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Bonding: General Concepts Continued

Understand the fundamental concepts of chemical bonding, bond energies, and covalent bonds. Learn about shared electron pairs, bond length, enthalpy change, and the Lewis electron bonding model.

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Bonding: General Concepts Continued

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  1. Bonding: General Concepts Continued

  2. 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.

  3. 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).

  4. Single 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

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

  6. 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.

  7. 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(bondsbroken) – n×D(bondsformed) • where  represents the sum of terms and D represents the bond energy per mole (n) of bonds. • D always has a positive sign.

  8. 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.

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

  10. 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.

  11. 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

  12. 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.

  13. 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.

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

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

  16. 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).

  17. 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

  18. 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

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