The Bohr Atom

1
The Bohr Atom

• Thompson plum pudding
• Rutherford nucleus in a sea of electrons.
• Bohr planetary model.
• Wave velocity (wavelength)(frequency).
• Review of wave mechanics.

2
The Bohr Atom

• The electromagnetic spectra figure 6.2, p. 132
• The type of wave is determined by its wavelength
  and frequency.
• Speed of light 2.998 x 108 m/s

3
The Bohr Atom

• A light wave has a frequency of 1.74 x 1017
  hertz. What is its wavelength? What type of light
  is it?

• Wavelength 1.72 x 10-9m
• (x ray)

4
Plancks Hypothesis

• Light is given off in bundles of energy called
  quanta.
• E h?
• Energy plancks constant wave frequency
• h 6.626 x 10-34 J/hz

5
Energy Problems

• A photon of light is found to have 3.63 x 10-22
  Joules of energy. What is the wavelength of this
  light wave?

? 5.47 x 10-4 m
6
Energy Problems

• A photon of light is found to have a wavelength
  of 5.00 x 102 nm. What is the energy of a photon
  of this light wave?

E 3.98 x 10-19 J
7
Plancks Hypothesis

• Atomic Spectra produced when an electron moves
  from a higher to lower energy level, giving off
  light in the process.

8
Plancks Hypothesis

• Atomic Spectra produced when an electron moves
  from a higher to lower energy level, giving off
  light in the process.
• ??E Ehi - Elo h? hc/?

9
Plancks Hypothesis

• Ex. For the yellow line in the sodium spectra (?
  589.0 nm), find its frequency, quantum energy,
  and the energy released by one mol of sodium
  electrons.What is the energy difference between
  two energy levels of Na?

10
Plancks Hypothesis

• ????5.090 x 1014s-1
• ?E 3.373 x 10-19 J
• For one mol of electrons ?E 203.1 kJ
• Hence a two energy level difference 203.1
  kJ/mol

11
Bohr Model

• Bohr postulated that an electron moves about the
  nucleus in a circular orbit of a fixed radius.

12
Bohr Model

• The emission spectra of hydrogenhydrogen absorbs
  energy when excited then gives it off when it
  returns to its ground state.

13
Bohr Model

• The ground state of an electron represents its
  lowest orbit. The excited state represents any
  other possible orbit.

14
Bohr Model

• Hydrogen only emits this absorbed energy at
  certain visible wavelengths. Bohr reasoned this
  related to certain allowed electron orbits.

15
Bohr Model

• To calculate the energy of an allowed energy
  level En (-2.180 x 10-18 J)/n2, where n 1,
  2, 3,
• RH Rydberg constant 2.180 x 10-18 J

16
Bohr Model

• In the Bohr atom, calculate the energy released
  as an electron moves from the third to the second
  energy level. What is the wavelength of the
  emitted light?

17
Bohr Model

• E3 -2.422 x 10-19 J E2 -5.450 x 10-19 J
• Ehi - Elo 3.028 x 10-19 J
• ?? hc/?E 6.560 x 10-7 m 656.0 nm (how does
  this compare to the Balmer series?)

18
Modern Atomic Structure

• Bohr emission spectra turned out to be several
  lines at each level, not singular lines.
• De Broglie wavelengths can be predicted based on
  the mass and velocity of a particle.

19
Modern Atomic Structure

• Wave/particle duality.
• Planck - waves can act like particles Ehn
• DeBroglie - hey then particles can act as waves,
  mc2 E hn, l h/mv.

20
Modern Atomic Structure

• Wave/particle duality.
• Experiments can only demonstrate one of these
  qualities at a time.
• Particle behavior photoelectric effect (solar
  powered calculator).

21
Modern Atomic Structure

• Wave/particle duality.
• Wave behavior refraction (changes speed in
  different media), defraction (bends around
  barriers), reflection

22
Modern Atomic Structure

• Heisenberg Uncertainty Principle both the
  momentum and position of a particle can not be
  precisely known at the same time.

