1. Chapter Seven: Atomic Structure and Periodicity

2. Arrangement of Electrons in Atoms
Reminder: electrons in atoms are arranged as…
SHELLS (n)
SUBSHELLS (l)
ORBITALS (ml)

3. Electrons in Atoms
When n = 1, then l = 0 This shell has a single orbital (1s) to which 2 electrons can be assigned.
When n = 2, then l = 0, 1
2s orbital electrons
three 2p orbitals 6 electrons
TOTAL = electrons

4. Electrons in Shells
The number of shells increases with the shell value (n).
Therefore each higher shell holds more electrons. (more orbitals)

5. Quantum Numbers Electrons in an atom are arranged by: (quantum number)
Principle energy levels
(shells) n
angular energy levels
(sub-shells) l
oriented energy levels
(orbitals) ml
See Chapter 7 Video Presentation Slides 4 & 5
Within an individual orbital, there may only be two electrons differentiated by their spin.
Spin states
(electrons) ms

6. Atomic Subshell Energies & Electron Assignments
Based on theoretical and experimental studies of electron distributions in atoms, chemists have found there are two general rules that help predict these arrangements:
Electrons are assigned to subshells in order of increasing “n + l” value.
2. For two subshells with the same value of “n + l” electrons are assigned first to the subshell of lower n.

7. Single-Electron Atom Energy Levels3p 3d
Energy 2s 2p
In a Hydrogen atom (1–electron) the orbitals of a subshell are equal in energy (degenerate) 1s

8. Multi-Electron Atom Energy Levels2p 3p 3d
For multi-electron atoms, screening results in the 4s-orbital having a lower energy that that of the 3d-orbital.
E = n + l
4s: n + l = = 4
vs.
3d: n + l = = 5

9. Effective Nuclear Charge, Z*
Z* is the net charge experienced by a particular electron in a multi-electron atom resulting from a balance of the attractive force of the nucleus and the repulsive forces of other electrons.
Z* increases across a period owing to incomplete screening by inner electrons
This explains why E(4s electron) < E(3p electron)
Z*  [ Z  (no. inner electrons) ]
Charge felt by 2s electron in: Li Z* = 3  2 = 1
Be Z* = 4  2 = 2
B Z* = 5  2 = and so on!

10. Electron Configurations & the Periodic Table

11. Electron Filling Order

12. Assigning Electrons to Subshells
In the H-atom, all subshells of same n have same energy.
In a multi-electron atom:
subshells increase in energy as value of n + l increases.
for subshells of same n + I, the subshell with lower n is lower in energy.
See Chapter 7 Video Presentation Slide 6

13. Orbital Filling Rules:
Aufbau Principle: Lower energy orbitals fill first.
Hund’s Rule: Degenerate orbitals (those of the same energy) are filled with electrons until all are half filled before pairing up of electrons can occur.
Pauli exclusion principle: Individual orbitals only hold two electrons, and each should have different spin.
“s” orbitals can hold 2 electrons
“p” orbitals hold up to 6 electrons
“d” orbitals can hold up to 10 electrons

14. Orbital Filling: The “Hund’s Rule”
Degenerate orbitals are filled with electrons until all are half-filled before pairing up of electrons can occur.
Consider a set of 2p orbitals:    2p
Electrons fill in this manner

15. Orbital Filling: The “Pauli Principle”
“Pauli exclusion principle”
Individual orbitals only hold two electrons, and each should have opposite spin.
Consider a set of 2p orbitals:       2p
 = spin up
 = spin down
*the convention is to write up, then down.
Electrons fill in this manner

16. Electron Configuration: Orbital Box Notation
Electrons fill the orbitals from lowest to highest energy.
The electron configuration of an atom is the total sum of the electrons from lowest to highest shell.
Example:
Nitrogen:
N has an atomic number of 7,
therefore 7 electrons
Orbitals 1s 2s 2p       
1s2
2s2
2p3
Electron Configuration (spdf) notation:

17. Atomic Electron Configurations

18. Group 4A Atomic number = 6 6 total electrons 1s2 2s2 2p2
Carbon
Group 4A Atomic number = 6 6 total electrons 1s2 2s2 2p2 1s 2s 2p      

