2. CHAPTER 4.1 THE DEVELOPMENT OF A NEW ATOMIC MODEL

3. PROPERTIES OF LIGHT  Light as a wave:  Visible light is a type of electromagnetic radiation, along with X-rays, ultraviolet and infrared light, microwaves, and radio waves.  These form the electromagnetic spectrum.  Waves have a repetitive nature and can be measured by wavelength( ) & frequency( ).  Wavelength unit is cm or nm.  Frequency unit is waves/sec or hertz (Hz).

4.  Frequency and wavelength are related to each other through the following equation: c = c = speed of light or 3.00 x 10 8 m/s As wavelength increases, frequency decreases and vice versa.

5. EXAMPLE  Determine the frequency ( ) of light whose wavelength ( ) is 6.87 x 10 -8 cm. c = c = 3.00 x 10 8 m/s = c/ Change cm to m and plug into equation = 3.00 x 10 8 m/s 6.87 x 10 -10 m =.437 x 10 18 ≈ 4.37 x 10 17 Hz

6.  Light as a particle:  Photoelectric effect is the emission of electrons from metal when light shines on it.  A quantum of energy is the minimum amount of energy that can be lost or gained by an atom.  A photon is a particle of electromagnetic radiation with no mass and carrying a quantum of energy.

7.  Max Plank and Einstein came up with a relationship between a quantum of energy and the frequency of radiation. E photon = h E = energy in Joules (J) h = Planck’s constant = 6.626 x 10 -34 Js = frequency in s -1 or Hz  In order for an electron to be ejected from a metal surface, it must be hit by a single photon possessing a certain minimum energy, which corresponds to a minimum frequency.

8. EXAMPLE  Determine the energy in Joules of a photon whose frequency is 4.85 x 10 15 Hz. E photon = h E = energy in Joules (J) h = Planck’s constant = 6.626 x 10 -34 Js = frequency in s -1 or Hz Plug into equation E photon = (6.626 x 10 -34 Js)  (4.85 x 10 15 Hz) E photon = 32.1361 x 10 -19 E photon = 3.21 x 10 -18 J

9. THE HYDROGEN-ATOM LINE-EMISSION SPECTRUM  When current is passed through a gas, it goes from the ground state to the excited state.  It emits light known as the emission-line spectrum.  When an excited hydrogen atom falls to its ground state, it emits a photon of radiation.

10. CH 4.2 NOTES THE QUANTUM MODEL OF THE ATOM

11. ELECTRONS AS WAVES  Behavior of electrons is similar to the behavior of waves.  Electron waves can only exist at certain frequencies.

12. HEISENBERG UNCERTAINTY PRINCIPLE  Werner Heisenberg had an idea on how to detect the location of electrons.  Heisenberg Uncertainty Principle: it is impossible to determine simultaneously both the position and velocity (speed) of an electron.

13. THE SCHRÖDINGER WAVE EQUATION  Developed an equation that treated electrons in atoms as waves and quantization of electron energies was an outcome of the equation.  Quantum Theory: mathematically describes the wave properties of electrons and other very small particles.

14. PRINCIPAL QUANTUM NUMBER N PRINCIPAL QUANTUM NUMBER ( N )  Main energy level (shell)  Size of the orbital  PERIOD #  Number of orbitals per main energy level is equal to n 2.  Number of electrons = 2n 2. Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem 1s1s 2s2s 3s3s

15. ANGULAR MOMENTUM QUANTUM NUMBER  Indicates the shape of the orbital, represented by l.  Values of l allowed are 0 and all positive integers less than or equal to n -1.  Ex. If n=2, shapes are l = 0 and l = 1  A sublevel consists of orbitals in a main energy level with the same value of l.

