1. Electrons in Atoms Chapter 5 General Chemistry

2. Objectives Understand that matter has properties of both particles and waves. Describe the electromagnetic spectrum in terms of wavelength and energy; identify regions of the electromagnetic spectrum.

3. Objectives Using Bohr’s model of the atom interpret changes (emission/absorption) in electron energies in the hydrogen atom corresponding to emission transitions between quantum levels. Write quantum numbers for electrons in atoms of elements Write the electron configurations for elements.

4. Review/Link to Previous Learning Atoms are the smallest part of an element that contains the properties of that element Atoms consist of protons, neutrons, and electrons Quantum Mechanics is the currently accepted model of the atom

5. Properties of Light

6. Electromagnetic Radiation

7. Electromagnetic Spectrum Electromagnetic (EM) radiation is a form of energy that exhibits wavelike behavior as it travels through space.

8. Electromagnetic Spectrum EM Spectrum: full range of frequencies of EM radiation Listed in order of increasing frequencies (decreasing wavelengths)

10. EM Spectrum & frequency Microwaves

11. Parts of EM Spectrum In order of increasing frequency: Radio waves Microwaves and Radar Infrared rays Visible light Ultraviolet rays X rays Gamma rays

12. Transverse Waves Visible light exhibits properties of wavelike motion. One type of wave is a transverse wave. Transverse waves are like taking a rope and moving it up and down. A transverse wave is a wave that causes the medium to vibrate at right angles to the direction that the wave travels. The particle in the medium travels perpendicular to the motion of the wave.

13. Trough

14. Parts of Transverse Wave Crest: highest point of the wave above the resting position Trough: lowest point of the wave below the resting position

15. Properties of Transverse Waves common properties of Transverse Waves: Frequency ( ν ) Wavelength ( λ) Speed (c)

16. Frequency Frequency ( ν) : the number of complete cycles of a wave passing a given point in a certain amount of time Measured in cycles per second or Hertz (Hz)

17. Wavelength Wavelength ( λ) is the distance between two adjacent crests of a wave. Measured in meters (m)

18. Calculate Wave Speed (Velocity) Speed of a any wave can be calculated by: v = speed (velocity) of wave (m/s) λ = wavelength (m) ν = frequency (Hz or cycles/s) v = λ x ν

19. Constant Speed of EM Waves All forms of EM radiation travel at a constant speed in a vacuum c = 3.00 x 10 8 m/sec (speed of light)

20. EM Radiation Speed Equation c = ν x λ c = 3.00 x 10 8 m/sec (speed of light) ν = frequency (Hz, 1 / sec, sec -1 ) λ = wavelength (meters)

21. Sample Problem Red light has a wavelength of 700. nm. What is the frequency of red light?

22. Sample Problem Red light has a wavelength of 700. nm. What is the frequency of red light? Answer: 4.29 x 10 14 Hz

23. Photoelectric Effect

24. Solar panel on calculator produced electrical current when light shines on it Photoelectric effect: emission of electrons from a metal caused by light striking the metal Experiments show only certain colors of light (only certain amount of energy) will allow electrical current to flow

25. Photoelectric Effect

26. Quanta of Energy A quantum is the minimum amount of energy that can be lost or gained by an atom Max Planck proposed relationship between quantum of energy and frequency of radiation (and color of light)

27. Energy in Quantum of Energy E = h x ν E = Energy (Joules) h = Planck’s Const. (6.63 x 10 -34 Joules x sec) ν = frequency (cycles/sec or Hertz or sec -1 )

28. Sample Problem Calculate the energy of a photon with a frequency of 2.85 x 10 12 Hz.

29. Sample Problem Calculate the energy of a photon with a frequency of 2.85 x 10 12 Hz. Answer: 1.89 x 10 -21 J

30. Study Buddy Review What is the Photoelectric Effect? How is a energy related to the color of light?

31. Energy and EM Radiation in Atoms

32. Spectra Observations of properties of light emitted by an atom after it absorbs extra energy Continuous spectrum Atomic emission spectrum

33. Line Spectra 1) For white light (sun or incandescent light bulb), you see continuous spectrum- rainbow of all colors Contains all colors of visible light Example: prism separates white light into rainbow colors. 2) Atomic Emission Spectrum (Line- emission spectrum) contains only certain colors or wavelengths, mostly black unique for each element “fingerprint” Example: pink glow from hydrogen gas tube that can be separated into series of lines

35. Recall Bohr’s Model of the Atom Thought atom was mostly empty space Nucleus in center is dense, positively charge Electrons move in orbits around the nucleus http://images.search.yahoo.com/search/imag es/

36. Hydrogen Atom Bohr’s model arose from hydrogen atomic emission spectra. If electrons move in specific orbits, they have a certain fixed amount of energy Observed as discrete bands of color

37. Why Bands of Color? Electrons can only absorb a certain amount of energy to enter the excited state (next energy level) Like rungs of ladder Can only move to specific orbit, which has a specific energy

38. Parts of Bohr model Ground state: lowest energy state of an atom Excited state: state in which an atom has a higher potential energy than it has in its ground state

39. How is Radiation Produced from Atom? When an excited electron falls back from an excited state to its ground state, it emits a photon of radiation. (E= h ν ) Since only specific frequencies of light are emitted from elements, the energy levels in atom are fixed. Electrons exist in only certain energy states.

41. Quantum Mechanics

43. Only for Hydrogen Bohr’s model of atom was successful only for elements with one electron

44. Heisenberg Continued building on Bohr model Uncertainty Principle – it is impossible to know both the exact position and the velocity of an electron or other particle simultaneously Act of measuring changes what you are trying to find Thus, no well-defined orbit (like Bohr had proposed) Best we can do is represent PROBABILITY of finding e- within a given space

45. Schrodinger Erwin Schrodinger developed a mathematical equation that allowed for wavelike behavior of electrons Energy of electrons is quantized

46. Quantum Mechanical Model of Atom atom is mostly empty space Nucleus in center is dense, positively charge Electrons are around the nucleus e- do not have a precise orbit (electron cloud) e- moves in wavelike motion with quantized amounts of energy

47. Quantum Numbers Quantum Numbers: specify the properties of atomic orbitals and properties of electrons in orbitals.

48. Principle Quantum Number (n) n: Principal energy level or energy number n = 1, 2, 3…(as you get farther from nucleus) maximum # of e- in energy level = 2n 2

49. Angular Momentum (l) l: Shape of orbitals (sublevels) named s, p, d, f s = spherical p = dumbbells d = clover leaf f = butterflies?? each sublevel has slightly different energy

50. Magnetic (m l ) m l = Orientation of orbital around nucleus s = 1 orbital p = 3 orbitals d = 5 orbitals f = 7 orbitals

51. Spin m s m s = Magnetic spin Only two directions of spin to create magnetic field Like N and S of magnet

52. Electron Configurations Electron configuration: Notation describing the most stable arrangement of e- around the nucleus Like writing the address for electrons around nucleus of atom

54. Three Rules for Electron Configurations 1. The Aufbau Principle e- added one at a time to the lowest energy levels 2. The Pauli Exclusion Principle orbital can hold at most 2 e- must have opposite spin 3. Hund’s Rule (The “bus seat” rule) e- like to be unpaired if possible &e- enter equal energy orbitals until all orbitals contain one e- with parallel spin, then they begin to pair up

55. Electron Configuration Write the electron configuration for the following: Oxygen Sodium copper