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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 Quantum numbers

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**Atomic Structure Atomic orbitals Periodic table Orbital energies**

Electron configuration and the periodic table Periodic table Periodic properties Energy

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**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

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Atomic Structure Atomic structure studied through atomic interaction with light Light: electromagnetic radiation carries energy through space moves at 3.00 x 108 m/s in vacuum wavelike characteristics

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**Electromagnetic Spectrum**

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Visible Spectrum

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**Wavelength () & Frequency ()**

amplitude = number of complete cycles to pass given point in 1 second

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**Energy c = x = 3.00 x 108 m/s long wavelength low frequency**

Low Energy High Energy short wavelength high frequency

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**Energy Mathematical relationship: E = h E = energy**

h = Planck’s constant: 6.63 x 10–34 J s = frequency in s–1

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**Energy Mathematical relationship: E = h c = x E =**

Energy: directly proportional to frequency inversely proportional to wavelength

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Problems 3-1, 2, & 3 a) Calculate the wavelength of light with a frequency = 5.77 x 1014 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?

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**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

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**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 Ephoton = hphoton = ho + KEelectron KEelectron = hphoton – ho Animation

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Problem 3-4 A given metal has a photoelectric threshold frequency of o = 1.3 x 1014 s1. If light of = 455 nm is used to produce the photoelectric effect, determine the kinetic energy of the electrons that are produced.

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**Bohr Model Line spectra Light through a prism continuous spectrum:**

Ordinary white light

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**Bohr Model Line spectra Light from gas-discharge tube**

through a prism line spectrum: H2 discharge tube

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**Line Spectra (emission)**

White light H He Ne

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**Line Spectra (absorption)**

Gas-filled tube Light source

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**Bohr Model For hydrogen: C = 3.29 x 1015 s–1**

Niels Bohr: Electron energy in the atom is quantized. n = 1, 2, 3,…. RH = 2.18 x 10–18 J

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**Bohr Model E = Ef – Ei = h Eatom = Eelectron = h Line spectrum**

Minus sign: free electron has zero energy Line spectrum Photoelectric effect:

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Bohr Energy Levels

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**Electrons All electrons have same charge and mass**

Electrons have properties of waves and particles (De Broglie)

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**Heisenberg Uncertainty Principle**

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

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**Schrödinger Equation Erwin Schrödinger 1926**

Wave functions with discrete energies Less empirical, more theoretical n En n wave functions or orbitals n2 probability density functions

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**Quantum Numbers Each orbital defined by 3 quantum numbers**

Quantum number: number that labels state of electron and specifies the value of a property

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**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…..

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**Quantum Numbers Azimuthal quantum number, (subshell) = 0, 1, 2… n–1**

n = 2, = 0 or 1 n = 3, = 0, 1, or 2 Etc. 1 2 3 4 s p d f g

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**Quantum Numbers Magnetic quantum number, m**

Describes the orientation of orbital in space m = –….+ If = 2, m = –2, –1, 0, +1, +2

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**Problem 3-5 Fill in the quantum numbers in the table below. n m 3**

3s 2 –2, –1, 0, 1, 2 2p

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**Schrödinger Equation Wave equations: **

Each electron has & E associated w/ it Probability Density Functions: 2 -graphical depiction of high probability of finding electron

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**Probability Density Functions**

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

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**Probability Density Functions**

Node: area of 0 electron density nodes Link to Ron Rinehart’s page

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**Electrons and Orbitals**

Pauli Exclusion Principle: no two electrons in the same atom may have the same quantum numbers Electron spin quantum number ms = ½ Electrons are spin paired within a given orbital

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**Electrons and Orbitals**

= 0, m = 0, ms = ½ 2 electrons possible: 1,0,0,+½ and 1,0,0,–½ 2 electrons per orbital 1s1 H 1s2 He

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**Electrons and Orbitals**

= 0, m = 0, ms = ½ 2,0,0, ½ 2 electrons possible = 1, m = –1,0,+1, ms = ½ 2,1,–1, ½ ,1,0, ½ 2,1,+1, ½ 6 electrons possible

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**Electron Configurations**

1s 2 electrons possible H 1e– 1s1 He 2e– 1s2

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**Electron Configurations**

2s 2 electrons possible Li 3e– 1s2 2s1 2s 1s Be 4e– 1s2 2s2 2s 1s

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**Electron Configurations**

2p = 1, m = –1, 0, +1 3 x 2p orbitals (px, py, pz): 6 electrons possible 2p B 5e– 1s2 2s2 2p1 2s 1s

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**Electron Configurations**

2p = 1, m = –1, 0, +1 3 x 2p orbitals (px, py, pz): 6 electrons possible 2p B 5e– 1s2 2s2 2p1 2s 1s

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**Electron Configurations**

2p = 1, m = –1, 0, +1 C 6e– 1s2 2s2 2p2 1s 2s 2p Hund’s Rule: for degenerate orbitals, the lowest energy is attained when electrons w/ same spin is maximized

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

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**Electron Configurations**

3s, 3p, 3d 1s 2s 2p 3s 3p Na 11e– 1s2 2s2 2p63s1

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**Electron Configurations**

3s, 3p, 3d 1s 2s 2p 3s 3p Mg 12e– 1s2 2s2 2p63s2

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**Electron Configurations**

3s, 3p, 3d 1s 2s 2p 3s 3p Al 13e– 1s2 2s2 2p63s23p1

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**Electron Configurations**

3s, 3p, 3d 1s 2s 2p 3s 3p Si 14e– 1s2 2s2 2p63s23p2

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**Electron Configurations**

3s, 3p, 3d 1s 2s 2p 3s 3p P 15e– 1s2 2s2 2p63s23p3

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**Electron Configurations**

3s, 3p, 3d 1s 2s 2p 3s 3p S 16e– 1s2 2s2 2p63s23p4

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**Electron Configurations**

3s, 3p, 3d 1s 2s 2p 3s 3p Cl 17e– 1s2 2s2 2p63s23p5

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**Electron Configurations**

3s, 3p, 3d 1s 2s 2p 3s 3p Ar 18e– 1s2 2s2 2p63s23p6

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**Electron Configurations**

3d vs. 4s Filling order 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 5g 6s 6p 6d 7s 7p

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**Electron Configurations**

4p K 3d 4s 3p 3s 2p 2s 1s

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**Electron Configurations**

4p Ca 3d 4s 3p 3s 2p 2s 1s

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**Electron Configurations**

4p 3d Sc 4s 3p 3s 2p 2s 1s

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**Electron Configurations**

4p Ti 3d 4s 3p 3s 2p 2s Link to OSU site 1s

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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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