1. Periodic Variation in Physical Properties
Sizes of atoms and ions
Ionization energy
Electron affinity
Metallic properties

2. Properties and Electronic Structure
Properties depend on
valence electrons
Effective nuclear charge (net charge on an electron)
Protons-electrons attraction (increases)
Electron-electron repulsion (decreases)
Penetration of the orbitals (greater penetration – increases the Zeff)
The shielding effect
Core electrons
Size of the atom

3. Effective Nuclear Charge
Effective Nuclear charge, Zeff: the actual nuclear charge a valence electron experiences.
In a many-electron atom,
the Zeff depends on two
factors:
1. Attraction between electrons and nucleus
2. Repulsion between electrons in orbitals

4. Effective Nuclear Charge
In SWE, electron is treated individually in a field of net nuclear charge that determines the Zeff..
The effective nuclear charge, Zeff, is found:
Zeff = Z − σ
Z =atomic number
σ =screening constant
~ # number of inner electrons, but not equal.

5. Effective Nuclear Charge (Zeff)
Proton-electron attraction increases the Zeff.
It depends on:
Size of the atom: Coulomb’s Law
Penetration of orbitals: (10% probability being next to the nucleus)

6. Effective Nuclear Charge (Zeff)
Electron-electron repulsion reduces the Zeff
Smaller Zeff on valence electron, easier to remove the electron from the atom

7. Effective Nuclear Charge on Valence Electrons
Expected Nuclear Charge by the valence electron in
Li atom; Be atom; B atom …..
But: Li. Zeff on the 3rd electron Why?
Zeff = Z – (shielding by other electrons)
Zeff = Z – σ

8. Effective Nuclear Charge on 2s Electron in Li

9. Shielding and (Zeff) in a Group
The electron on the outside energy level has to look through all the other energy levels to see the nucleus: it is shielded from the nucleus by all the inner electrons
Less attraction between the valence electron and the nucleus
Lower effective nuclear charge
If shielding were completely effective, Zeff = 1
Why isn’t it? 24

10. Zeff in a Group Electrons enter new shell (energy level)
Size of atom increases
Zeff decreases: decreased force of attraction (Coulomb’s Law – inversely proportional to r2)
In a Group: Zeff decreases; shielding Increases.

11. Shielding in a Group
The more positive the nucleus (higher n, higher atomic number), the closer the inner orbitals to the nucleus
provides for better shielding of the outer electrons (easier to remove electron in Na than in Li or H) by the core electrons.
In lithium 1s orbital is the same shape as a hydrogen 1s orbital, but it is smaller because the electron is more strongly attracted to the nucleus.
The sodium 1s is even smaller.

12. Shielding: Electron Repulsion and Orbital Energy
Extent of shielding depends on:
1. Core e- provide more shielding than valence electrons
Electrons in the same shell ( same n), extend of shielding depends on l (penetration effects);
s > p > d > f
Correlates with degree of penetration.
3. Electrons in the same sub-shell (same value of n and l) do not effectively shield one another. Electrons are in the degenerate orbitals

13. Shielding in Period
The electron on the outside energy level has to look through all the other energy levels to see the nucleus
A second electron in the same energy level has almost the same shielding.
The third one has almost the same shielding
Zeff (effective nuclear charge) increases as atomic number increases in the Period.
Why? 29

14. Shielding Effects: Examples
First Ionization energies in kJ/mol
He atom (1s2 + 2 protons): 2372
He+ ion: (1s1 + 2 protons): ( it is not 2X 2372, but larger))
In He, the second electron repels the first, making it easier to remove (one electron shields the other from the full effect of the nucleus); not in He+1 (there is only one electron)
In Li (1s22s1): 520; Li+2 (1s1): 2954
In Li: we have two 1s electrons. They shield very effectively the electron in 2s and smaller radius, stronger attraction

15. Sizes of Atoms
The bonding atomic radius is defined as one-half of the distance between covalently bonded nuclei.

16. Atomic Radii: Covalent Compounds

17. Atomic Size in a Group (shielding dominates)
Each new member has one more level of inner electrons
Inner electrons shield the outer electrons very effectively
Zeff increases very slightly (more protons)
Atoms get larges, as n-increases
Atoms radii increase in a group

18. Atomic Size in a Period: Zeff Dominates
Electrons added to the same shell
shielding by inner electrons changes very slightly, if at all.
Outer electrons shield each other poorly
Zeff rises significantly, electrons pulled closer to nucleus
Atomic radius decreases in a period.
Na Mg Al C Si P Cl Ar

19. Atomic Radii of the Main Group Elements35

20. Overall Atomic Radius (nm) Atomic Number Rb K Na Li Kr Ar Ne H 10 GaZn V Ar Ne H 10
Atomic Number

21. Radii in Transition Elements
Size shrinks for the first two to three members because of increased nuclear charge
After: the size remains relatively constant as repulsion of d-electrons (increased radius) counteracts the increase in Zeff.
d-electrons shield very well, but p-orbital penetrates much more than d-orbital: thus Ga (135 pm) is much smaller than Ca (197 pm)
Another anomaly: 13Al (143 pm) versus 31Ga (135 pm). Filling the d-orbitals causes major contraction.