16. MAGNETIC QUANTUM NUMBER  Indicates the orientation of an orbital around the nucleus, represented by m.  Values of m are whole numbers, including zero, from –l to +l

17. SHAPES OF S, P, AND D-ORBITALS s orbital p orbitals d orbitals

18. MAXIMUM NUMBER OF ELECTRONS IN EACH SUBLEVEL Maximum Number of Electrons In Each Sublevel Maximum Number SublevelNumber of Orbitals of Electrons s 1 2 p 3 6 d 5 10 f 7 14 LeMay Jr, Beall, Robblee, Brower, Chemistry Connections to Our Changing World, 1996, page 146

19. Copyright © 2006 Pearson Benjamin Cummings. All rights reserved.

20. Electron capacities

27. Copyright © 2006 Pearson Benjamin Cummings. All rights reserved.

28. SPIN QUANTUM NUMBER  A single orbital has a maximum of two electrons, and those electrons must have opposite spins.  It has two values +1/2 and -1/2

29. CH 4.3 ELECTRON CONFIGURATIONS  Electron Configurations: the arrangement of electrons in an atom.  Each element has a unique electron configuration.  Ground-state Electron Configuration is the lowest energy arrangement of the electrons for an element.

30. GENERAL RULES Aufbau Principle  Electrons fill the lowest energy orbitals first.  “Lazy Tenant Rule”

31. GENERAL RULES  Hund’s Rule  Within a sublevel, place one electron per orbital before pairing them.  “Empty Bus Seat Rule” RIGHT WRONG Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

32. GENERAL RULES  Pauli Exclusion Principle  Each orbital can hold TWO electrons with opposite spins. Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem Wolfgang Pauli

33. ELECTRON FILLING IN PERIODIC TABLE 1 2 3 4 5 6 7 s d p s f

34.  really should include He, but He has the properties of the noble gases, and has a full outer level of electrons. ELEMENTS IN THE S - BLOCKS s2s2 s1s1 He

35. THE P-BLOCK p1p1 p2p2 p3p3 p4p4 p5p5 p6p6

36. TRANSITION METALS - D BLOCK d1d1 d2d2 d3d3 s1d5s1d5 d5d5 d6d6 d7d7 d8d8 s 1 d 10 d 10 Note the change in configuration.

37. F - BLOCK  Called the “inner transition elements”  Lanthanides and Actinides

39. PERIODIC PATTERNS  Example - Hydrogen s-block1st Period 1s 1 # element in block Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

40. NOTATION  Orbital Notation N 7e - Electron Configuration 1s 2 1s 2 2s 2 2s 2 2p 3 1s 2s 2p Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem N 14.0067 7

41. ELECTRON FILLING IN PERIODIC TABLE K4s1K4s1 Ca 4s 2 Sc 3d 1 Ti 3d 2 V3d3V3d3 Mn 3d 5 Fe 3d 6 Co 3d 7 Ni 3d 8 Cr 3d 4 Cu 3d 9 Zn 3d 10 Ga 4p 1 Ge 4p 2 As 4p 3 Se 4p 4 Br 4p 5 Kr 4p 6 1 2 3 4 s d p s Cr 4s 1 3d 5 Cu 4s 1 3d 10 4s3d Cr 4s 1 3d 5 4s3d Cu 4s 1 3d 10 Cr 3d 5 Cu 3d 10

42. ORDER IN WHICH SUBSHELLS ARE FILLED WITH ELECTRONS (FIG 19 PG 116) 1s2s3s4s5s6s7s1s2s3s4s5s6s7s 2p 3p 4p 5p 6p 7p 3d4d5d6d 3d4d5d6d 4f5f 4f5f 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d … 2 2 6 2 6 2 10 6 2 10

43. SHORTHAND ELECTRON CONFIGURATION Germanium  Example - Germanium [Ar]4s 2 3d 10 4p 2 Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem Ge 72.61 32

44. NOBLE GAS NOTATION  Longhand Configuration Shorthand Configuration S 16e - Valence Electrons Core Electrons S16e - [Ne] 3s 2 3p 4 1s 2 2s 2 2p 6 3s 2 3p 4 Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem S 32.066 16