1. Periodic Trends
Chapter 7

2. Development of Periodic Table
Elements in the same group generally have similar chemical properties.
Properties are not identical, however.

3. Development of Periodic Table
Mendeleev, for instance, predicted the discovery of germanium (which he called eka-silicon) as an element with an atomic weight between that of zinc and arsenic, but with chemical properties similar to those of silicon.

4. What are the electronic Configurations of the “always” ions?
What does Isoelectronic mean?

5. Electron Configurations of Cations and Anions
Of Representative Elements
Na [Ne]3s1
Na+ [Ne]
Atoms lose electrons so that cation has a noble-gas outer electron configuration.
Ca [Ar]4s2
Ca2+ [Ar]
Al [Ne]3s23p1
Al3+ [Ne]
H 1s1
H- 1s2 or [He]
Atoms gain electrons so that anion has a noble-gas outer electron configuration.
F 1s22s22p5
F- 1s22s22p6 or [Ne]
O 1s22s22p4
O2- 1s22s22p6 or [Ne]
N 1s22s22p3
N3- 1s22s22p6 or [Ne]
8.2

6. Ground State Electron Configurations of the Elements
ns2np6
Ground State Electron Configurations of the Elements
ns1
ns2np1
ns2np2
ns2np3
ns2np4
ns2np5
ns2
d10 d1 d5 4f 5f
8.2

7. Cations and Anions Of Representative Elements+1 +2 +3 -3 -2 -1
8.2

8. Na+, Al3+, F-, O2-, and N3- are all isoelectronic with Ne
Na+: [Ne]
Al3+: [Ne]
F-: 1s22s22p6 or [Ne]
O2-: 1s22s22p6 or [Ne]
N3-: 1s22s22p6 or [Ne]
Na+, Al3+, F-, O2-, and N3- are all isoelectronic with Ne
What neutral atom is isoelectronic with H- ?
H-: 1s2
same electron configuration as He
8.2

9. Who is the biggest? Who is the smallest? Why?
Atomic Radii Trends
Who is the biggest?
Who is the smallest?
Why?

10. Atomic Radii
8.3

11. 8.3

12. 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.

13. 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.

14. 0 < s < Z (s = shielding constant)
Effective nuclear charge (Zeff) is the “positive charge” felt by an electron.
Zeff = Z - s
0 < s < Z (s = shielding constant)
Zeff  Z – number of inner or core electrons
Zeff
Core Z
Radius Na Mg Al Si 11 12 13 14 10 1 2 3 4
186
160
143
132
Within a Period
as Zeff increases
radius decreases
8.3

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

16. Sizes of Ions Cations are smaller than their parent atoms.
The outermost electron is removed and repulsions are reduced.

17. Sizes of Ions Anions are larger than their parent atoms.
Electrons are added and repulsions are increased.

18. Sizes of Ions Ions increase in size as you go down a column.
Due to increasing value of n.

19. Sizes of Ions
In an isoelectronic series, ions have the same number of electrons.
Ionic size decreases with an increasing nuclear charge.

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

21. 8.3

22. Do Metals lose or gain electrons?
What do we call the energy required to do this process?

23. I1 first ionization energy
Ionization energy is the minimum energy (kJ/mol) required to remove an electron from a gaseous atom in its ground state.
I1 + X (g) X+(g) + e-
I1 first ionization energy
I2 + X+(g) X2+(g) + e-
I2 second ionization energy
I3 + X2+(g) X3+(g) + e-
I3 third ionization energy
I1 < I2 < I3
8.4

24. 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.

25. The next slide has patterns.
Find them.

26. 8.4 Filled n=1 shell Filled n=2 shell Filled n=3 shell

27. 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.

28. General Trend in First Ionization Energies
Increasing First Ionization Energy
Increasing First Ionization Energy
8.4

29. What implications can we make from these trends?
What is a stable electronic configuration?
Which electron is removed from a transition element first?
Other questions answered?

30. 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
Some repulsion by s electrons.

31. 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.

32. Electron Configurations of Cations of Transition Metals
When a cation is formed from an atom, electrons are removed from the highest n orbital! Do not just remove electrons from the “end” of the electronic configuration!
Why is this so?
Fe: [Ar]4s23d6
Mn: [Ar]4s23d5
Fe2+: [Ar]4s03d6 so … [Ar]3d6
Mn2+: [Ar]4s03d5 so … [Ar]3d5
Fe3+: [Ar]4s03d5 so … [Ar]3d5
8.2

33. What is the energy called for adding an electron to an atom?
What are the trends?
Why?

34. Electron affinity is the negative of the energy change that occurs when an electron is accepted by an atom in the gaseous state to form an anion.
X (g) + e X-(g)
F (g) + e F-(g)
DH = -328 kJ/mol
EA = +328 kJ/mol
O (g) + e O-(g)
DH = -141 kJ/mol
EA = +141 kJ/mol
8.5

35. 8.5

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

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

38. 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.

39. 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.

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

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

42. Metals Compounds formed between metals and nonmetals tend to be ionic.
Metal oxides tend to be basic.

43. 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.

44. Nonmetals
Substances containing only nonmetals are molecular compounds.
Most nonmetal oxides are acidic.

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

46. Group Trends

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

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

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

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

51. Group 1A Elements (ns1, n  2)
M M+1 + 1e-
2M(s) + 2H2O(l) MOH(aq) + H2(g)
4M(s) + O2(g) M2O(s)
Li and Na make oxides
M(s) + O2(g) MO2 (s)
K, Rb, Cs make super oxides
Increasing reactivity
8.6

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

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

54. Group 2A Elements (ns2, n  2)
M M+2 + 2e-
Be(s) + 2H2O(l) No Reaction
Mg(s) + 2H2O(g) Mg(OH)2(aq) + H2(g)
M(s) + 2H2O(l) M(OH)2(aq) + H2(g) M = Ca, Sr, or Ba
Increasing reactivity
8.6

55. Group 3A Elements (ns2np1, n  2)
4Al(s) + 3O2(g) Al2O3(s)
2Al(s) + 6H+(aq) Al3+(aq) + 3H2(g)
8.6

56. Group 6A Oxygen, sulfur, and selenium are nonmetals.
Tellurium is a metalloid.
The radioactive polonium is a metal.

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

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

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

60. 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

61. Group 7A Elements (ns1np5, n  2)
X + 1e X-1
X2(g) + H2(g) HX(g)
Increasing reactivity
8.6

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

63. 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.