1. Chapter 7
Atomic Structure

2. Niels Bohr He said the atom was like a solar system.
The electrons were attracted to the nucleus because of opposite charges.
Didn’t fall in to the nucleus because it was moving around.

3. The Bohr Model Doesn’t work. Only works for hydrogen atoms.
Electrons don’t move in circles.
The quantization of energy is right, but not because they are circling like planets.

4. The Bohr Ring Atom
He didn’t know why but only certain energies were allowed.
He called these allowed energies energy levels.
Putting Energy into the atom moved the electron away from the nucleus.
From ground state to excited state.
When it returns to ground state it gives off light of a certain energy.

5. Hydrogen spectrum Developed the quantum model of the hydrogen atom.
Emission spectrum because these are the colors it gives off or emits.
Called a line spectrum.
Developed the quantum model of the hydrogen atom.
There are just a few discrete lines showing
656 nm
434 nm
410 nm
486 nm

6. Spectrum The range of frequencies present in light.
White light has a continuous spectrum.
All the colors are possible.

7. Light
Made up of electromagnetic radiation.

8. Kinds of EM waves There are many
different l (wavelength) and n (Frequency)
Light is only the part our eyes can detect. G a m m a R a d i o R a y s w a v e s

10. Waves Waves have 3 primary characteristics:
1. Wavelength: distance between two peaks in a wave.
2. Frequency: number of waves per second that pass a given point in space.
3. Speed: speed of light is ´ 108 m/s.

11. Parts of a wave Wavelength l
Frequency = number of cycles in one second
Measured in hertz 1 hertz = 1 cycle/second

13. Wavelength and frequency can be interconverted.
c = n l
n = frequency (s-1, Hz, cyc/s, or waves/s )
l = wavelength (m)
c = speed of light (m/s)
in a vacuum is 3.00 x 108 m/s

14. The speed of light
c = ln
What is the wavelength of light with a frequency 5.89 x 105 Hz?
What is the frequency of blue light with a wavelength of 484 nm?

15. 509.34M
6.214

17. The Bohr Ring Atom
n = 4
n = 3
n = 2
n = 1

18. Only certain energies are allowed for the hydrogen atom.
Energy in the atom is quantized.

19. Planck’s Constant
Transfer of energy is quantized, and can only occur in discrete units, called quanta.
DE = change in energy, in J
h = Planck’s constant, 6.63 ´ J s
n = frequency, in s-1
l = wavelength, in m

20. The Bohr Model Calculating the energy of an electrons in an energy level or during a jump
n is the energy level
for each energy level the energy is different
E = x J /n2
Not Hyrogen
Z is the nuclear charge, which is +1 for hydrogen.
E = x J (Z2 / n2 )

21. When you want the change.
DE = Efinal - Einitial

22. Examples When the electron moves from one energy level to another.
Calculate the energy need to move an electron from its 1 to the third energy level.
Calculate the energy released when an electron moves from n= 4 to n=2 in a hydrogen atom.

23. Einstein is next
Said electromagnetic radiation is quantized in particles called photons.
Each photon has energy = hn = hc/l
Combine this with E = mc2
You get the apparent mass of a photon.
m = h / (lc)
De Broglie Equation

24. What is the wavelength?
of an electron with a mass of x kg traveling at x 106 m/s?

25. The Quantum Mechanical Model
A totally new approach.
De Broglie said matter could be like a wave.

26. There are only certain allowed waves.
In the atom there are certain allowed waves called electrons.
1925 Erwin Schroedinger described the wave function of the electron.
A lot of math but what is important are the solution.

27. Schroedinger’s Equation
Solutions to the equation are called orbitals.
These are not Bohr orbits.
Each solution is tied to a certain energy.
These are the energy levels.

28. There is a limit to what we can know
We can’t know how the electron is moving or how it gets from one energy level to another.
The Heisenberg Uncertainty Principle.
There is a limit to how well we can know both the position and the momentum of an object.

