Chemical Bonding

Chemical Bonds

  • Ionic bonds: held together by electrostatic interaction between ions of opposite charge; generally metal + nonmetal
  • Covalent bonds: sharing of electrons between two atoms; generally nonmetal
  • Metallic bonds: each atom bonded to several neighbors; electrons move freely which leads to good electrical and thermal conductivity

Lewis Structures

  • The electrons involved in bonding are the valence electrons
  • The number of valence electrons is the same as the group number
  • Lewis Symbol: the element’s chemical symbol + valence electron “dots”

The Octet Rule

  • Atoms try to obtain a noble gas configuration
  • Noble gases are very stable having high IE, low (less negative…near zero) EA
  • An octet has full s and p subshells
  • Significant number of exceptions to the octet rule; but, useful guideline
  • Exceptions to the octet rule:
    • Molecules and polyatomic ions containing an odd number of electrons
    • Molecules and polyatomic ions in which an atom has fewer than an octet of valence electrons. (especially B and Be)
    • Molecules and polyatomic ions in which an atom has more than an octet of valence electrons.

Ionic Compounds Characteristics

  • Brittle
  • High melting point
  • Crystalline
  • Cleave along planes
  • Form 3D lattice
  • Lattice formation is highly exothermic
    • The energy released by lattice formation more than makes up for the endothermic nature of ionization energies

Lattice Energy

  • Lattice energy: the energy required to completely separate one mole of a solid ionic compound into its gaseous ions.
  • The magnitude of the lattice energy depends on the electrostatic potential energy
  • Lattice energy increases as the charges on the ions increases •
  • Lattice energy increases as the radii of the ions decreases
  • The ionic charge factor is more important; ionic radii vary over a limited range

Covalent Bonding

  • Covalent bonds: atoms share electrons.
  • There are several electrostatic interactions in these bonds:
    • Attractions between electrons and nuclei
    • Repulsions between electrons
    • Repulsions between nuclei

Multiple Bonds

  • Single bond: shares one pair of electrons
  • Double bond: shares two pairs of electrons
  • Triple bond: shares three pairs of electrons
  • Bond length: distance between nuclei in a bond

Electronegativity

  • Electronegativity gives an estimate of whether a bond is nonpolar covalent, polar covalent or ionic.
  • Electronegativity: ability of an atom, in a molecule, to attract electrons to itself
  • Pauling scale of electronegativity-based on thermochemical data
  • Fluorine is the most electronegative element with a value of 4.0
  • Oxygen is next with a value of 3.5; N and Cl are 3.0
  • Cesium is the least electronegative element with a value of 0.7

Polar Covalent Bonds

  • When two atoms share electrons unequally, a bond dipole results.
  • Dipole moment: produced by two equal but opposite charges separated by a distance, r
    • It is reported in debyes (D).
    • If charge separation exists, the molecule is said to be a polar molecule
  • The bond moment can be represented as a vector
    • Vector: a quantity having both direction & magnitude
  • The greater the difference in electronegativity, the more polar is the bond.