22. Atomic Radius: Examples
Using only the periodic table, rank each set of main-group elements in order of decreasing atomic size:
A) Ca, Mg, Sr
B) K, Ga, Ca
C) Br, Rb, Kr
D) Sr, Ca, Rb

23. Order the following according to increasing atomic radius.
Ge Si Se Cl
Ge < Si < Se < Cl
Se < Si < Ge < Cl
3. Si < Cl < Ge < Se
4. Cl < Si < Se < Ge
5. Si < Ge < Se < Cl

24. Ionic Radius

25. Ionic Radius Ionic radius: the radius of a cation or an anion.
Determines the physical and chemical properties of an ionic compound such as
Crystal structure
Melting point
solubility

26. Cation is always smaller than atom from which it is formed.
Anion is always larger than atom from which it is formed.
8.3

27. 8.3

28. Sizes of Ions Ionic size depends upon: Nuclear charge.
Number of electrons.
Orbitals in which electrons reside.

29. Sizes of Ions
In an isoelectronic series, ions have the same number of electrons.
Ionic size decreases with an increasing nuclear charge.
All have the configuration 1s12s22p6 (10 electrons)

30. Radii of Atoms and Corresponding Ions

31. Ionic Radius: Examples
In each of the following pairs, indicate which one of the two species is larger:
A) N3- or F-
B) Mg+2 or Ca+2
C) Fe+2 or Fe+3
Explain your choice.

32. Order the following according to increasing atomic/ionic radius.
N3- Li+ C O2-
C < Li+ < O2- < N3-
N3- < O2- < C < Li+
3. Li+ < C < N3- < O2-
4. Li+ < C < N3- < O2-
5. Li+ < C < O2- < N3-

33. Ionization Energy
Amount of energy required to remove an electron from the ground state of a gaseous atom or ion.
First ionization energy is that energy required to remove first electron.
Second ionization energy is that energy required to remove second electron, etc.
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34. 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.

35. 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.

36. Filled n=1 shell
Filled n=2 shell
Filled n=3 shell
Filled n=4 shell
Filled n=5 shell

37. 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, shielding is almost the same.

38. Trends in First Ionization Energies
However, there are two apparent discontinuities in this trend.

39. Trends in First Ionization Energies
The first occurs between Groups IIA and IIIA.
Electron removed from p-orbital rather than s-orbital
Electron farther from nucleus
Small amount of repulsion by s electrons.

40. Trends in First Ionization Energies
The second occurs between Groups VA and VIA.
Electron removed comes from doubly occupied orbital.
Repulsion from other electron in orbital helps in its removal.

41. Trends in Successive IE
The effective nuclear charge increases as you remove electrons. Thus IE3>IE2>IE1
Big jump after all outer electrons removed.
It takes much more energy to remove a core electron than a valence electron because there is less shielding, smaller size (energy shell removed)
greater effective nuclear charge.

42. Successive Ionization Energies (kJ/mol)
Na Mg Al Si P S Cl Ar IE IE IE IE IE IE IE

43. Explain the trend in IE For Mg For Al IE1 = 735 kJ/mole

44. Which will have the highest ionization energy?Al Si

45. Which will be the largest?
I = ionization energy
1. I1 of Na
2. I2 of Na
3. I1 of Mg
4. I2 of Mg
5. I3 of Mg #5

46. Electron Affinity
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47. Electron Affinity, kJ/mol
Electron affinity: the energy change that occurs when an electron is accepted by an atom in gaseous state.
X(g) + e- → X-(g)
A large negative value indicates a strong attraction between the atom and the added electron
Cl(g) + e- → Cl-(g) ΔE = -349 kJ/mol
A positive value indicates the addition of electron is unfavorable
Ne(g) + e- → Ne- ΔE = 29 kJ/mol

48. Trends in Electron Affinity
In general, electron affinity becomes more exothermic as you go from left to right across a row.

49. Trends in Electron Affinity
There are again, however, two discontinuities in this trend.

50. Trends in Electron Affinity
The first occurs between Groups IA and IIA.
Added electron must go in p-orbital, not s-orbital.
Electron is farther from nucleus and feels repulsion from s-electrons.

51. Trends in Electron Affinity
The second occurs between Groups IVA and VA.
Group VA has no empty orbitals.
Extra electron must go into occupied orbital, creating repulsion.

