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Lecture 21: Chemical Bonding

Lecture 21: Chemical Bonding. Reading: Zumdahl 13.1-13.3 Outline Types of Chemical Bonds Electronegativity Bond Polarity and Dipole Moments. Chemical Bonds.

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Lecture 21: Chemical Bonding

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  1. Lecture 21: Chemical Bonding • Reading: Zumdahl 13.1-13.3 • Outline • Types of Chemical Bonds • Electronegativity • Bond Polarity and Dipole Moments

  2. Chemical Bonds • In broad terms, a chemical bond is a term used to characterize an interaction between two atoms that results in a reduction in Gibbs free energy for the system relative to the isolated atoms. • The degree of energy reduction or “stabilization” is given by the energy required to break the bond (known as the “bond energy”)

  3. Chemical Bonds (cont.) Stabilization

  4. Chemical Bonds (cont.) • Prototypical bonding example: NaCl (Sodium Chloride) • When NaCl is heated to the point of melting, one can demonstrate that the resulting solution conducts electricity. • This observation demonstrates that the solution (liquid NaCl) contains charged species. Those species are Na+ and Cl-.

  5. Chemical Bonds (cont.) • Why the formation of Na+ and Cl-? • In short, Na+ and Cl- are more energetically stable than atomic Na and Cl. • With the transfer of an electron from Na to Cl, two ions of opposite charge are produced. The Coulombic attraction between these ions is largely responsible for the stabilization.

  6. Chemical Bonds (cont.) • What is the extent of Coulombic stabilization? +1 -1 Na+ and Cl- x Na 0.276 nm -504 kJ/mol

  7. Chemical Bonds (cont.) • This example was for a collection of NaCl dimers in the gas phase….solid NaCl is a bit different. • Each ion is surrounded by six ions of the opposite charge.

  8. Chemical Bonds (cont.) • NaCl is an example of “ionic bonding”. In this case an amount of charge approaching that for an electron is “transferred” from one atom to its bonding partner. • Ionic bonding is one limit in the spectrum of bonding. The second limit is a bond in which electrons are “shared” rather than being transferred. • Electrons are shared in a covalent bond.

  9. Chemical Bonds (cont.) • The H-H example we saw previously is an example of covalent bonding.

  10. Chemical Bonds (cont.) • If ionic bonding is one limit, and covalent bonding is the other limit, what lies in the middle? • Polar covalent bonds: bonds in which electrons are shared, but the probability distribution is skewed toward the atom with greater electron affinity. • HF is a typical example.

  11. Electronegativity • From the previous examples of bonding, we need some way to predict what pattern of bonding we would expect. • Clearly, the affinity of an atom for electrons is a critical parameter. • Electronegativity: The ability of an atom in a molecule to attract shared electrons to itself.

  12. Electronegativity • Electronegativity can be defined in many ways. Pauling model is the most widely used. • Idea: compare the bond energy of an “HX” molecule to that of the average of an HH bond and an XX bond: Expected energy = [(H-H energy)(X-X energy)]1/2 D = (H-X)experimental - (H-X)expected D > 0: ionic character D = 0: covalent

  13. Electronegativity (cont.) • Pauling used this approach to develop a scale, where F = 4.0 (flourine has largest electronegativity). F = 4 O = 3.4 Cl = 3.2 C = 2.6 H = 2.2 Na = 0.9

  14. Electronegativity (cont.) • The key idea is this: the greater the electronegativity difference between two atoms, the more ionic the bond. • Example: Which of the following compounds is expected to demonstrate intermediate bonding behavior (i.e., polar covalent). Cl-Cl O-H Na-Cl 0 2.3 1.2 Delect

  15. Which atomic pair will for the most ionic bond?

  16. Dipole Moments • The above discussion involved bonds in which electrons were shared, but shared unequally in polar, covalent bonds • In the HF example, when placed in an electric field the HF atoms will align. • This observation demonstrates that the centers of negative and positive charge do not coincide.

  17. Dipole Moments

  18. Dipole Moments (cont.) • When the centers of negative and positive charge are separated, we say that the molecule has a dipole moment.

  19. Dipole Moments (cont.) • The dipole moment (m) is defined as: m = QR Charge magnitude Separation distance R + center

  20. Dipole Moments (cont.) • The units of dipole moment are generally the Debye (D): 1 D = 3.336 x 10-30 C.m • Example, the dipole moment of HF is 1.83 D. What would it be if HF formed an ionic bond (bond length = 92 pm)? m = (1.6 x 10-19 C)(9.2 x 10-11 m) = 1.5 x 10-29 C.m x (1D/3.336 x 10-30 C.m) = 4.4 D

  21. Dipole Moments (cont.) • Molecular gemoetry is a critical factor in determining if a molecule has a dipole moment:

  22. Which compound is polar?

  23. Dipole Moments (cont.) • Molecular gemoetry is a critical factor in determining if a molecule has a dipole moment: No net dipole moment. Dipoles add as vectors!

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