1. Unit 4 Energy and the Quantum Theory

2. I.Radiant Energy Light – electrons are understood by comparing to light 1. radiant energy 2. travels through space 3. makes you feel warm Light has properties of waves and particles

3. Waves Electromagnetic waves – what light is Electromagnetic radiation 1. light 2. x-rays 3. gamma rays 4. radio waves

4. Waves Amplitude: height of a wave (intensity) Wavelength: distance between crests (λ) Frequency: number of wavelengths passing a point in a given amount of time (cycles/ sec., Hz) (ν) Speed: light ( 3.00 X 10 8 m/s = c)

5. Waves

7. C = νλ λ = 633 nm (red) What is the frequency? 633nm x 1m/1x10 9 nm= 6.33x10 -7 m ν = 3.00x10 8 m/s/6.33x10 -7 m = 4.74x10 14 Hz

8. Electromagnetic Spectrum

9. II. Quantum Theory Planck – proposed that a fixed amount of energy can be absorbed or emitted by atoms, and called the energy a quantum E = hν h = 6.63 x 10 -34 J-s Planck’s constant

10. Figure 11.14: Continuous and discrete energy levels.

11. Figure 11.15: The difference between continuous and quantized energy levels.

12. Photoelectric Effect Light causes electrons to shoot off metals – not any kind of light will do Example: with Na red light never works but violet light always works

13. Figure 11.6: Photons of red and blue light.

14. Einstein and the photon Photon: particle of light electron can either use the photon or it can’t; no partial use is allowed Energy of a photon is high when the frequency is high higher frequency is more dangerous low frequency is not dangerous

15. Dual nature of light Light behaves as both a wave and a particle

16. III. More About the Atom Line spectrum: contains only certain colors or wavelengths Atomic emission spectrum: a line spectrum for the elements; fingerprint of elements, unique for each element

17. Figure 11.11: Colors and wavelengths of photons in the visible region.

18. Bohr’s model Electrons have certain orbits that correspond to quanta of energy Quantum number (n): energy level Ground state (n=1): lowest energy level, closest to nucleus Excited state (n>1): higher energy level, further from nucleus

19. Figure 11.17: The Bohr model of the hydrogen atom.

20. Bohr’s model When an electron absorbs energy from light, it jumps to a higher energy level. When it falls back to the ground state, it releases the energy as light.

21. Figure 11.8: An excited lithium atom emitting a photon of red light to drop to a lower energy state.

22. Figure 11.13: Each photon emitted corresponds to a particular energy change.

23. Matter Waves DeBroglie proposed the idea of matter waves. Since light behaves like a wave and particle, matter should also act like a wave and particle. There is a dual nature to matter. Electron microscopes work because matter behaves like a wave.

24. Heisenberg’s Uncertainty Principle The position and momentum of a moving object (electron) cannot be measured and known at exactly the same time. We can only tell the probability of finding it in a certain area.

25. IV. New Approach to Atom Quantum-mechanical model – uses probability to describe electron motion Electron density – cloud that shows the probability of finding an electron in a certain area

26. Figure 11.19: The orbital that describes the hydrogen electron in its lowest possible energy state.

27. Orbitals Area where electron is most likely to be found (90%) s – spherical p – dumbbell d – four leaf clover f - complex

28. Figure 11.20: The hydrogen 1s orbital.

29. Figure 11.25: The three 2p orbitals.

30. Figure 11.28: The shapes and labels of the five 3d orbitals.

31. Orbitals & Energy Principal energy levels – described by quantum number (n) Sublevels n=1 1 sublevel 1s n=2 2 sublevels 2s, 2p n=3 3 sublevels 3s, 3p, 3d n=4 4 sublevels 4s, 4p, 4d, 4f 2s is bigger than 1s but same shape

32. Figure 11.24: The relative sizes of the 1s and 2s orbitals of hydrogen.

33. Figure 11.27: Relative sizes of the spherical 1s, 2s, and 3s orbitals of hydrogen.

34. Orbitals & Energy Sublevel orbitals s 1 p 3 d 5 f 7

35. Electron spin A spinning charge produces a magnetic field, so a spinning electron is like a little magnet. Opposite spins will cancel each other. Only two electrons can fit in an orbital. They are called “paired electrons”.

37. Maximum number of electrons in sublevels s 2 p 6 d 10 f 14

38. V. Electron Configurations Show where electrons are located in the atom and how much energy they possess.

39. The Aufbau Principle Electrons are added one at a time to the lowest energy orbitals available until all electrons have been accounted for. Examples: C has 6 electrons Cl has 17 electrons

41. The Pauli Exclusion Principle There can only be two electrons per orbital; they must have opposite spins. ( paired & unpaired electrons)

42. Hund’s Rule Electrons fill up orbitals in the same sublevel with one electron before they pair up.

43. Orbital diagram

44. Orbital diagram for Oxygen