Understanding Electronic Structure of Atoms: Light and Quantum Concepts

1. 14 November 2011 • Objective: You will be able to: • describe evidence for the current theory of the electronic structure of atoms. • Homework: p. 312 #3, 4, 5, 6, 7, 9, 16, 19, 25, 32

2. Electronic Structure of Atoms

3. Next Units: • Electron configuration • Trends on the periodic table • Ionic/covalent bonding • Chemical reactivity

4. In order to understand these things • we’ll study the electronic structure of atoms

5. The Wave Nature of Light • electromagnetic radiation (a.k.a. light) is a form of energy with wave and particle characteristics. It moves through a vacuum at the speed of light • speed of light: 3.00x108 m/s

6. To describe waves… • wavelength (λlamda): the distance between two adjacent peaks of a wave • frequency (v): the number of wavelengths that pass a given point in a second

7. Electromagnetic Spectrum

8. electromagnetic spectrum includes all wavelengths of radiant energy • visible spectrum: the part of the electromagnetic spectrum that is visible to the human eye (wavelengths between 400 and 700 nm)

9. Quantized Energy and Photons • quantum (a.k.a. photon) is a specific particle of light energy that can be emitted or absorbed as electromagnetic radiation. • Energy of a photon E=hv • Energy is quantized – matter is allowed to emit or absorb energy in discrete amounts, whole number multiples of hv.

10. How are these things related to electromagnetic radiation? E=hc/λ

11. Example 1 Calculate the energy (in joules) of • a photon with a wavelength of 5.00x104 nm (infrared region) • a photon with a wavelength of 5.00x10-2 nm (x-ray region)

12. Example 2 • What is the frequency and the energy of a single photon? • What is the energy of a mole of photons of light having a wavelength of 555 nm?

13. Problem • The energy of a photon is 5.87x10-20 J. What is its wavelength, in nanometers?

14. Homework • p. 312 #3, 4, 5, 6, 7, 9, 16, 19, 25, 32

15. 15 November 2011 • Take Out Homework • Objective: You will be able to: • describe and explain experimental evidence for energy levels • Homework Quiz: The energy of a photon is 3.98x10-19 J. What color light do you observe?

16. Agenda • Homework Quiz • Hand back tests • Line spectra and the Bohr model of the atom Homework: p. 313 #23, 24, 25, 26, 30, 31, 35, 36

17. Line Spectra and the Bohr Model • atomic emission spectrum (a.k.a. line spectrum): a pattern of discrete lines of different wavelengths that result when the light energy emitted from energized atoms is passed through a prism • Each element produces a characteristic or identifiable pattern

19. Demo • Emission spectra of common cations • Note: we don’t have a way to separate all the wavelengths of light into discrete lines of color, so we’re just seeing all those lines of color blended together. • http://www.youtube.com/watch?v=2ZlhRChr_Bw&feature=related

20. So, why do we see these discrete lines of color? • Bohr model of the atom: energies are quantized. Electrons move in circular, fixed energy orbits around the nucleus. • Usually, electrons are in the most stable “ground” state. • When energy (a photon) is added, they “jump” up to the “excited” state • They fall back down, and release that photon.

21. Multimedia • http://www.youtube.com/watch?v=45KGS1Ro-sc • http://www.colorado.edu/physics/2000/quantumzone/lines2.html

22. Homework • p. 313 #23, 24, 25, 26, 30, 31, 35, 36

23. 16 November 2011 • Objective: You will be able to: • explain how line spectra give evidence for the existence of energy levels • explain how quantum mechanics describes electron configuration

24. Agenda • Homework Quiz • Go over homework • How do atoms emit photons? • Quantum mechanics: how do we describe where the electrons are?! • Writing orbital notation and electron configuration Homework: p. 313 #23-26, 30, 35, 48, 53, 60, 63,

25. Energy levels

26. Wave Behavior of Matter • Like light, electrons have characteristics of both waves and particles. Because a wave extends into space, its location is not precisely defined. • uncertainty principle: it is impossible to simultaneously determine the exact position and momentum of an electron. • we can only determine the probability of finding an electron in a certain region of space.

27. Quantum Mechanics and Atomic Orbitals • quantum mechanical model: mathematical model that incorporates both the wave and particle characteristics of electrons in atoms. • quantum numbers: describe properties of electrons and orbitals • each electron has a series of four quantum numbers

28. Table of Quantum Numbers

29. Table of quantum numbers and orbital designations

30. Pauli Exclusion Principle • Two electrons in an atom can’t have the same four quantum numbers • Two electrons per orbital, with opposite spins

31. Representations of Orbitals • orbital: calculated probability of finding an electron of a given energy in a region of space

32. p orbitals

33. d orbitals

34. orbital ≠ orbit

35. 17 November 2011 • Objective: You will be able to: • write the orbital and electron configuration for any element • describe several exceptions to the orbital filling rules • Homework Quiz: Describe, as completely as you can in a paragraph or two, the evidence that convinced Neils Bohr of the existence of energy levels instead of a cloud of electrons.

36. Agenda • Homework Quiz • Go over homework • Electron configuration notation • Problem Set Unit 4 Quiz Weds.

37. Atoms with more than one electron • like hydrogen • electron-electron repulsions cause different sublevels to have different energies

38. Order those orbitals fill

39. Electron Configuration • distribution of electrons among various orbitals of an atom

41. Rules for Writing E- Config. • at the ground state • Fill the lowest energy level first. Electrons in the same orbital must have opposite spins. Total number of electrons = atomic number • Only two electrons per orbital! • Do not pair electrons in a orbitals of the same energy until each orbital has one electron of the same spin (Hund’s rule) • Label each sublevel with the energy level number and letter of the sublevel

42. Examples • phosphorus • calcium • iron

43. Paired-ness of Electrons • Paramagnetic: an atom having one or more unpaired electrons • Ex: Li, B, C… • Diamagnetic: all electrons in an atom are paired. • Ex:

44. Excited-State Configuration • has a higher energy than the ground-state electron configuration. • One or more electrons occupy higher energy levels than predicted by the rules • Ex: Iron in an excited state:

45. Electron Configuration and the Periodic Table • Elements with similar electron configurations arranged in columns

46. Examples • Write the electron configuration for palladium • Write the electron configuration for osmium

47. Condensed Electron Config. • shows only the electrons occupying the outermost sublevels • preceded by the symbol for the noble gas in the row above the element • Example: calcium • Example: iodine

48. Unusual Electron Configs. • Cr and Mo: ground state valence electrons are arranged s1d5 rather than s2d4 • a half filled d orbital is more stable than a more-than-half-filled d orbital • Cu, Ag and Au have s1d10 ground state configs because of the stability of a fill d orbital

49. 21 November 2011 • Objective: You will be able to: • describe the electronic structure of an atom and make associated calculations.

50. Agenda • Math with exponents (#6) • Problem set work time Homework: Problem set due tomorrow Quiz Mon. on all electronic structure, calculations, evidence for Bohr’s theory…