29. Quantum Numbers There are many solutions to Schroedinger’s equation
Each solution can be described with quantum numbers that describe some aspect of the solution.
Principal quantum number (n) size and energy of of an orbital.
Has integer values >0

30. Quantum numbers Angular momentum quantum number l .
shape of the orbital
integer values from 0 to n-1
l = 0 is called s can hold 2 e-
l = 1 is called p can hold 6 e-
l =2 is called d can hold 10e-
l =3 is called f can hold 14e-
l =4 is called g can hold 18e-

31. Quantum numbers Magnetic quantum number (m l)
integer values between - l and + l
Orientation of the orbital to each other
Electron spin quantum number (m s)
Can have 2 values.
either +1/2 or -1/2
Spin of an electron

32. Pauli Exclusion Principle
In a given atom no two electrons can have the same set of 4 quantum #s.
Since only 2 values of ms are allowed thefore an orbital can hold no more then two electrons

33. N=3 l =1 ml= -2 Ms =+1/2
2,3,-1,+1/2
2,1,0,-1/2
3,4,-1,+1/2
4,3,-3,+1/2
1,0,0,-1/2

34. Write each set of quantum #s that describe the electrons in an atom of Boron .

35. S orbitals

36. P Orbitals

37. D orbitals

38. F orbitals

39. P orbitals

40. F orbitals

42. Increasing energy He with 2 electrons 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p3d 4d 5d 7p 6d 4f 5f
He with 2 electrons

43. Do Now
Write each set of quantum #s that describe the valence electrons in an atom of Al

44. Aufbau Principle Aufbau is German for building up.
As the protons are added one by one, the electrons fill up hydrogen-like orbitals.
Fill up in order of energy levels.

45. Details
Valence electrons- the electrons in the outermost energy levels (not d).
Core electrons- the inner electrons.
Hund’s Rule- The lowest energy configuration for an atom is the one have the maximum number of of unpaired electrons in the orbital.
C 1s2 2s2 2p2

46. Write the electron configuration for
C, Cu
Identify the following atom.
1s2 2s2 2p5 3s2 3p3

47. Paramagnetic - Attracted by a magnet because of unpaired electrons
Paramagnetic vs. Diamagnetic
S Ca
Paramagnetic - Attracted by a magnet because of unpaired electrons
Diamagnetic - Weakly repelled by magnet because of paired electrons

48. Exceptions Cr = [Ar] 4s1 3d5 Half filled orbitals.
Scientists aren’t sure of why it happens
same for Cu [Ar] 4s1 3d10
Draw the Orbital diagrams following Hund’s rule

49. Ion Configurations
To form cations from elements remove 1 or more e- from subshell of highest n
P [Ne] 3s2 3p e > P3+ [Ne] 3s2 3p0

50. Ion Configurations
For transition metals, remove ns electrons and then (n - 1) electrons.
Fe [Ar] 4s2 3d6
loses 2 electrons ---> Fe2+ [Ar] 4s0 3d6
loses 3 electrons ---> Fe [Ar] 4so 3d5

51. The Periodic Table
Developed independently by German Julius Lothar Meyer and Russian Dmitri Mendeleev (1870”s).
Didn’t know much about atom.
Put in columns by similar properties.
Predicted properties of missing elements.

52. Information Contained in the Periodic Table
Each group member has the same valence electron configuration (these electrons primarily determine an atom’s chemistry).
Certain groups have special names (alkali metals, halogens, etc).
Metals and nonmetals are characterized by their chemical and physical properties.

53. Broad Periodic Table Classifications
Representative Elements (main group): filling s and p orbitals (Na, Al, Ne, O)
Transition Elements: filling d orbitals (Fe, Co, Ni)
Lanthanide and Actinide Series (inner transition elements): filling 4f and 5f orbitals (Eu, Am, Es)

54. General Periodic Trends
Atomic Radii
Ionization energy
Electron Affinity

55. Reasons for General Periodic Trends
Higher Z*.
Larger orbitals.
More sheilding .

56. Shielding
Electrons on the higher energy levels tend to be farther out.
Have to look through the other electrons to see the nucleus.
They are less effected by the nucleus.
lower effective nuclear charge
If shielding were completely effective, Zeff = 1
Why isn’t it?

