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BUSINESS 1.EXAM 2THURSDAY NOVEMBER 4, 2010 2.MATERIAL COVERED: CHAPTERS 4, 5 & 6 3.TIME:7:00PM-8:00PM 4.WHERE:(TO BE ANNOUNCED LATER) 5.WHAT TO BRING:CALCULATOR,

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Presentation on theme: "BUSINESS 1.EXAM 2THURSDAY NOVEMBER 4, 2010 2.MATERIAL COVERED: CHAPTERS 4, 5 & 6 3.TIME:7:00PM-8:00PM 4.WHERE:(TO BE ANNOUNCED LATER) 5.WHAT TO BRING:CALCULATOR,"— Presentation transcript:

1 BUSINESS 1.EXAM 2THURSDAY NOVEMBER 4, MATERIAL COVERED: CHAPTERS 4, 5 & 6 3.TIME:7:00PM-8:00PM 4.WHERE:(TO BE ANNOUNCED LATER) 5.WHAT TO BRING:CALCULATOR, ONE PAGE OWN NOTES 6.CONFLICT IN SCHEDULE?CONTACT ME TO MAKE SEPARATE TIME

2 wavelength Visible light wavelength Ultraviolet radiation Amplitude Node Chapter 6: Electromagnetic Radiation

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5 Long wavelength --> small frequency low energy Short wavelength --> high frequency high energy

6 The electromagnetic spectrum.

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8 Which has the longest wavelength? 1.Infrared 2.Ultraviolet 3.X-rays 4.Radio waves

9 Rank the following in order of increasing frequency: microwaves radiowaves X-rays blue light red light UV light IR light

10 Waves have a frequency Use the Greek letter “nu”,, for frequency, and units are “cycles per sec” All radiation: = c c = velocity of light = 3.00 x 10 8 m/sec Long wavelength  small frequency Short wavelength  high frequency

11 –So, given the frequency of light, its wavelength can be calculated, or vice versa. The Wave Nature of Light The product of the frequency, (waves/sec) and the wavelength, (m/wave) would give the speed of the wave in m/s. –In a vacuum, the speed of light, c, is 3.00 x 10 8 m/s. Therefore,

12 The Wave Nature of Light What is the wavelength of yellow light with a frequency of 5.09 x s -1 ? (Note: s -1, commonly referred to as Hertz (Hz) is defined as “cycles or waves per second”.)

13 What is the frequency of violet light with a wavelength of 408 nm? The Wave Nature of Light

14 What is the wavelength of WONY? What is the wavelength of cell phone radiation? Frequency = 850 MHz What is the wavelength of a microwave oven? Frequency = 2.45 GHz

15 An anode (+) attracts e - Current is measured vacuum Anode (+) Metal cathode (-) window “Light” can cause ejection of e - from a metal surface. The Photoelectric Effect

16 Einstein proposed that “light”: is quantized. behaves like a stream of massless particles. photons –G. N. Lewis later named them photons. Imagine photons (balls) hitting e - embedded in glue. If the E of the ball: is low, it can’t eject an e -. exceeds the strength of the glue, an e - is released Higher intensity = more photons (balls). If E > threshold, more balls eject more e -. The Photoelectric Effect

17 Energy of radiation is proportional to frequency h = Planck’s constant = x Js Light acts as if it consists of particles called PHOTONS, with discrete energy. E = h Quantization of Energy

18 Relationships: E = h

19 Short wavelength light has: 1.High frequency and low energy 2.High frequency and high energy 3.Low frequency and low energy 4.Low frequency and high energy

20 Rank the following in order of increasing photon energy: microwaves radiowaves X-rays blue light red light UV light IR light

21 Energy of Radiation What is the frequency of UV light with a wavelength of 230 nm? What is the energy of 1 photon of UV light with wavelength = 230 nm?

22 What is the energy of a photon of 525 nm light? x J x J x J x J

23 Radio Wave Energy What is the energy of a photon corresponding to radio waves of frequency x 10 6 s -1 ?

24 What is the energy of a mole of 230 nm photons? Can this light break C-C bonds with an energy of 346 kJ/mol?

25 Does 1200 nm light have enough energy to break C-C bonds?

26 Where does light come from? Excited solids emit a continuous spectrum of light Excited gas-phase atoms emit only specific wavelengths of light (“lines”)

27 Light emitted by solids

28 Light emitted by hydrogen gas

29 The Bohr Model of Hydrogen Atom Light absorbed or emitted is from electrons moving between energy levels Only certain energies are observed Therefore, only certain energy levels exist –This is the Quanitization of energy levels

30 Excited atoms emit light of only certain wavelengths The wavelengths of emitted light depend on the element. Emission spectra of gaseous atoms

31 Line spectra of atoms

32 Energy Adsorption/Emission

33 For H, the energy levels correspond to: Constant = 2.18 x J Energy level diagram:

34 Each line corresponds to a transition: Example: n=3  n = 2

35 Balmer series Explanation of line spectra

36 Bohr Model of the Hydrogen Atom Heated solid objects emit continuous spectra. Excited atomic gases emit line spectra. Each element has a unique pattern.

37 Bohr (1913). The hydrogen e - : Orbits the nucleus. Different orbits are possible with different quantized E values: E = −R H n = 1, 2, 3...  1n21n2 Bohr Model of the Hydrogen Atom Rydberg constant x J Niels Bohr ground state If the e - has n = 1 (lowest, most negative E), the atom is in its ground state. If ionized (e - removed), n =  ( E = 0).

38 wavelength (nm) absorption:ΔE > 0, n ↑ emission: ΔE < 0, n ↓ Bohr’s model exactly predicts the H-atom spectrum. Bohr Model of the Hydrogen Atom Energy ultraviolet emission visible emission infrared emission ultraviolet absorption 0 -¹⁄ 9 R H -¼R H -RH-RH n∞321n∞321

39 Example Calculate the energy and wavelength (in nm) for an H- atom n = 4 → n = 2 transition. Bohr Model of the Hydrogen Atom 1nf21nf2 1ni21ni2 H-atom transitions: ΔE = −R H –

40 Calculate E and wavelength (nm) for an H-atom n = 4 → n = 2 transition. Bohr Model of the Hydrogen Atom


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