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3-Atomic Structure Overview Characteristics of Atoms Interaction b/tw matter and light –Photoelectric Effect Absorption and Emission Spectra Electron behavior.

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Presentation on theme: "3-Atomic Structure Overview Characteristics of Atoms Interaction b/tw matter and light –Photoelectric Effect Absorption and Emission Spectra Electron behavior."— Presentation transcript:

1 3-Atomic Structure Overview Characteristics of Atoms Interaction b/tw matter and light –Photoelectric Effect Absorption and Emission Spectra Electron behavior Quantum numbers

2 Atomic Structure Atomic orbitals –Orbital energies –Electron configuration and the periodic table Periodic table –Periodic properties –Energy

3 Characteristics of Atoms Atoms possess mass Atoms contain positive nuclei Atoms contain electrons Atoms occupy volume Atoms have various properties Atoms attract one another Atoms can combine with one another to form molecules

4 Atomic Structure Atomic structure studied through atomic interaction with light Light: electromagnetic radiation –carries energy through space –moves at 3.00 x 10 8 m/s in vacuum –wavelike characteristics

5 Electromagnetic Spectrum

6 Visible Spectrum

7 Wavelength ( ) & Frequency ( ) = number of complete cycles to pass given point in 1 second amplitude

8 Energy c = x = 3.00 x 10 8 m/s long wavelength  low frequency short wavelength  high frequency Low Energy High Energy

9 Energy Mathematical relationship: E = h E = energy h = Planck’s constant: 6.63 x 10 –34 J s = frequency in s –1

10 Energy Mathematical relationship: E = h c = x E = Energy:directly proportional to frequency inversely proportional to wavelength

11 Problems 3-1, 2, & 3 1.a) Calculate the wavelength of light with a frequency = 5.77 x s –1 b) What is the energy of this light? 2. Which is higher in energy, light of wave- length of 250 nm or light of 5.4 x 10 –7 m? 3. a) What is the frequency of light with an energy of 3.4 x 10 –19 J? b) What is the wavelength of light with an energy of 1.4 x 10 –20 J?

12 Photoelectric Effect Light on metal surface Electrons emitted Threshold frequency, o If < o, no photoelectric effect If > o, photoelectric effect As , kinetic energy of electrons 

13 Photoelectric Effect Einstein: energy  frequency If < o electron doesn’t have enough energy to leave the atom If > o electron does have enough energy to leave the atom Energy is transferred from light to electron, extra is kinetic energy of electron E photon = h photon = h o + KE electron KE electron = h photon – h o Animation

14 Problem 3-4 A given metal has a photoelectric threshold frequency of o = 1.3 x s  1. If light of = 455 nm is used to produce the photoelectric effect, determine the kinetic energy of the electrons that are produced.

15 Bohr Model Line spectra Light through a prism  continuous spectrum: Ordinary white light

16 Bohr Model Line spectra Light from gas-discharge tube through a prism  line spectrum: H 2 discharge tube

17 Line Spectra (emission) White light H He Ne

18 Line Spectra (absorption) Light source Gas-filled tube

19 Bohr Model For hydrogen: C = 3.29 x s –1 Niels Bohr: Electron energy in the atom is quantized. n = 1, 2, 3,…. R H = 2.18 x 10 –18 J

20 Bohr Model  E atom =  E electron = h  E = E f – E i = h Line spectrum Photoelectric effect: Minus sign: free electron has zero energy

21 Bohr Energy Levels

22 Electrons All electrons have same charge and mass Electrons have properties of waves and particles (De Broglie)

23 Heisenberg Uncertainty Principle Cannot simultaneously know the position and momentum of electron  x = h Recognition that classical mechanics don’t work at atomic level.

24 Schrödinger Equation Erwin Schrödinger 1926 Wave functions with discrete energies Less empirical, more theoretical  n E n  n wave functions or orbitals  n 2 probability density functions

25 Quantum Numbers Each orbital defined by 3 quantum numbers Quantum number: number that labels state of electron and specifies the value of a property

26 Quantum Numbers Principal quantum number, n (shell) Specifies energy of electron (analogous to Bohr’s n) Average distance from nucleus n = 1, 2, 3, 4…..