23
Modern Atomic Structure

• Therefore, we can only refer to the probability
  of finding an electron in a region we cannot
  specify the path.

24
Modern Atomic Structure

• Schrodinger wave equations (y2) can be used to
  predict the region of probability for locating an
  electron.

25
Modern Atomic Structure

• An Electron moves at high velocities usually on
  the surface of this region.
• An electron effectively fills the surface. (fan
  analogy)

26
Modern Atomic Structure

• Quantum numbers are used to describe the location
  of electrons in atoms.
• Importance model of atoms and bonding theory.

27
Modern Atomic Structure

• Principle Quantum Number, n energy level.
• The higher the number the larger the region.
• Corresponds to the periodic table.

28
Modern Atomic Structure

• Principle energy level the value of n is the
  main factor that determines the energy of an
  electron and its distance from the nucleus.
• Maximum electron capacity of a level 2n2

29
Modern Atomic Structure

• Second Quantum Number, l refers to energy
  sublevels.
• The number of sublevels equals the principal
  quantum number.

30
Modern Atomic Structure

• Sublevels do not have the same energy.
• Sublevels from one principal level can overlap
  sublevels from another. Figure 6.7, p. 151

31
Modern Atomic Structure

• 3rd Quantum Number, m refers to the orientation
  of the suborbital.
• s,p,d and f orbitals.
• Degenerate orbitals (geometries and orientations
  - capacities)

32
Modern Atomic Structure

• Each orbital has the capacity of two electrons.
  The s orbitals are spherically symmetric about
  the nucleus p orbitals are dumbell shaped and at
  right angles to each other.

33
Modern Atomic Structure

• 4th Quantum Number refers to the spin.
• ms 1/2 or -1/2

34
Modern Atomic Structure

• Pauli Exclusion Principle each electron can be
  described by a unique set of 4 quantum numbers.

35
Modern Atomic Structure

• n primary energy level
• l sublevels
• 0 s-orbital
• 1 p-orbital
• 2 d-orbital.

36
Modern Atomic Structure

• m orientation of the orbital ( -l to l)
• ex. p-orbital
• px -1
• py 0
• pz 1

37
Modern Atomic Structure

• spin 1/2 or -1/2
• ex
• 1st electron 1 0 0 1/2.
• 2nd electron 1 0 0 -1/2
• 3rd electron 2 0 0 1/2

38
Modern Atomic Structure

• Hunds Rule electrons fill unoccupied degenerate
  orbitals before pairing.
• Find the quantum numbers for the 5th and 7th
  electrons.

39
Modern Atomic Structure

• Basic electron configurations.
• Orbital diagrams.
• Core configurations.

40
Modern Atomic Structure

• Note 1. 2 e- in an orbital have opposed spins.
• 2. When several orbitals of the same sublevel are
  available, electrons enter one at a time with
  parallel spins.

41
Modern Atomic Structure

• Pneumonic device for remembering the filling
  order.

42
Modern Atomic Structure

• Use the periodic table to locate the outermost
  electrons of an atom.
• S-block and p-block elements match the row.

43
The Transition Metals

• d-block elements (groups 3-12).
• Energy sublevel overlap ex - 4s vs. 3d
• Multiple valence

44
The Periodic Table

• Brightly colored compounds and solutions.

45
The Periodic Table

• Lanthanoids elements 57-70, begins 4f block.
• Actinoids elements 89-102, begins 5f block.
• Rare earth metals.

46
The Periodic Table

• Nature tends towards stability.
• Atoms seek bonding situations that result in
  stable electron configurations (ex. Share
  electrons).

47
The Periodic Table

• Octet rule eight electrons in the outer level (s
  ps?) render an atom unreactive.
• Atoms seek to lose, gain, or share electrons to
  seek a stable octet of electrons.

