Bonding: General Concepts Continued

Covalent Bonding Model Remember chemical bonds can be viewed as forces that cause a group of atoms to behave as a unit. Bonds result from the tendency of a system to seek its lowest possible energy. Bonds occur when collections of atoms are more stable (lower in energy) than the separate atoms.

Covalent Bond energies and Chemical Reactions Bond Energies Bond energy: the energy required to break a given chemical bond. To break bonds, energy must be added to the system (endothermic). To form bonds, energy must be released (exothermic).

Single bond < Double bond < Triple bond one pair of electrons shared. Double bond: two pairs of electrons shared. Triple bond: three pairs of electrons shared. Bond Energies Single bond < Double bond < Triple bond

Shared Electron Pairs and Bond Length As the number of shared electrons increases, the bond length shortens.

Bond Energy and Enthalpy Bond energy values can be used to calculate approximate energies for reactions. Example: calculate the change in energy that accompanies the following reaction: H2 (g) + F2 (g) → 2HF (g) To form HF, one H-H bond and one F-F bond must be broken and two H-F bonds must be formed.

Remember for bonds to be broken energy must be added to the system – an endothermic process – and carries a positive sign. Formation of a bond releases energy – an exothermic process – and carries a negative sign. Enthalpy change: H = n×D(bonds broken) – n×D(bonds formed) where  represents the sum of terms and D represents the bond energy per mole (n) of bonds. D always has a positive sign.

In the case of the formation of HF, ∆H = DH-H + DF-F – 2DH-F = (1 mol x 432 kJ/mol) + (1 mol x 154 kJ/mol) - (2 mol x 565 kJ/mol) = -544 kJ Thus, when 1 mol H2 (g) and 1 mol F2 (g) react to form 2 mol HF (g), 544 kJ of energy should be released. When this result is compared to the result for the reaction when using the standard enthalpy of formation for HF (-542 kJ) the use of bond energies works well.

The Covalent Chemical Bond: A Model Models are attempts to explain how nature operates on the microscopic level based on experiences in the macroscopic world.

Fundamental Properties of Models A model does not equal reality. Models are oversimplifications, and are therefore often wrong. Models become more complicated and are modified as they age. We must understand the underlying assumptions in a model so that we don’t misuse it. When a model is wrong, we often learn much more than when it is right.

Localized Electron Bonding Model A molecule is composed of atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms. Electron pairs are assumed to be localized on a particular atom or in the space between two atoms: Lone pairs – pairs of electrons localized on an atom Bonding pairs – pairs of electrons found in the space between the atoms

Localized Electron Bonding Model has three parts: Description of valence electron arrangement (Lewis structure). Prediction of geometry (VSEPR model). Description of atomic orbital types used to share electrons or hold lone pairs.

Lewis Structure G. N. Lewis (1875-1946) Shows how valence electrons are arranged among atoms in a molecule. Reflects central idea that stability of a compound relates to noble gas electron configuration.

Duet Rule Hydrogen forms stable molecules where it shares two electrons.

Octet Rule Elements form stable molecules when surrounded by eight electrons.

Steps for Writing Lewis Structures Sum the valence electrons from all the atoms. Use a pair of electrons to form a bond between each pair of bound atoms. Atoms usually have noble gas configurations. Arrange the remaining electrons to satisfy the octet rule (or duet rule for hydrogen).

Steps for Writing Lewis Structures Sum the valence electrons from all the atoms. (Use the periodic table.) Example: H2O 2 (1 e–) + 6 e– = 8 e– total Use a pair of electrons to form a bond between each pair of bound atoms. Example: H2O

Atoms usually have noble gas configurations Atoms usually have noble gas configurations. Arrange the remaining electrons to satisfy the octet rule (or duet rule for hydrogen). Examples: H2O, PBr3, and HCN