19. Electron Configuration in the 3rd period:
Orbital box notation: 1s 2s 2p 3s 3p 1s 2s 2p 3s 3p
This corresponds to the energy level diagram:
See Chapter 7 Video Presentation Slide 7

20. Electron Configuration in the 3rd Period
Orbital box notation: 1s 2s 2p 3s 3p
Aluminum: Al (13 electrons)                      
spdf Electron Configuration  
1s2
2s2
2p6
3s2
3p1  

21. Noble Gas Notation [Ne] Full electron configuration spdf notation
The electron configuration of an element can be represented as a function of the core electrons in terms of a noble gas and the valence electrons.
Full electron configuration spdf notation
1s22s22p63s23p2
Orbital Box Notation 3s 3p
[Ne]    
Noble gas Notation
[Ne] 3s23p2

22. 1s22s2 = [He], 1s22s22p6 = [Ne], 1s22s22p63s23p6 = [Ar]...
Noble Gas Notation
The innermost electrons (core) can be represented by the full shell of noble gas electron configuration:
1s22s2 = [He], 1s22s22p6 = [Ne], 1s22s22p63s23p6 = [Ar]...
The outermost electrons are referred to as the “Valence” electrons”.
Element Full Electron Config. Noble Gas Notation Mg
1s2 2s2 2p6 3s2
[Ne] 3s2

23. Transition Metal
All 4th period and beyond d-block elements have the electron configuration [Ar] nsx (n - 1)dy
Where n is the period and x, y are particular to the element.
Chromium
Iron
Copper

24. Transition Elements Ni:
Electron Configurations are written by shell even though the electrons fill by the periodic table:
Ni:
last electron to fill:
3d8
electron configuration by filling:
1s2
2s2
2p6
3s2
3p6
4s2
3d8
electron configuration by shell: (write this way)
1s2
2s2
2p6
3s2
3p6
3d8
4s2

26. Lanthanides & Actinides
f-block elements: These elements have the configuration [core] nsx (n - 1)dy (n - 2)fz Where n is the period and x, y & z are particular to the element.
Uranium:
[Rn] 7s2 6d1 5f3
Cerium:
[Xe] 6s2 5d1 4f1

27. Some Anomalies
Some irregularities occur when there are enough electrons to half-fill s and d orbitals on a given row.
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28. Some Anomalies For instance, the electron configuration for copper is
[Ar] 4s1 3d10
rather than the expected
[Ar] 4s2 3d9.
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29. Some Anomalies
This occurs because the 4s and 3d orbitals are very close in energy.
These anomalies occur in f-block atoms, as well.
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30. Filling Rules Electrons fill the lowest energy levels first.
Aufbau Principle:
Electrons fill the lowest energy levels first.
Pauli Exclusion Principle:
Electrons can fill two/orbital, providing they have opposite spins ( )
Hund’s Rule:
Orbitals of the same energy level (p,d,f) distribute electrons 1/orbital before adding the second.

31. 1s22s22p63s2 [Kr]5s24d105p5 1s22s22p3 [Rn]7s1 [Ar]4s23d10
Who are they?
1s22s22p63s2
[Kr]5s24d105p5
1s22s22p3
[Rn]7s1
[Ar]4s23d10

32. Who are they? 1s22s22p63s2 Magnesium [Kr]5s24d105p5 Iodine
1s22s22p3 Nitrogen
[Rn]7s1 Francium
[Ar]4s23d10 Zinc

33. What’s wrong?
Mg: [Ar]3s2 Fe: 1s22s22p63s23p64s24d6 Al: [Ne] 3s3

34. Mistakes!
Mg: [Ar]3s2 Fe: 1s22s22p63s23p64s24d6 Al: [Ne] 3s3

35. Corrections
Mg: [Ne]3s2 Fe: 1s22s22p63s23p64s23d6 Al: [Ne] 3s23p1

36. Valence Electrons: Electrons at the highest energy level.
Electrons available to be gained/lost/shared in a chemical reaction
Valence electron configuration: nsxpx (total of 8 valence electrons)
Representation of valence electrons in Lewis dot notation:

