Unit III

1. Unit III Periodicity and Introduction to Bonding

2. Chapter 6: The Structure of the Atom • To understand the electronic structure of atoms, one must understand electromagnetic radiation • Electromagnetic Radiation • Propagating electromagnetic wave that carries energy through space • Characterized by it’s wavelength and frequency

3. Electromagnetic Radiation wavelength Visible light Amplitude wavelength Node Ultaviolet radiation

4. Waves The distance between corresponding points (successive crests) on adjacent waves is the wavelength (). The number of waves passing a given point per unit of time (usually per seconds) is the frequency(). 1 second -1 = 1/s = Hertz For waves traveling at the same velocity, the longer the wavelength, the smaller the frequency.

5. Speed of Waves • In a vacuum, velocity = 2.9979 x 108 m/s • Speed of light (c) = 2.9979 x 108 m/s

6. Speed of Waves c=   • = frequency (s1, Hz, cyc/s, or waves/s) • = wavelength (m) • c= speed of light (m/s)

7. Quantization of Energy The wave nature of light does not explain how an object can glow when its temperature increases.

8. Max Planck • Max Planck explained it by assuming that an object can gain or lose energy by absorbing/emitting energy in packets called quanta.

9. Planck’s Constant • Transfer of energy is quantized, and can only occur in discrete units, called quanta. • E = change in energy, in J • h = Planck’s constant 6.626  1034 J s •  = frequency, in s1 •  = wavelength, in m

10. Einstein • Einstein used this assumption to explain the photoelectric effect. • He concluded that energy is proportional to frequency: E = h where h is Planck’s constant, 6.626  10−34 J-s

11. The Nature of Energy • Therefore, if one knows the wavelength of light, one can calculate the energy in one photon, or packet, of that light: c =  E = h

12. Math c =  E = h E = hc/ How much energy does a quanta with a frequency of 5.09x1014s-1 possess? E = h E= (6.626  10−34 J●s)(5.09x1014s-1) E= 3.37  10−19 J

13. Practice Exercises • Exercise 6.2 Compare the energy of a mole of photons of orange light (625nm) with the energy of a mole of photons of microwave radiation having a frequency of 2.45GHz (1GHz = 109s-1). Which has the greater energy?

14. The Nature of Energy • Another mystery in the early 20th century involved the emission spectra observed from energy emitted by atoms and molecules.

15. Line Spectra of Excited Atoms • Excited atoms emit light of only certain wavelengths • The wavelengths of emitted light depend on the element.

16. Continuous vs.Line Spectra • Continuous spectrum:Contains allthe wavelengths of light. • Line (discrete) spectrum: Contains only some of the wavelengths of light.

17. Line Spectra

18. The Nature of Energy • Niels Bohr adopted Planck’s assumption and explained these phenomena in this way: • Electrons in an atom can only occupy certain orbits (corresponding to certain energies). • Electrons in permitted orbits have specific, “allowed” energies; these energies will not be radiated from the atom. • Energy is only absorbed or emitted in such a way as to move an electron from one “allowed” energy state to another; the energy is defined by E = h

19. ( ) - 1 nf2 E = −RH 1 ni2 The Nature of Energy The energy absorbed or emitted from the process of electron promotion or demotion can be calculated by the equation: where RH is the Rydberg constant, 2.18  10−18 J, and ni and nf are the initial and final energy levels of the electron.