Atomic Emission Spectra

2. Atomic Emission Spectra • Zumdahl2: p. 290-299 Atoms Let go A range

3. What is light? • White light: reflection of all colors • Black light: absorption of all colors • Colors are each a different wavelength (λ: lamda) of light

4. Colors • Different wavelengths of light are seen as different colors. • Different colors indicate (show) different energy levels.

5. c=fλ (velocity of light = frequency x wavelength) • the greater the frequency the shorter the wavelength • ΔE = hf • (energy lost by the electron = h(constant) x frequency • Frequency (and thus, color) of the light depends on the amount of energy lost by the electron.

6. c=fλ

8. When atoms are “exited” (energy is added) they produce light. • Not white or all-colored light, but one color at a time.

9. Study the light emitted (produced) by atoms and ions to deduce (find out) the structure of atoms. • When an atom is “excited” its electrons gain energy and move to a higher energy level. To return to a lower energy level, electrons must lose energy. They do this by giving off light.

11. Continuous spectrum: all wavelengths of visible light contained in white light. • Light emitted by an atom can be separated into a line spectrum that shows exactly what frequencies of light are present.

13. Increasing frequency Further from nucleus Closer to nucleus

14. Because the light emitted from atoms is a line spectrum (not a continuous spectrum) we determine that: • There are “discrete” (separate) energy levels for each atom that can only produce light of certain wavelengths (this is NOT ordinary white light!).

17. Increasing frequency (f) (increasing energy)

18. Increasing frequency (f) (increasing energy)

20. Hydrogen • Only certain energy levels can occur (not a continuous spectrum)

21. Energy Level Diagram • The larger the difference in energy, the greater the frequency (thus, the more purple the light).

22. Increasing potential energy Visible Frequency

23. convergence:the lines in a spectrum converge (get closer together) as frequency increases. • related to how much energy is required to remove the electron from the atom (ionize)

24. Increasing frequency Further from nucleus Closer to nucleus

25. Stop

26. Electronic Structure Energy Levels Shells

27. Most stable = closest to nucleus

28. 1st energy level = 2 • 2nd energy level = 8 • Electronic structure: number of electrons in each orbital

30. H=1 • O=2,6 (two electrons in the first energy level, six in the second) • Al=2,8,3 • Cl= • Ca= • Different isotopes have the same electronic structure and the same chemical properties!

31. Electron Behavior • Valence shell: outer shell of an atom • determine the physical and chemical properties of an atom Valence Shell

33. How many electrons in valence shell? • Al • Ne • Li • Ca

34. Stop here 

35. HL Topic Electronic Structure of Atoms Zumdahl2: p. 307-312

36. Electronic Structure • Energy levels A. Sub-levels 1. Orbital a. Spin

37. Energy Levels • Major shells (layers) around the nucleus • filled before higher levels are filled • 1st: 2 electrons • 2nd: 8 electrons • 3rd: 8 electrons

38. Sub-levels • Different shapes • s – sphere • one orbital • p – figure eight • three orbitals • d – • five orbitals • f – • seven orbitals

40. p Sub-level • p sub-level has three orbitals • px, py, pz

41. d and f sub-levels have very complex shapes

42. Orbitals • Each orbital can hold two electrons. • Electrons spin in opposite directions

44. Energies of sub-levels

47. Electronic structure of atoms

48. Depending on where an electron was around the nucleus, it had different energy states. Ground state: An orbit near the nucleus: not very exited at all. Excited State: An orbit farther away from the nucleus: much more potential for giving off energy. Energy States

49. If electrons are both waves and particles, where are they around the atom? It is impossible to figure out both the position and velocity of an electron, at the same time. We CAN figure out the probability that an electron is in any one spot at any given time. Heisenberg Uncertainty Principle

50. Electron Configuration • The arrangement of electrons in an atom • Each element’s atoms are different • Arrangement with the lowest energy= ground state electron configuration • How do we figure out what the ground state electron configuration looks like?