Atomic Structure & Light: Waves, Particles, and Quantum Mechanics

1. Chapter 7 Atomic Structure

2. Light • Made up of electromagnetic radiation • Waves of electric and magnetic fields at right angles to each other.

3. Parts of a wave Wavelength (units???) l Frequency (ν) = number of cycles in one second (units???) Measured in hertz(Hz) 1 hertz = 1 cycle/second (or s-1)

4. Frequency = n

5. Kinds of EM waves • There are many • different l and n • Radio waves, microwaves, x rays and gamma rays are all examples • Light is only the part our eyes can detect • l → ← n G a m m a R a d i o w a v e s R a y s

6. The speed of light • in a vacuum is 2.998 x 108 m/s. • = c. • c = ln =2.998 x 108 m/s • What are the units for l? • For n? • Calculate the wavelength of light with a frequency 5.89 x 105 Hz. • What is the frequency of blue light with a wavelength of 484 nm?

7. In 1900: • Matter and energy were seen as different from each other in fundamental ways: • Matter was particles • Energy could come in waves, with any frequency. • Max Planck found that the cooling of hot objects couldn’t be explained by viewing energy as a wave.

8. Energy is Quantized • Planck found DE came in chunks with size hn, so • DE = nhν • where n is an integer. • and h is Planck’s constant (h = 6.626 x 10-34 J·s) • these packets of hν are called quantum.

9. Einstein is next • Said electromagnetic radiation is quantized in particles called photons. • Each photon has energy = hν = hc/l • Combine this with E = mc2, and you get the apparent mass of a photon: • m = h / (lc).

10. So which is it? • Is energy a wave like light, or a particle? • Yes • This concept is called the Wave -Particle duality. • What about the other way, is matter a wave? • Yes • Who was the first scientist to specify that electrons are both particles (matter) and waves? • Werner Heisenberg

11. Matter as a wave • Using the velocity (v) instead of the frequency (ν) we get • De Broglie’s equation l = h/mv • Using this, we can calculate the wavelength of an object.

12. Examples • The laser light of a CD is 7.80 x 102 m. What is the frequency of this light? • What is the energy of a photon of this light? • What is the apparent mass of a photon of this light? • What is the energy of a mole of these photons?

13. What is the wavelength? • of an electron with a mass of 9.11 x 10-31kg traveling at 1.0 x 107m/s? • of a softball with a mass of 0.10 kg moving at 125 mi/hr?

14. How do they know? • When light passes through, or reflects off, a series of thinly spaced lines, it creates a rainbow effect because the waves interfere with each other.

15. A wave moves toward a slit.

16. Comes out as a curve

17. with two holes

18. Two Curves with two holes

19. Two Curves with two holes Interfere with each other

20. Two Curves with two holes Interfere with each other crests add up

21. Several waves

22. Several waves Several Curves

23. Several waves Several waves Several Curves Interference Pattern

24. What will an electron do? • It has mass, so it is matter. • A particle can only go through one hole • A wave goes through both holes • We know light shows interference patterns. So what about electrons?

25. Electron as Particle Electron “gun”

26. Electron as wave Electron “gun”

27. Which did it do when tested? • It made the diffraction pattern, so the electron is a wave. • Led to Schrödingers equation

28. What will an electron do? • An electron does go though both slits, and makes an interference pattern. • It behaves like a wave. • Other matter has wavelengths too short to notice.

29. Spectrum • Def: the range of frequencies present in light. • White light has a continuous spectrum, creating a rainbow. • All the colors are possible.

30. Hydrogen spectrum • Called emission spectrum because these are the colors it gives off or emits. • Also called a line spectrum, because there are just a few discrete lines showing 656 nm 434 nm 410 nm 486 nm Spectrum

31. What this means • Only certain energies are allowed for the hydrogen atom (can only give off certain energies). • Use DE = hn = hc / l • Remember, energy in the atom is quantized.

32. Niels Bohr • Developed the quantum model of the hydrogen atom. • He said the atom was like a solar system (“planetary model”). • The electrons were attracted to the nucleus because of opposite charges, but didn’t fall into the nucleus because they were moving around it in their assigned orbitals (or energy levels.)

33. The Bohr Ring Atom • He didn’t know why but only certain energies were allowed. • He called these allowed energies energy levels. • Putting energy into the atom moved the electron away from the nucleus, from ground state to excited state. • When it returns to ground state it gives off light of a certain energy.

34. The Bohr Ring Atom n = 4 n = 3 n = 2 n = 1

35. The Bohr Model • n is the energy level (in integers). • Z is the nuclear charge, which is +1 for hydrogen. • E = -2.178 x 10-18J (Z2 / n2 ) • n = 1 is called the ground state • when the electron is removed, n = ¥ • E = 0

36. The Bohr Model • When the electron moves from one energy level to another • ΔE = Efinal - Einitial • ΔE = -2.178 x 10-18 J · (Z2 / nf2 - Z2 / ni2) • ΔE = -2.178 x 10-18 J · Z2 (1/ nf2 - 1/ ni2)

37. Examples • Calculate the energy need to move an electron from its ground state to the third energy level. • Calculate the energy released when an electron moves from n = 4 to n = 2 in a hydrogen atom. • Calculate the energy released when an electron moves from n = 5 to n = 3 in a He+1 ion.

38. When is it true? • Only for hydrogen atoms and other monoelectronic species(????). • Why the negative sign? • To increase the energy of the electron you move it further from the nucleus. (Would this be considered kinetic or potential energy?) • the maximum energy an electron can have is zero, at an infinite distance.

39. The Bohr Model: • doesn’t work. • only works for hydrogen atoms. • WHY? Electrons don’t move in circles! • The quantization of energy is right, but not because they are circling like planets.