1. Electronic Structure and the Periodic Table

2. Electromagnetic Radiation
Wave nature of light
wavelength - l
distance from the top (crest) of one wave to the top of the next wave
units of distance - m,cm, Å
1 Å = 1 x m = 1 x 10-8 cm
frequency - u
number of crests or troughs that pass a given point per second
units of 1/time - s-1

3. Electromagnetic Radiation
Speed of the wave, v
Frequency multiplied by wavelength
V = l u
For light, speed = c
relationship for electromagnetic radiation - c = l u
c = velocity of light x 108 m/s

4. Electromagnetic Radiation

5. Electromagnetic Radiation
What is the frequency of green light of wavelength 5200 Å?

6. Electromagnetic Radiation
What is the frequency of green light of wavelength 5200 Å?

7. Electromagnetic Radiation
Max Planck
energy is quantized
light has particle character
Planck’s equation

8. Electromagnetic Radiation
What is energy of a photon of green light with wavelength 5200 Å?

9. Electromagnetic Radiation
What is energy of a photon of green light with wavelength 5200 Å?

10. Electromagnetic Radiation
What is energy of a photon of green light with wavelength 5200 Å?

11. Atomic Spectra & Bohr Theory
emission spectrum
electric current passing through a gas in a vacuum tube (at very low pressure) causes the gas to emit light
emission or bright line spectrum

12. Line Spectra
Radiation composed of only one wavelength is monochromatic
Radiation that spans an array of different wavelengths is continuous
White light is continuous

13. Atomic Spectra & Bohr Theory
absorption spectrum
shining a beam of white light through a sample of gas gives an absorption spectrum
shows the wavelengths of light that have been absorbed

14. Atomic Spectra & Bohr Theory
spectra are fingerprints of elements
use spectra to identify elements
can even identify elements in stars

15. Atomic Spectra & Bohr Theory
“how atoms talk to us”
we have to interpret their language
Bohr, Schrodinger, and Heisenberg were some of the first scientists to translate the language of atoms

16. Atomic Spectra & Bohr Theory
An orange line of wavelength 5890 Å is observed in the emission spectrum of sodium. What is the energy of one photon of this orange light?

17. Atomic Spectra & Bohr Theory
An orange line of wavelength 5890 Å is observed in the emission spectrum of sodium. What is the energy of one photon of this orange light?

18. Atomic Spectra & Bohr Theory
Rydberg equation
empirical equation that relates the wavelengths of the lines in the hydrogen spectrum (Equ. 6.4 text)

19. Atomic Spectra & Bohr Theory
Neils Bohr incorporated Planck’s quantum theory into the H spectrum explanation
Postulates of Bohr’s theory

20. Atomic Spectra & Bohr Theory
Atom has a number of definite and discrete energy levels (orbits) in which an electron may exist without emitting or absorbing electromagnetic radiation.
increasing radius of orbit increases the energy
K<L<M<N<O......

21. Atomic Spectra & Bohr Theory
An electron may move from one discrete energy level (orbit) to another and in doing so monochromatic radiation is emitted or absorbed in accordance with the following equation.
E absorbed as electron jumps to higher orbit
E emitted as electron falls to lower orbit

22. Atomic Spectra & Bohr Theory
An electron moves in a circular orbit about the nucleus and its motion is governed by the ordinary laws of mechanics and electrostatics, with the restriction that the angular momentum of the electron is quantized (can only have certain discrete values).
angular momentum = mvr = nh/2p
h = Planck’s constant n = 1,2,3,4,...(energy levels)
v = velocity of electron m = mass of electron
r = radius of orbit

23. Atomic Spectra & Bohr Theory
Bohr theory correctly explains H emission spectrum
fails for all other elements
just not an adequate theory

24. The Origin of Spectral Lines
light of a characteristic wavelength (& frequency) is emitted when electron falls from higher E (orbit) to lower E (orbit)
Origin of the emission spectrum
light of a characteristic wavelength (& frequency) is absorbed when electron jumps from lower E (orbit) to higher E (orbit)
origin of absorption spectrum

25. The Wave Nature of the Electron
Louis de Broglie -1925
electrons have wave-like properties
their wavelengths are described by the de Broglie relationship

26. The Wave Nature of the Electron
verified by Davisson & Germer two years later
electrons (in fact - all particles) have both a particle and a wave like character
wave-particle duality is a fundamental property of submicroscopic particles

27. Quantum Mechanical Picture
Werner Heisenberg
Uncertainty Principle
It is impossible to determine simultaneously both the position & momentum of an electron.

28. Quantum Mechanical Picture
devices for detecting motion of electron disturbs its position
like measuring position of a car with a wrecking ball
must speak of electrons in terms of probability functions

29. Quantum Numbers Basic Postulates of Quantum Theory
Atoms and molecules can exist only in certain energy states. In each energy state, the atom or molecule has a definite energy. When an atom or molecule changes its energy state, it must emit or absorb just enough energy to bring it to the new energy state (the quantum condition).

