlap of two parallel p-orbitals.

• Example: double bond has one sigma and one pi bond; triple bond has one sigma and two pi bonds.

Hybridization of Orbitals

Concept and Application

Hybridization is the process by which atomic orbitals mix to form new hybrid orbitals suitable for bonding. This concept explains the observed shapes and bond angles in organic molecules.

• Carbon: Electronic configuration:

• sp3 hybridization: Tetrahedral geometry, as in methane ().

• sp2 hybridization: Trigonal planar geometry, as in ethene ().

• sp hybridization: Linear geometry, as in acetylene ().

Additional info: Hybridization helps explain the equivalence of bonds in organic molecules and the observed molecular geometries.