37. Formation of ions & Valence electrons
Ions are formed when atoms either:
1. Cation (+): Give up (lose) electrons
Anion (-): Gain electrons
Overall goal: stable noble gas notation
Elements with < 4 valence electrons- form cations
Elements with > 4 valence electrons for anions.
Non-metals with 4 valence electrons- do not form ions
Noble gases (8 valence electrons) – are unreactive

38. Ions & Electron Configuration
Atoms or groups of atoms that carry a charge
Cations- positive charge
Formed when an atom loses electron(s)
Na  Na+ + e-
Na+ : 1s22s22p6 (10 e-) = [Ne]
Anions –negative charge
Formed when an atom gain electrons(s)
F + e-  F-
F- : 1s22s22p6 (10 e-) = [Ne]

39. Isoelectronic series:
Ions that contain the same number of electrons.
Example: Al3+, Mg 2+, Na+, F-, O2-, N3-

40. Transition Metals & Ions
Transition metals: multi-valent ions
(more than one possible charge)
Electrons are removed from the highest quantum number first:
Example: Cu+ & Cu 2+
Cu+: [Ar] 3d10
Cu2+ : [Ar] 3d9

41. Electron Spin & Magnetism
Diamagnetic Substances: Are NOT attracted to a magnetic field
Paramagnetic Substances: ARE attracted to a magnetic field.
Substances with unpaired electrons are paramagnetic.
See Chapter 6 Video Presentation Slide 24

42. Electron Configurations of Ions
Ions with UNPAIRED ELECTRONS are PARAMAGNETIC (attracted to a magnetic field). Ions without UNPAIRED ELECTRONS are DIAMAGNETIC (not attracted to a magnetic field).
Fe3+ ions in Fe2O3 have 5 unpaired electrons.
This makes the sample paramagnetic.

43. Practice: Electron Configuration
Write the following electron configurations
Silicon- orbital notation
Strontium – long (spdf) notation
Bismuth – noble gas notation
Silver – noble gas notation
O 2- - orbital notation
Fe2+ - noble gas notation

44. Practice #2: Electron Configuration
Write the following electron configurations
Chromium- long (spdf) notation
Sulfur – orbital notation
Ag+ – noble gas notation
Americium – noble gas notation
Mg 2+ - orbital notation
Krypton – noble gas notation

45. Development of Periodic Table
Elements in the same group generally have similar chemical properties.
Physical properties are not necessarily similar, however.
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46. Periodic Trends
In this chapter, we will rationalize observed trends in:
Sizes of atoms and ions.
Ionization energy.
Electron affinity.
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47. General Periodic Trends
Atomic and ionic size
Ionization energy
Electron affinity
Higher effective nuclear charge
Electrons held more tightly
See Chapter 7 Video Presentation Slides 8, 9, & 10
Larger orbitals.
Electrons held less
tightly.

48. Effective Nuclear Charge
In a many-electron atom, electrons are both attracted to the nucleus and repelled by other electrons.
The nuclear charge that an electron experiences depends on both factors.
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49. Effective Nuclear Charge
The effective nuclear charge, Zeff, is found this way:
Zeff = Z − S
where Z is the atomic number and S is a screening constant, usually close to the number of inner electrons.
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50. What Is the Size of an Atom?
The bonding atomic radius is defined as one- half of the distance between covalently bonded nuclei.
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51. Sizes of Atoms Bonding atomic radius tends to…
…decrease from left to right across a row
(due to increasing Zeff).
…increase from top to bottom of a column
(due to increasing value of n).
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52. Ionization Energy
The ionization energy is the amount of energy required to remove an electron from the ground state of a gaseous atom or ion.
The first ionization energy is that energy required to remove first electron.
A(g) + energy  A+ + e-
The second ionization energy is that energy required to remove second electron, etc.
A+(g) + energy  A2+ + e-
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53. Ionization Energy
It requires more energy to remove each successive electron.
When all valence electrons have been removed, the ionization energy takes a quantum leap.
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54. Trends in First Ionization Energies
As one goes down a column, less energy is required to remove the first electron.
For atoms in the same group, Zeff is essentially the same, but the valence electrons are farther from the nucleus.
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55. Trends in First Ionization Energies
Generally, as one goes across a row, it gets harder to remove an electron.
As you go from left to right, Zeff increases.
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56. Trends in First Ionization Energies
However, there are two apparent discontinuities in this trend.
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57. Trends in First Ionization Energies
The first occurs between Groups IIA and IIIA.
In this case the electron is removed from a p-orbital rather than an s-orbital.
The electron removed is farther from nucleus.
There is also a small amount of repulsion by the s electrons.
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58. Electron Affinity ( EAH)
Electron affinity is the energy change accompanying the addition of an electron to a gaseous atom:
Cl(g) + e−  Cl−
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59. Trends in Electron Affinity
In general, electron affinity becomes more exothermic as you go from left to right across a row.
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60. Trends in Electron Affinity
There are again, however, two discontinuities in this trend.
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61. Trends in Electron Affinity
The first occurs between Groups 1 and 2.
The added electron must go in a p-orbital, not an s-orbital.
The electron is farther from nucleus and feels repulsion from the s-electrons.
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62. Trends in Electron Affinity
The second occurs between Groups IVA and VA.
Group VA has no empty orbitals.
The extra electron must go into an already occupied orbital, creating repulsion.
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63. Ions: Write the noble gas notation for the following ions: Mg2+ P3-
Fe3+