52. Summary
Reactive nonmetals (Groups 6A and 7A): have high ionization energies and high negative electron affinities. Gain electrons easily, lose electrons with difficulties.
Reactive metals (Groups 1A and 2A) have low ionization energies and slightly negative (exothermic) electron affinities. Lose electrons lightly and gain electrons with difficulties.
Noble Gases: very high IE and slightly positive electron affinities. Do not lose or gain electrons easily.
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53. Metallic Character Metals Metalloids Nonmetals malleable & ductile
shiny, lustrous
conduct heat and electricity
most oxides basic and ionic
form cations in solution
lose electrons in reactions - oxidized
Metalloids
Also known as semi-metals
Show some metal and some nonmetal properties
Nonmetals
brittle in solid state
dull
electrical and thermal insulators
most oxides are acidic and molecular
form anions and polyatomic anions
gain electrons in reactions - reduced

54. General Trends in Chemical Reactivity: DIAGONAL RULE
First member in each group differs from the rest of the group
It resembles element to its right and next period (Li-Mg; Be-Al, B-Si). Has to do with size. Called diagonal relationship.

55. Properties of Metal, Nonmetals, and Metalloids

56. Metals versus Nonmetals
Differences between metals and nonmetals tend to revolve around these properties.

57. Metallic Character Metals Lose electrons to become positive ions
Low ionization energy
Low electron affinity
Good reducers
Nonmetals
High IE
Low EA
Gain (to become negative ions)
or share electrons to form covalent compounds
Oxidizers

58. Metals versus Nonmetals
Metals tend to form cations.
Nonmetals tend to form anions.

59. Metals
Tend to be lustrous, malleable, ductile, and good conductors of heat and electricity.

60. Metals Compounds formed between metals and nonmetals tend to be ionic.
Metal oxides tend to be basic.
Form bases when reacted with water.

61. Nonmetals
Dull, brittle substances that are poor conductors of heat and electricity.
Tend to gain electrons in reactions with metals to acquire noble gas configuration.

62. Nonmetals
Substances containing only nonmetals are molecular compounds.
Most nonmetal oxides are acidic.
Form acids with water.

63. Metalloids Have some characteristics of metals, some of nonmetals.
For instance, silicon looks shiny, but is brittle and fairly poor conductor.

64. Group Trends

65. Alkali Metals Soft, metallic solids.
Name comes from Arabic word for ashes.

66. Alkali Metals Found only as compounds in nature.
Have low densities and melting points.
Also have low ionization energies.

67. Alkali Metals
Their reactions with water are famously exothermic.

68. Alkali Metals
Alkali metals (except Li) react with oxygen to form peroxides.
K, Rb, and Cs also form superoxides:
K + O2  KO2
Produce bright colors when placed in flame.

69. Alkaline Earth Metals
Have higher densities and melting points than alkali metals.
Have low ionization energies, but not as low as alkali metals.

70. Alkaline Earth Metals
Be does not react with water, Mg reacts only with steam, but others react readily with water.
All react with acids
Reactivity tends to increase as go down group.

71. Group 3A Elements (ns2np1, n≥2)
B is a metalloid, does not react with oxygen or water; does not form binary ionic compounds
Rest of elements mostly behave as metals
2Al(s) + 6H+(aq) → 2 Al+3(aq) + 3H2(g)

72. Group 4A/5A Elements Group VA: Group 5A: Carbon: nonmetal
Silicon and germanium: metalloids
Rest (Pb and Sn) metallic: react with acids, but not with water.
Variable oxidation states for Pb and Sn.
Group 5A:
N, P, nonmetals; As, and Sb metalloids
Rest are metals. P exists as P4. N2 forms many oxides, P only 2.

73. Group 6A Oxygen, sulfur, and selenium are nonmetals.
Oxygen – strong oxidizer.
Tellurium is a metalloid.
The radioactive polonium is a metal.

74. Oxygen Two allotropes: Three anions:
O3, ozone
Three anions:
O2−, oxide
O22−, peroxide
O21−, superoxide
Tends to take electrons from other elements (oxidation)

75. Sulfur Weaker oxidizing agent than oxygen.
Most stable allotrope is S8, a ringed molecule.

76. Group VIIA: Halogens Prototypical nonmetals
Name comes from the Greek halos and gennao: “salt formers”

77. Group VIIA: Halogens Large, negative electron affinities
Therefore, tend to oxidize other elements easily
React directly with metals to form metal halides
Chlorine added to water supplies to serve as disinfectant

78. Group VIIIA: Noble Gases
Astronomical ionization energies
Positive electron affinities
Therefore, relatively unreactive
Monatomic gases

79. Group VIIIA: Noble Gases
Xe forms three compounds:
XeF2
XeF4 (at right)
XeF6
Kr forms only one stable compound:
KrF2
The unstable HArF was synthesized in 2000.

80. Hydrogen: in a class by itself
Acts as a metal: hydrated ion in solution
Acts as a nonmetal: forms hydrides
NaH(s) + H2O(l) → 2NaOH(aq) + H2(g)
NaH(s) + H2O(l) → 2Na+(aq) + OH-(aq) + H2(g)
Most important reaction:
2H2(g) + O2(g) → 2H2O(l)

81. Which will produce a basic solution in water?
1. CO2
2. P2O5
3. BaO
4. XeO3
5. SO2

82. Periodic Trends Interactive
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83. The End