30. Quantum Numbers
Atoms or molecules emit or absorb radiation (light) as they change their energies. The frequency of the light emitted or absorbed is related to the energy change by a simple equation.

31. Quantum Numbers
The allowed energy states of atoms and molecules can be described by sets of numbers called quantum numbers.

32. Quantum Numbers
Quantum numbers are solutions of the Schrodinger, Heisenberg & Dirac equations
electron wave functions
Four quantum numbers are necessary to describe energy states of electrons in atoms

33. Quantum Numbers Principal quantum number - n
n = 1, 2, 3, 4, “shells”
n = K, L, M, N,
electron’s energy depends principally on n

34. Quantum Numbers Subsidiary Quantum number - l
l = 0, 1, 2, 3, 4, 5, (n-1)
l = s, p, d, f, g, h, (n-1)
tells us the shape of orbitals
volume that the electrons occupy 90-95% of the time

35. Quantum Numbers Magnetic quantum number - ml l = 0, ml = 0
ml = - l, (- l + 1), (- l +2),.....0, ,(l -2), (l -1), l
l = 0, ml = 0
only 1 value
s orbital
l = 1, ml = -1,0,+1
3 values
p orbitals

36. Quantum Numbers l = 2, ml = -2,-1,0,+1,+2
5 values
d orbitals
l = 3, ml = -3,-2,-1,0,+1,+2, +3
7 values
f orbitals
theoretically, we can continue this series on to g,h,i, orbitals

37. Quantum Numbers Spin Quantum Number - ms
ms = +1/2 or -1/2
ms = ± 1/2
tells us the spin and orientation of the magnetic field of the electrons
Wolfgang Pauli
Exclusion Principle
No two electrons in an atom can have the same set of 4 quantum numbers.

38. Atomic Orbitals
regions of space where the probability of finding an electron about an atom is highest
described by either
n (1,2,3,4,5,...) or
letters (K,L,M,N,O,...)
s orbitals
spherically symmetric
one s orbital per n level
l = value of ml

39. Atomic Orbitals
s orbitals

40. Atomic Orbitals p orbitals
start with n = 2
3 mutually perpendicular peanut shaped volumes
directed along the axes of a Cartesian coordinate system
3 per n level,
px, py, pz
l = 1
ml = -1,0, values of ml

41. Atomic Orbitals
p orbitals

42. Atomic Orbitals d orbitals
start with n = 3
4 clover leaf shaped and 1 peanut shaped with a doughnut around it
on Cartesian axes and rotated 45o

43. Atomic Orbitals d orbitals
start with n = 3
4 clover leaf shaped and 1 peanut shaped with a doughnut around it
on Cartesian axes and rotated 45o
5 per n level
l = 2
ml = -2,-1,0,+1,+2 5 values of ml

44. Atomic Orbitals
d orbitals

45. Atomic Orbitals f orbitals most complex shaped orbitals
start with n = 4
most complex shaped orbitals
7 per n level, complicated names
l = 3
ml = -3,-2,-1,0,+1,+2, values of ml
important effects in lanthanides & actinides

46. Atomic Orbitals
f orbitals

47. Atomic Orbitals spin effects
every orbital can hold up to two electrons
one spin up ­ one spin down ¯
spin describes the direction of their magnetic field
unpaired electrons have their spins aligned ­­ or ¯¯

48. Atomic Orbitals paired electrons have spins unaligned ­¯
2 electrons in same orbital must be paired
consequence of Pauli Exclusion Principle

49. Atomic Orbitals number of orbitals per n level is given by n2
maximum number of electrons per n level is 2n2

50. Atomic Orbitals Energy Level # of Orbitals Max. # of e- n n2 2n2 1 1 2

51. Electronic Configurations
Aufbau Principle - The electron that distinguishes an element from the previous element enters the lowest energy atomic orbital available.

52. Electronic Configurations

53. Electronic Configurations
Aufbau Principle

54. Electronic Configurations
use mnemonic

55. Electronic Configurations
use periodic chart - best method

56. Electronic Configurations
Write the electronic configuration for EVERY ATOM on the PERIODIC TABLE!!

57. Synthesis Question
What is the atomic number of the element that should theoretically be the noble gas below Rn?
The 6 d’s are completed with element 112 and the 7d’s are completed with element Thus the next noble gas (or perhaps it will be a noble liquid) should be element 118.

58. Group Question
In a universe different from ours, the laws of quantum mechanics are the same as ours with one small change. Electrons in this universe have three spin states, -1, 0, and +1, rather than the two, +1/2 and -1/2, that we have. What two elements in this universe would be the first and second noble gases? (Assume that the elements in this different universe have the same symbols as in ours.)