48
The Periodic Table

• An atom having a filled or half-filled sublevel
  is slightly more stable than an atom without.
• Full sublevels are more stable than half-filled.

49
The Periodic Table

• Full outer levels are more stable than full
  sublevels.

50
The Periodic Table

• Electron promotion an electron can be promoted
  to a slightly higher sublevel in order to produce
  a full or half filled sublevel.
• Ex - Cr and Cu (p. 154)

51
The Periodic Table

• Groups families columns of related elements.
  Each family has a similar outer electron
  configuration. Ex. s2 and p5 elements.

52
The Periodic Table

• The properties are predictable and repeat
  themselves. They are based on electronic
  configuration.

53
Atomic Radii

• The distance from the nucleus to the outermost
  orbital.
• Increases as you move down a group increased
  principle energy level.

54
Atomic Radii

• Decreases as you move across a period.
• Z effective - the larger the charge of a nucleus,
  the greater the pull on the electrons.

55
Atomic Radii

• Z effective - this greater attractive force pulls
  the electrons slightly closer to the nucleus and
  accounts for the trend in atomic radii across a
  period.

56
Radii of Ions

• Ions charged particles which are the result of
  adding or subtracting electrons from a neutral
  atom.

57
Radii of Ions

• Cations ions with a positive charge (metals).
• Anions ions with a negative charge (non-metals).

58
Radii of Ions

• Atoms will add or subtract electrons to complete
  their outermost energy level (they seek
  stability a full octet of electrons).

59
Radii of Ions

• Nature does whatever is easiest.
• Ex. It is easier for potassium to lose 1 electron
  rather than gain 7 to complete its octet.

60
Radii of Ions

• Ionic radii is based on whether an atom will add
  or subtract electrons when it ionizes.

61
Radii of Ions

• Cations are smaller than their neutral atoms.
• Anions are larger than their neutral atoms.

62
Radii of Ions

• Ionic radii increase down a group and decrease
  across a period. p. 252 (notice the group 18
  elements)

63
Oxidation Numbers

• Predicting oxidation states for the main group
  elements (groups 1-2, 13-18) this will be based
  on the tendency of the group towards stability.

64
Oxidation Numbers

• Predicting oxidation states for the transition
  elements - theory vs reality
• Ex. Zn and Ag

65
Oxidation Numbers

• Transition elements
• 1. Can lose one or both of its two s-shell
  electrons first.
• 2. Can lose each individual d-shell electrons
  only after the outer s-shell is empty

66
Oxidation Numbers

• Transition elements
• 3. Will not lose d-electrons if that shell is
  half-filled.
• Ex. Scandium (2, 3)
• vs Titanium (2, 3, 4)
• Vs Copper (1, 2)

67
Ionization Energy

• Ionization energy the energy required to remove
  one electron from a gaseous atom.

68
Ionization Energy

• Trends decreases down a group, increases across
  a period ( takes more energy to remove an
  electron from an element that usually forms an
  anion).

69
Ionization Energy

• Factors affecting ionization energy
• Nuclear charge
• Shielding effect
• Radius
• Sublevel

70
Ionization Energy

• Nuclear charge the larger the Z effective, the
  greater the force attracting the electrons
  greater ionization energy.

71
Ionization Energy

• Shielding effect - core electrons shield the
  attractive nuclear force from outer electrons,
  lessening ionization energy. Electron/electron
  repulsion can increase the effect.

72
Ionization Energy

• Radius- the greater the atomic radii, the lower
  the ionization energy.
• Sublevels- removing an electron from a
  half-filled or full sublevel requires more energy.

73
Electron Affinity

• Electron Affinity an atoms ability to attract
  additional electrons.

74
Electron Affinity

• Metals have low electron affinities. Non-metals
  have high electron affinities. The trend
  increase across a period and decreases down a
  group. Why?

75
Ionization Energy

• Multiple ionization energies the second and
  third ionization energies can give clues as to
  atomic structure. Ex. Al vs Mg vs Na