As a bond forms each atom shares one electron with

Info icon This preview shows pages 63–66. Sign up to view the full content.

View Full Document Right Arrow Icon
As a bond forms, each atom shares one electron with the other atom, but both electrons of the bond count towards each atom’s octet. Bonds formed in this way are stabilizing for two reasons. Firstly, each electron in a bond is simultaneously attracted to both nuclei of the bond, whereas before a bond was formed, each electron was only attracted to its own nucleus. Secondly, because bonded electrons may occupy a larger volume of space than non- bonded electrons, the energies of bonded atoms are lower than non-bonded atoms. A chemical bond may, from another point of view, form via the overlap of atomic orbitals, each of which contains a single electron. When two atoms are infinitely far apart, as shown in the diagram below, the energy of the two atoms is 0 J. As they approach each other, the two orbitals begin to overlap and the total potential energy decreases. With decreasing distance between the atoms, the degree of overlap between the two orbitals increases, the incipient chemical bond increases in strength, and the total potential energy decreases. If the atoms approach each other too closely, the potential energy rises due to repulsion between the positively charged nuclei. At the internuclear distance of minimum energy, the attractive and repulsive forces effectively balance and the chemical bond has its greatest possible strength. The negative of the potential energy at this point is the bond dissociation energy, the energy that the bond must absorb to separate the atoms to infinite distance. The internuclear distance at this minimum potential energy is called the bond length.
Image of page 63

Info icon This preview has intentionally blurred sections. Sign up to view the full version.

View Full Document Right Arrow Icon
Chemistry 132 Lab Manual Page 64 (Figure 12.59 from Zumdahl and DeCoste) The internuclear distance of two bonded atoms is not fixed. In a diatomic molecule, the two atoms vibrate along the internuclear axis. When atoms in a polyatomic molecule vibrate, several characteristic vibrations are possible. These vibrations are called normal modes of vibration. Shown below from left to right are the asymmetric stretching, symmetric stretching, and bending modes of vibration for water. (The next two diagrams were taken from Chapter 17 in Physical Chemistry for the Chemical and Biological Sciences by Raymond Chang, 2000, University Science Books). Each of these normal modes of vibration oscillates at a characteristic frequency, n , which is shown below the mode of vibration. The frequencies, n , are given in units of wavenumbers. A wavenumber, n , equals 1/ l where l is the wavelength in cm. You will remember that n = c/ l , where c is the speed of light. Therefore, wavenumbers are proportional to frequency and c is the proportionality constant. When frequency is expressed in wavenumbers, n represents how many waves fit within one cm.
Image of page 64
Chemistry 132 Lab Manual Page 65 When vibrating atoms of a molecule interact with light whose frequency is the same as the characteristic frequency, n , the molecule absorbs the energy of the light. This interaction does not change the frequency of vibration, but the amplitude of the vibrations increases. That is, the normal mode of vibration does not change, but more
Image of page 65

Info icon This preview has intentionally blurred sections. Sign up to view the full version.

View Full Document Right Arrow Icon
Image of page 66
This is the end of the preview. Sign up to access the rest of the document.

{[ snackBarMessage ]}

What students are saying

  • Left Quote Icon

    As a current student on this bumpy collegiate pathway, I stumbled upon Course Hero, where I can find study resources for nearly all my courses, get online help from tutors 24/7, and even share my old projects, papers, and lecture notes with other students.

    Student Picture

    Kiran Temple University Fox School of Business ‘17, Course Hero Intern

  • Left Quote Icon

    I cannot even describe how much Course Hero helped me this summer. It’s truly become something I can always rely on and help me. In the end, I was not only able to survive summer classes, but I was able to thrive thanks to Course Hero.

    Student Picture

    Dana University of Pennsylvania ‘17, Course Hero Intern

  • Left Quote Icon

    The ability to access any university’s resources through Course Hero proved invaluable in my case. I was behind on Tulane coursework and actually used UCLA’s materials to help me move forward and get everything together on time.

    Student Picture

    Jill Tulane University ‘16, Course Hero Intern