Electronic Structure and the Periodic Table

1. Electronic Structureand the Periodic Table Unit 6 Honors Chemistry

2. Electromagnetic Waves: Electromagnetic waves: progressive, repeating disturbances that come from the movement of electric charges Electromagnetic Waves & Light

3. Wavelength and Frequency • Wavelength (, lambda): distance between any two points in a wave • measured in any distance unit (mainly nm or m: 1 nm = 1x10-9 m)

4. Wavelength Can be Measured in One of Two Ways…

5. Wavelength and Frequency • Frequency (; pronounced nu): the number of cycles of the wave that pass through a point in a unit of time • Measured in sec-1 (/sec) • 1 sec-1 = 1 Hertz (Hz)

6. Illustration of Frequency

7. Wavelength is indirectly proportional to frequency As Wavelength increases, frequency _________________. As Wavelength decreases, frequency _________________.

8. Amplitude • Note: height of wave is amplitude (intensity or brightness of wave) • Amplitude is INDEPENDENT of frequency or wavelength!

9. Speed • Speed (c): The speed of light! c = 3.00 x 108 m/s (rounded to 3 sig figs)

10. Equation • One equation relates speed, frequency and wavelength: c =  

11. Example • The wavelength of the radiation which produced the yellow color of sodium vapor light is 589.0 nm. What is the frequency of this radiation?

12. The electromagnetic spectrum • complete range of wavelengths and frequencies • mostly invisible

13. What is color? TED Talk: What is color?

14. The visible/continuous spectrum • continuous spectrum: components of white light split into its colors, ROY G BIV • from 390 nm (violet) to 760 nm (red)

15. Line Spectra • Pattern of lines produced by light emitted by excited atoms of an element • unique for every element • used to identify unknown elements

16. How do we see color? TED Talk: How we see color

17. Max Planck • Light is generated as a stream of particles called PHOTONS • Equation: E (Energy of a photon)= h (h =Plank’s constant= 6.626x10-34Js)

18. Relationships in Planck’s Eqn. E = h •  High frequency, low λ, high E. Low frequency, high λ, low E.

19. Photoelectric effect – Nobel Prize in Physics 1921 to Einstein Occurs when light strikes the surface of a metal and electrons are ejected. Practical uses: Automatic door openers

20. Photoelectric Effect: Conclusion Light not only has wave properties but also has particle properties. These massless particles, called photons, are packets of energy.

21. Example 6.2 Using the frequency calculated in the previous example, calculate the energy, in joules, of a photon emitted by an excited sodium atom. Calculate the energy, in kilojoules, of a mole of excited sodium atoms.

22. Bohr’s Hydrogen Atom: A Planetary Model Niels Bohr: Proposed planetary model. Electrons “orbit” the nucleus like planets around the sun. NOT current model of atom but used to explain some features of atom.

23. Ground State vs. Excited State • ground state: all electrons in lowest possible energy levels • excited state: an electron that has absorbed energy and moved to a higher energy level • This is a temporary state!!

24. Explanation of Line Spectra & Equation Niels Bohr • Energy of an electron is quantized: can only have specific values. • Energy proportional to energy level.

25. Explanation of Line Spectra Electron will drop from excited state to ground state and will emit energy as a photon.

26. Explanation of Line Spectra • Type of photon emitted by electron depends on energy difference of energy levels Elevel = -RH1 – 1 (nhi)2 (nlow)2 AND Elevel = h = hc/ (h: Planck’s constant, 6.626 x 10-34 J sec/photon)

27. Flaw in Bohr’s Model Only works well for 1 electron species (H atom). Does not explain fine structure of line spectra.

28. Wave-Particle Duality • Light has properties of both WAVES and PARTICLES. • most matter has undetectable wavelengths (1000 kg car at 100 km/hr has  = 2.39 x 10-38 m) • This work led to the development of the electron microscope

29. Quantum Mechanics • Quantum mechanics: atomic structure based on wave-like properties of the electron • Schrödinger: wave equation that describes hydrogen atom

30. Heisenberg Uncertainty Principle • The exact location of an electron cannot be determined (if we try to observe it, we interfere with the particle) • You can know either the location or the velocity but not both • Electrons exist in electron clouds and not on specific rings or orbits

31. Quantum Numbers • Four quantum numbers are a mathematical way to represent the most probable location of an electron in an atom • analogy... state = energy level, n city = sublevel, l address = orbital, ml house number = spin, ms

32. Principal Quantum Number: n • Always a positive integer (1,2, 3…7) • Indicates size of orbital, or how far electron is from nucleus • Similar to Bohr’s energy levels or shells • Larger n value = larger orbital or distance from nucleus

33. n = 1 n = 2 n = 3 n = 4 n = 5 n = 6 n = 7 The Periodic Table and Shells n = row number on periodic table for a given element

34. Angular Momentum Quantum Number: l • positive integer from zero to n-1 • Sublevel within an energy level; indicates shape of orbital • 0 = s • 1 = p • 2 = d • 3 = f

35. Types of Sublevels s p d

36. Magnetic Quantum Numbers: ml • integer from -l to +l • Indicates orientation of orbital in space • Orbital = electron containing area

37. Spin Quantum Number: ms • Two values only: + ½ or -½ • 2 electrons max. allowed in each orbital • (Pauli Exclusion Principle) • Indicates spin of electron; spins of each electron must be opposite

38. REVIEW: QUANTUM NUMBERS n ---> level 1, 2, 3, 4, ... l ---> sublevel 0, 1, 2, ... n - 1 ml ---> orbital -l ... 0 ... +l ms ---> electron spin +½ and -½ Every Electron has four!

39. Orbitals • No more than 2 e- assigned to an orbital • Orbitals grouped in s, p, d (and f) subshells s orbitals p orbitals d orbitals

40. Capacities of levels, sublevels, and orbitals—see packet

41. Example Example 6.6 Give the n and l values for the following orbitals: a. 3p b. 4s

42. Example Example 6.8 What are the possible ml values for the following orbitals: a. 3p b. 4f

43. Shapes of Atomic Orbitals

44. Shapes of Atomic Orbitals s = spherical p = peanut d = dumbbell (clover) f = flower

45. Multielectron Atoms In the hydrogen atom the subshells(sublevels) of a principal energy level or shell are at the same energy level. Previous Equation: En = –RH /n2

46. Multielectron Atoms In a multielectron atom, only the orbitalsare at the same energy level: the sublevels are at different energy levels!

47. The increasing energy order of sublevels is generally: s < p < d < f

48. Overlapping subshells At higher energy levels, sublevels overlap. Note: 4s vs. 3d!

49. EOS Introduction to Electron Configuration Definition:describes the distribution of electrons among the various orbitals in the atom Represents the most probable location of the electron!

50. Electron Configurations • The system of numbers and letters that designates the location of the electrons • 3 major methods: • Full electron configurations • Abbreviated/Noble Gas configurations • Orbital diagram configurations