64. Write the noble gas notation for the following ions:
Mg2+ [He]2s22p6
P3- [Ne] 3s23p6
Fe3+ [Ar] 3d5

65. Sizes of Ions Ionic size depends upon: The nuclear charge.
The number of electrons.
The orbitals in which electrons reside.
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66. Sizes of Ions Cations are smaller than their parent atoms.
The outermost electron is removed and repulsions between electrons are reduced.
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67. Sizes of Ions Anions are larger than their parent atoms.
Electrons are added and repulsions between electrons are increased.
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68. Sizes of Ions Ions increase in size as you go down a column.
This is due to increasing value of n.
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69. Sizes of Ions
In an isoelectronic series, ions have the same number of electrons.
Ionic size decreases with an increasing nuclear charge.
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70. Problem: Rank the following ions in order of decreasing size? Na+,
Mg2+, F–
O2–

71. Problem: Rank the following ions in order of decreasing size? Na+,
Mg2+, F–
O2–
Ion # of protons # of electrons ratio of e/p
Na+
N3-
Mg2+ F–
O2–

72. Problem: Rank the following ions in order of decreasing size? Na+,
Mg2+, F–
O2–
Ion # of protons # of electrons ratio of e/p
Na+
N3-
Mg2+ F–
O2– 11 7 12 9 8

73. Problem: Rank the following ions in order of decreasing size? Na+,
Mg2+, F–
O2–
Ion # of protons # of electrons ratio of e/p
Na+
N3-
Mg2+ F–
O2– 11 7 12 9 8 10

74. Problem: Rank the following ions in order of decreasing size? Na+,
Mg2+, F–
O2–
Ion # of protons # of electrons ratio of e/p
Na+
N3-
Mg2+ 11 7 12 9 8 10
0.909
1.43
0.833
1.11
1.25

75. Na+
N3-
Mg2+ F–
O2–
Ion
0.909
1.43
0.833
1.11
1.25
e/p ratio
Since N3- has the highest ratio of electrons to protons, it must have the largest radius.

76. Since Mg2+ has the lowest, it must have the smallest radius.
Na+
N3-
Mg2+ F–
O2–
Ion
0.909
1.43
0.833
1.11
1.25
e/p ratio
Since N3- has the highest ratio of electrons to protons, it must have the largest radius.
Since Mg2+ has the lowest, it must have the smallest radius.
The rest can be ranked by ratio.

77. Na+
N3-
Mg2+ F–
O2–
Ion
0.909
1.43
0.833
1.11
1.25
e/p ratio
Since N3- has the highest ratio of electrons to protons, it must have the largest radius.
Since Mg2+ has the lowest, it must have the smallest radius.
The rest can be ranked by ratio.
N3-
>
O2–
> F–
>
Na+
>
Mg2+
Decreasing size
Notice that they all have 10 electrons: They are isoelectronic (same electron configuration) as Ne.

78. Summary of Periodic Trends
Moving through the periodic table:
Atomic radii
Ionization Energy
Electron Affinity
Down a group
Increase
Decrease
Becomes less exothermic
Across a Period
Becomes more exothermic