27 Quantum Numbers Azimuthal quantum number, (subshell) = 0, 1, 2… n–1 n = 1, = 0 n = 2, = 0 or 1 n = 3, = 0, 1, or 2 Etc spdfg

28 Quantum Numbers Magnetic quantum number, m Describes the orientation of orbital in space m = – ….+ If = 2, m = –2, –1, 0, +1, +2

29 Problem 3-5 Fill in the quantum numbers in the table below. n m 3003s 2 –2, –1, 0, 1, 2 0 2p

30 Schrödinger Equation Wave equations:  Each electron has  & E associated w/ it Probability Density Functions:  2 -graphical depiction of high probability of finding electron

31 Probability Density Functions  energy  2 probability density function s, p, d, f, g 1s 2s 3s Node: area of 0 electron density Link to Ron Rinehart’s page

32 Probability Density Functions 2p Node: area of 0 electron density 3p nodes Link to Ron Rinehart’s page

33 Electrons and Orbitals Pauli Exclusion Principle: no two electrons in the same atom may have the same quantum numbers Electron spin quantum number m s =  ½ Electrons are spin paired within a given orbital

34 Electrons and Orbitals n = 1 = 0, m = 0, m s =  ½ 2 electrons possible: 1,0,0,+½ and 1,0,0,–½ 2 electrons per orbital 1s 1 H 1s 2 He

35 Electrons and Orbitals n = 2 = 0, m = 0, m s =  ½ 2,0,0,  ½ 2 electrons possible n = 2 = 1, m = –1,0,+1, m s =  ½ 2,1,–1,  ½ 2,1,0,  ½2,1,+1,  ½ 6 electrons possible

36 Electron Configurations n = 1 1s2 electrons possible H1e – 1s 1  He 2e – 1s 2 

37 Electron Configurations n = 2 2s2 electrons possible  Li 3e – 1s 2 2s 1 1s  2s  Be 4e – 1s 2 2s 2 1s  2s

38 Electron Configurations n = 2 2p = 1, m = –1, 0, +1 3 x 2p orbitals (p x, p y, p z ): 6 electrons possible B 5e – 1s 2 2s 2 2p 1 1s 2s 2p

39 Electron Configurations n = 2 2p = 1, m = –1, 0, +1 3 x 2p orbitals (p x, p y, p z ): 6 electrons possible  B 5e – 1s 2 2s 2 2p 1 1s  2s 2p

40 Electron Configurations n = 2 2p = 1, m = –1, 0, +1 C 6e – 1s 2 2s 2 2p 2  1s   2s 2p Hund’s Rule: for degenerate orbitals, the lowest energy is attained when electrons w/ same spin is maximized

41 Problem 3-6 Write electron configurations and depict the electrons for N, O, F, and Ne.

42 Electron Configurations n = 3 3s, 3p, 3d Na 11e – 1s 2 2s 2 2p 6 3s 1  1s  2s  2p  3s 3p

43 Electron Configurations n = 3 3s, 3p, 3d Mg 12e – 1s 2 2s 2 2p 6 3s 2  1s  2s  2p  3s 3p

44 Electron Configurations n = 3 3s, 3p, 3d Al 13e – 1s 2 2s 2 2p 6 3s 2 3p 1  1s  2s  2p   3s 3p

45 Electron Configurations n = 3 3s, 3p, 3d Si 14e – 1s 2 2s 2 2p 6 3s 2 3p 2  1s  2s  2p   3s 3p

46 Electron Configurations n = 3 3s, 3p, 3d P 15e – 1s 2 2s 2 2p 6 3s 2 3p 3  1s  2s  2p   3s  3p

47 Electron Configurations n = 3 3s, 3p, 3d S 16e – 1s 2 2s 2 2p 6 3s 2 3p 4  1s  2s  2p  3s  3p

48 Electron Configurations n = 3 3s, 3p, 3d Cl 17e – 1s 2 2s 2 2p 6 3s 2 3p 5  1s  2s  2p  3s  3p

49 Electron Configurations n = 3 3s, 3p, 3d Ar 18e – 1s 2 2s 2 2p 6 3s 2 3p 6  1s  2s  2p  3s  3p

50 Electron Configurations 3d vs. 4s Filling order 1s 2s2p 3s3p3d 4s4p4d4f 5s5p5d5f5g 6s6p6d 7s7p

51 Electron Configurations  1s  2s  2p  3s  3p  4s 4p 3d K

52 Electron Configurations  1s  2s  2p  3s  3p  4s 4p 3d Ca

53 Electron Configurations  1s  2s  2p  3s  3p  4s 4p  3d Sc

54 Electron Configurations  1s  2s  2p  3s  3p  4s 4p  3d Ti Link to OSU siteOSU

55 Problem 3-7 Write the electron configurations for the transition metals V – Zn. Fill in the corresponding boxes to denote the electronic spin.


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