57. Atomic Radii

58. Ion Sizes
Does the size go up or down when gaining an electron to form an anion?

59. Ion Sizes Forming an anion.
ANIONS are LARGER than the atoms from which they come.
The electron/proton attraction has gone DOWN and so size INCREASES.
Trends in ion sizes are the same as atom sizes.

60. Ion Sizes Does the size go
up or down when losing an electron to form a cation?

61. Ion Sizes . CATIONS are SMALLER than the atoms from which they come.
Forming a cation.
. CATIONS are SMALLER than the atoms from which they come.
The electron/proton attraction has gone UP and so size DECREASES.

62. Trends in Ion Sizes

63. Periodic Trends
Ionization energy the energy required to remove an electron from a gaseous atom
Highest energy electron removed first.
First ionization energy (I1) is that required to remove the first electron.
Second ionization energy (I2) - the second electron
etc. etc.

64. Ionization Energy
IE = energy required to remove an electron from an atom in the gas phase.
Mg (g) kJ ---> Mg+ (g) + e-

65. Ionization Energy
IE = energy required to remove an electron from an atom in the gas phase.
Mg (g) kJ ---> Mg+ (g) + e-
Mg+ (g) kJ ---> Mg2+ (g) + e-

66. Trends in ionization energy
for Mg
I1 = 735 kJ/mole
I2 = 1445 kJ/mole
I3 = 7730 kJ/mole
The effective nuclear charge increases as you remove electrons.
It takes much more energy to remove a core electron than a valence electron because there is less shielding.

67. Across a Period Generally from left to right, I1 increases because
there is a greater nuclear charge with the same shielding.
As you go down a group I1 decreases because electrons are farther away, more shielding

68. The first ionization energy for phosphorous is 1060 kj/mol and sulfur is 1005kj/mol . Why?

69. Which atom has the largest first ionization energy,and which one has the smallest second ionization energy ? Explain your choice . Ne Na Mg

70. Trends in Ionization Energy
1st Ionization energy (kJ/mol)
500
1000
1500
2000
2500 He Ne Ar Kr 1 3 5 7 9 11 13 15 17 19 21 23 25 27 29 31 33 35 H Li Na K
Atomic Number

71. Explain this trend For Al I1 = 580 kJ/mole I2 = 1815 kJ/mole

72. Electron Affinity A few elements GAIN electrons to form anions.
Electron affinity is the energy involved with the addition of an electron in the gaseous state
A(g) + e- ---> A-(g) E.A. = DE

73. Trends in Electron Affinity
Affinity for electron increases across a period (EA becomes more negative).
Affinity decreases down a group (EA becomes less negative).

74. Increasing Periodic Trends
electronegativity, ionization energy, electron affinity
atomic radii
ionization energy,
electron affinity,
& electronegativity
ionic & atomic
radii
Increasing Periodic Trends

75. Penetration effect
The outer energy levels penetrate the inner levels so the shielding of the core electrons is not totally effective.
from most penetration to least penetration the order is
ns > np > nd > nf (within the same energy level).
This is what gives us our order of filling, electrons prefer s and p.

76. Trends in Ionization Energy
IE increases across a period because Z* increases.
Metals lose electrons more easily than nonmetals.
Metals are good reducing agents.
Nonmetals lose electrons with difficulty.

77. Trends in Ionization Energy
IE decreases down a group
Because size increases.
Reducing ability generally increases down the periodic table.

78. Trends in Ionization Energy
1st Ionization energy (kJ/mol)
500
1000
1500
2000
2500 He Ne Ar Kr 1 3 5 7 9 11 13 15 17 19 21 23 25 27 29 31 33 35 H Li Na K
Atomic Number

79. It is not that simple
Zeff changes as you go across a period, so will I1
Half filled and filled orbitals are harder to remove electrons from.
here’s what it looks like.

80. First Ionization energy
Atomic number

81. First Ionization energy
Atomic number

82. First Ionization energy
Atomic number