Chemistry Chemical Bonding

1. ChemistryChemical Bonding

2. The Development of Atomic Models • Dalton – solid, indivisible mass

3. Thomson – “plum-pudding” model Negatively charged e- (raisins) stuck in positively charged proton dough No neutrons

4. Rutherford – electrons surrounding a dense nucleus

5. Bohr model – elctrons arranged in symmetrical orbits around the nuclues “planetary model” Electrons in a given path have a fixed energy level

6. 5. Quantum mechanical model – modern mathematical description of the atom Sodium atom:

7. Energy level – region around the nucleus where the electron is likely to be moving. An electron can jump from one level to another by absorbing energy.

8. Quantum – the amount of energy required to move an electron from its present energy level to the next higher one “quantum leap”

9. Quantum mechanical model – uses mathematical equations to describe the location and energy of electrons in an atom • Developed by Erwin Schrodinger • Electrons are not in definite paths • Their location is described in terms of probability of being in a certain region • Electron cloud (ceiling fan) • Conventionally, the border is drawn at 90% probability

10. Atomic orbital – region in space that an electron is likely to be in Electrons can be described by a series of 4 quantum numbers.

11. 1. Principle quantum number (n) • Describes the energy level • Values of 1, 2, 3, 4, etc.

12. 2. Azimuthal quantum number ( l ) • Describes the shape of atomic orbitals • Sublevels • Values of 0 to n-1 • 0 = s, 1 = p, 2 = d, 3 = f

13. s = spherical, p = peanut shape, d&f = more complex shapes d = “daisy”f = “fancy” So if n = 1, then l can be 0 (s) = 1 sublevel n = 4, then l can be 0 (s), 1 (p), 2(d) , 3(f), = 4 sublevels

14. s orbitals – spherical

15. p orbitals – dumbbell-shaped

16. d orbitals – daisy-shaped

17. f orbitals – fancy shapes

18. 3. Magnetic quantum number (ml) • Orientation of the orbital in space • Values of –l to +l So s has 1 orbital p has 3 d has 5 f has 7

19. 4. Spin quantum number (ms) • Values of +½ and -½ • Each orbital can hold 2 electrons with opposite spins • Since spinning charged objects create a magnetic field, the electrons must spin opposite directions to minimize repulsion

20. electron spins

21. Ex. How many orbitals are in the following? A. 3p D. 4p B. 2s E. 3d C. 4f F. 3rd energy level How many electrons can be in each of the above?

22. Are these possible?

23. Electron Configuration – ways in which electrons are arranged around nuclei of atoms.

24. Rules that govern filling of atomic orbitals: 1. Aufbau principle – electrons enter orbitals of lowest energy first.

26. 2. Pauli exclusion principle – An atomic orbital can describe at most two electrons. They must have opposite spins.

27. 3. Hund’s rule – When electrons occupy orbital of equal energy, one electron enters each orbital until all orbitals contain one electron with parallel spins.

29. 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 4f

30. Na = 11 1s22s22p63s1 Cd =48 1s22s22p63s23p64s23d104p65s24d10

31. Practice: Write the electron configuration for the following elements: Li O Sc

32. More practice: Identify each of the following atoms on the basis of its electron configurations. a) 1s22s22p6 neon b) 1s22s22p63s1 sodium c) [Kr] 5s24d2 zirconium d) [Xe] 6s24f6 samarium

33. Ground state – lowest energy level for an electron. (Normal, nonexcited state)

34. Exceptional Electron Configurations: Cr: expected: 1s22s22p63s23p64s23d4 actual: 1s22s22p63s23p64s13d5 Cu: expected: 1s22s22p63s23p64s23d9 actual: 1s22s22p63s23p64s13d10 Half-filled energy levels are more stable than other partially filled energy levels. There are other exceptions.

35. Light and Atomic Spectra Electromagnetic radiation – a series of energy waves that includes radio waves, microwaves, visible light, infrared, and ultraviolet light, X-rays, and gamma rays.

36. Electromagnetic Spectrum

37. Parts of a wave: Wavelength, l crest trough

38. Amplitude – height of the wave from the origin to the crest. Wavelength - l - distance between the crests Frequency – f (or n) – the number of wave cycles to pass a given point per unit of time. The units of frequency are 1/s, s-1, or Hertz (Hz)

39. c= lf where c = speed of light = 3.00x108 m/s or 3.00x1010 cm/s As l increases, f decreases.

40. As wavelength increases, frequency decreases. As wavelength decreases, frequency increases.

42. Ex. A certain wavelength of yellow light has a frequency of 2.73x1016 s-1. Calculate its wavelength. Convert to nm. C = lf l = c/f l = (3.00x108 m/s)/(2.73x1016 s-1) l = 1.10x10-8 m

43. Spectrum – series of colors produced when sunlight is separated by being passed through a prism. ROYG. BIV Red: longest wavelength, lowest frequency Violet: shortest wavelength, highest frequency

45. Atomic emission spectrum – series of lines of colored light produced by passing light emitted by an “excited atom through a prism.” This can be used to identify the element. The atomic emission spectrum of hydrogen shows three series of lines. The lines in the UV region (Lyman series) represent electrons falling to n=1, lines in the visible region (Balmer series) represent electrons falling to n=2 and lines in the IR region (Paschen series) represent electrons falling to n=3.

48. Max Planck found that the energy emitted or absorbed by a body changes only in small discrete units he called quanta. He determined that the amount of radiant energy, E, absorbed or emitted by a body is proportional to the frequency of the radiation. E=hf E = energy (J) f = frequency h = Planck’s constant, 6.626x10-34 Js

49. Einstein studied the photoelectric effect whereby light of sufficient frequency shining on a metal causes current to flow. The amplitude of the radiation was not important, the frequency was. This told him that light must be in particles, each having a given energy. Einstein proposed that electromagnetic radiation can be viewed as a stream of particles called photons: E=hf

50. Photoelectric Effect light Electron (photons) metal