Electronic Structure

1. Electronic Structure

2. Wave Nature of Light • Electromagnetic Radiation • Gamma Rays, Visible Light • Moves Through Vacuum at 3.00x108m/s (c) • Wavelength (λ, m) = distance between successive peaks or troughs • Frequency (f, s-1) = how often a wave passes through a particular point • c = f ·λ, or v · λ

3. Wavelength Practice • The brilliant red colors seen in fireworks are due to the emission of light with wavelengths around 650nm when strontium salts such as Sr(NO3)2 and SrCO3 are heated. Calculate the frequency of red light of wavelength 6.50x102nm. • A FM radio station broadcasts electromagnetic radiation at a frequency of 103.4MHz. Calculate the wavelength of this radiation. (1MHz=106s-1)

4. Quantize Energy and Photons • Wave model explains much of the behavior of light but not all: • Black body radiation – Emission of light from hot objects • Photoelectric Effect – Emission of electrons from metal surfaces • Emission Spectra – Emission of light from excited atoms

5. Black Body Radiation

6. Hot Objects and Quantization of Energy • When objects are heated they emit light • Red hot (cooler) → white hot (hotter) • Light only emitted at certain wavelengths • Max Planck declared that energy can only be emitted or absorbed in packets (quanta, photon) • E= h · v • h = Planck's Constant 6.63x10-34J·s • Energy emitted at whole number multiples of hv • Think walking up and down stairs • Why don't we notice this? Why does energy seem to flow continuously for us?

7. Photoelectric Effect and Photons • When light shines on an object, electrons are emitted • Light has to have a specific energy, frequency and wavelength in order for e- to be emitted • Photons are absorbed • Too little energy – nothing happens • Just right amount – electrons are emitted • A little too much – electrons are emitted and excess used as kinetic energy

8. Calculation Practice • The blue color of fireworks is often achieved by heating copper (I) chloride to about 1200°C. Then the compound emits blue light having a wavelength of 450nm. Calculate the frequency and quantum of energy that is emitted at 4.50x102nm by CuCl.

9. Spectra • Radiation emitted from a source contains various λ's • When separated into its different λ's a spectrum is formed • Two types of spectrum • Continuous • Line

10. Continuous Spectra

11. Line Spectra

12. Bohr Model • Assumed electrons orbit nucleus in circular patterns • To move between levels energy is absorbed or emitted • Ground State – electron at lowest energy • Excited State – electron is at higher energy state • Bohr model only accurately explains hydrogen

13. Wave Behavior of Matter • Lights is both a wave and a particle • Louis de Broglie believed matter could have wave properties • Applied idea to electrons • λ=h/m·v • Works for all matter so why don't we observed this in our everyday lives?

14. Uncertainty Principle • If matter can act as a wave we should be able to calculate position and velocity • Heisenberg determined that we cannot know both position and velocity of subatomic matter. • Solving Schrödinger's equation gives use probabilities of location. • The solutions correspond to the orbitals

15. Quantum Numbers • n – principle quantum number • Whole numbers – 1,2,3,.... • l – azimuthal quantum number • From 0 to n-1 • Determines shape of orbital • 0 → s = shape • 1 → p = principle • 2 → d = diffuse • 3 → f = fundamental

16. Quantum Numbers Cont. • ml – magnetic quantum number • Goes from -l to l including 0 • Determines orientation of orbital • ms – spin quantum number • Two values +1/2 and -1/2 • Determines spin of electron • No two electrons can have the same four quantum numbers – Pauli Exclusion Principle

17. Possible Quantum Numbers

18. Quantum Numbers Example • Which of the following sets of quantum numbers are not allowed? For each incorrect set, state why it is incorrect. • n = 3, l = 3, ml = 0, ms = -1/2 • n = 4, l = 3, ml = 2, ms = -1/2 • n = 4, l = 1, ml = 1, ms = +1/2

19. Quantum Numbers Practice • Which of the following sets of quantum numbers are not allowed? For each incorrect set, state why it is incorrect. • n = 2, l = 1, ml = -1, ms = -1 • n = 5, l = -4, ml = 2, ms = +1/2 • n = 3, l =1, ml = 2, ms = -1/2

20. Quantum Number Practice • What is the designation for the subshell with n = 5 and l = 1? How many orbitals are in this subshell? Indicate the values of ml for each of these orbitals.

21. Representation of Orbitals • S – orbital is a sphere

22. Representation of Orbitals • P Orbital

23. Representation of Orbitals • D Orbital

24. Atoms with Multiple Electrons • Shapes of orbitals remain the same • Energy of orbitals varies • For a given n, energy increases as l increases

25. Electron Configuration • Governed by three rules • Pauli Exclusion • Hund's Rule – for orbitals with same energy, lowest energy is attained when the number of electrons with the same spin is maximized • AUFBAU – Energy shells are filled from lowest energy to highest energy

26. Orbital Diagram

27. Where Orbitals are Filled

28. Filling Order

29. Electron Configuration Example • Write the electron configuration for the following elements: • Silicon • Chromium • Iodine

30. Electron Configuration Practice • Write the electron configuration for the following elements: • Copper • Sulfur • Tin

31. Exceptions to Filling Rules • Chromium • Copper • Silver • Molybdenum

32. Homework • 2, 6, 10, 14, 20, 22, 26, 46, 52, 54, 60, 68

33. Electron Configurations