1. ELECTROMAGNETIC RADIATION

2. The Wave Nature of Light Much of our present understanding of the electronic structure of atoms has come from analysis of the light emitted or absorbed by substances

3. Electromagnetic Radiation Radiant energy which carries energy through space. All types of electromagnetic radiation move through a vacuum at a speed of 3.00 x 10 8 m/s

4. Wave-like Nature of Electromagnetic Radiation Electromagnetic radiation is measured in wavelenghts.

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

6. Since all electromagntic radiation travels at the same velocity in vacuum, c, its frequency,, is inversely proportional to its wavelength,   = c

7. Electromagnetic Radiation Waves have a frequencyWaves have a frequency Use the Greek letter “nu”,, for frequency, and units are “cycles per sec”Use the Greek letter “nu”,, for frequency, and units are “cycles per sec” All radiation: = cAll radiation: = c where c = velocity of light = 3.00 x 10 8 m/secwhere c = velocity of light = 3.00 x 10 8 m/sec Note that long wavelength = small frequencyNote that long wavelength = small frequency Short wavelength = high frequencyShort wavelength = high frequency

8. Electromagnetic Spectrum Indicates the wavelenghts of electromagnetic radiation characteristic of various regions of the electromagnetic spectrum

9. Electromagnetic Radiation Note that long wavelength = small frequency Short wavelength = high frequency increasing wavelength increasing frequency See Screen 7.4

10. Atomic Line Spectra and Niels Bohr Bohr’s greatest contribution to science was in building a simple model of the atom. It was based on an understanding of the SHARP LINE SPECTRA of excited atoms. Niels Bohr (1885-1962)

11. Bohr’s Model of the Hydrogen Atom Line Spectra Produced when gases are placed under reduced pressure in a tube and a high voltage is applied - colored lines, separated by black regions are produced

12. Line Spectra In 1885, Johann Balmer observed that the four lines of the hydrogen spectrum fit a formula

13. Visible lines in H atom spectrum are called the BALMER series. High E Short Short High High Low E Long Long Low Low Line Spectra of Excited Atoms

14. = C( 1/2 2 -1/n 2 ) n = 3,4,5,6 C = 3.29 x 10 15 s -1 Predicts the frequency of each line of the hydrogen line spectra

15. Bohr also assumed the electron could “jump” from one allowed energy state to another. Energy is absorbed when electron moves to a higher energy state. Energy is emitted when when electron moves from higher to a lower energy state

16. Orbital Energies E n = (-R H )(1/n 2 ) n = 1,2,3,4…. R H = Rydberg constant (2.18 x 10 -18 J) n = principle quantun number

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

18. Atomic Spectra and Bohr 1.Any orbit should be possible and so is any energy. 2.But a charged particle moving in an electric field should emit energy. One view of atomic structure in early 20th century was that an electron (e-) traveled about the nucleus in an orbit.

19. Bohr Model stated that electrons can only exist in certain discrete orbits — called stationary states. Each electron is restricted to QUANTIZED energy states. n = quantum no. = 1, 2, 3, 4,....

20. Atomic Spectra and Bohr Only orbits where n = integral no. are permitted.Only orbits where n = integral no. are permitted. Results can be used to explain atomic spectra.Results can be used to explain atomic spectra.

21. Atomic Spectra and Bohr If electrons are in quantized energy states, then E of states can have only certain values. If electrons are in quantized energy states, then  E of states can have only certain values. This explain sharp line spectra.

22. Calculate E for an electron “falling” from high energy level (n = 2) to low energy level (n = 1). Calculate  E for an electron “falling” from high energy level (n = 2) to low energy level (n = 1). E = E final - E initial = -C[(1/1 2 ) - (1/2) 2 ]  E = E final - E initial = -C[(1/1 2 ) - (1/2) 2 ] E = -(3/4)C  E = -(3/4)C

23. C has been found from experiment (and is now called R, the Rydberg constant) R (= C) = 1312 kJ/mol or 3.29 x 10 15 cycles/sec so, E of emitted light = (3/4)R = 2.47 x 10 15 sec -1 = (3/4)R = 2.47 x 10 15 sec -1 and l = c/n = 121.6 nm

24. Atomic Line Spectra and Niels Bohr Bohr’s theory was a great accomplishment. Rec’d Nobel Prize, 1922 Problems with theory — theory only successful for H.theory only successful for H. introduced quantum idea artificially.introduced quantum idea artificially. Niels Bohr (1885-1962)