1. Chapter 7 Periodic Properties of the Elements
Chemistry, The Central Science, 10th edition
Theodore L. Brown; H. Eugene LeMay, Jr.; and Bruce E. Bursten
Chapter 7 Periodic Properties of the Elements
John D. Bookstaver
St. Charles Community College
St. Peters, MO
 2006, Prentice Hall, Inc.

2. November 9
For Monday finish chapter 7 and do all the questions in the worksheet.
Test on chapter 6 and 7 on Tuesday.
Equation book review Chapter 2 and 3

3. Anomalies Cr 3d5 4s1 Mo Cu Ag Au
The energies between 3d and 4s are close, and to have a sublevel half filled or completely filled gives stability to the atom.

4. Paramagnetism
Elements and compounds that have unpaired electrons are attracted to a magnet. The effect is weak but it can be observed.
Liquid O2 attracted to a magnet.

5. Diagmagnetism
Substances with no unpaired electrons experiment a slight repulsion when subjected to a magnetic field.

6. http://books. google. com/books

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

8. Development of Periodic Table
Dmitri Mendeleev and Lothar Meyer independently came to the same conclusion about how elements should be grouped in 1869.

9. PERIODIC ARRANGEMENT
Mendeleev arranged the elements according to their atomic masses.
In 1913 Henry Moseley developed the concept of ATOMIC NUMBERS which lead to the current arrangement according to their ATOMIC NUMBER.

10. 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.

11. 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.

12. Effective Nuclear Charge
The effective nuclear charge is smaller than the actual nuclear charge because the effective nuclear charge Zeff accounts for the repulsion of the electron by the other electrons in the atom.
Zeff < Z

13. Core electrons shield or screen the outer electrons from the attraction of the nucleus.

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

15. 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

16. Examples – Place each group of elements in order of increasing atomic radius:
S, Al, Cl, Mg, Ar, Na
K, Li, Cs, Na, H
Ca, As, F, Rb, O, K, S, Ga

17. Examples – Place each group of elements in order of increasing atomic radius:
S, Al, Cl, Mg, Ar, Na
Ar < Cl < S < Al < Mg < Na
K, Li, Cs, Na, H
H < Li < Na < K < Cs
Ca, F, As, Rb, O, K, S, Ga
F < O < S < As < Ga < Ca < K < Rb

18. Electron Configurations of Ions
Cations: electrons removed from orbital with highest principle quantum number, n, first:
Li (1s2 2s1)  Li+ (1s2)
Fe ([Ar]3d6 4s2)  Fe3+ ([Ar]3d5)
Anions: electrons added to the orbital with highest n:
F (1s2 2s2 2p5)  F- (1s2 2s2 2p6)

19. Write electron configurations for the following ions:
Al3+
S2-
Li+
Br-
Fe2+
Fe3+

20. Write electron configurations for the following ions:
Al3+ 1s22s22p6
S2- [Ne]3s23p6
Li+ 1s2
Br- [Ar]4s23d104p6
Fe2+ [Ar]3d6
Fe3+ [Ar]3d5

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

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

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

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

25. November 10 DO NOW Write the formula for
Potassium Permanganate and Sodium hypochlorite
Write the electronic configuration of both cations. Indicate to what element are each isoelectronic.
Which one has the largest radius?

26. Homework New Atomic Theory Exams Questions
is a file in the John Bowne site. Has similar questions to AP test. Practice them. The more you do them the better you’ll do in the test.
Test Unit 6 and 7 Tuesday

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

28. Examples – Choose the larger species in each case:
Na or Na+
Br or Br-
N or N3-
O- or O2-
Mg2+ or Sr2+
Mg2+ or O2-
Fe2+ or Fe3+

29. Examples – Choose the larger species in each case:
Na or Na+
Br or Br-
N or N3-
O- or O2-
Mg2+ or Sr2+
Mg2+ or O2-
Fe2+ or Fe3+

30. 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.

31. IONIZATION ENERGY
THE GREATER THE IONIZATION ENERGY THE MORE DIFFICULT IT IS TO REMOVE AN ELECTRON!
It requires more energy to remove each successive electron.
When all valence electrons have been removed, the ionization energy takes a quantum leap (removing core electrons require huge amounts of energy).

33. 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.

34. 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.

35. Irregularities N 1402 [He] 2s2 2p3
2p3 this configuration is more stable because half filled orbitals minimize repulsions between electrons O
[He] 2s2 2p4
The decrease is due to repulsions of paired electrons in 2p4

36. Trend across a period IE icreases
The energy needed to remove an electron increases because the effective nuclear charge increases and the atomic radius decreases.
Irregularities
Be B 801
[He] 2s [He] 2s2 2p1

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

38. 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.

39. 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.

40. Examples – Put each set in order of increasing first ionization energy:
P, Cl, Al, Na, S, Mg
Ca, Be, Ba, Mg, Sr
Ca, F, As, Rb, O, K, S, Ga

41. Examples – Put each set in order of increasing first ionization energy:
P, Cl, Al, Na, S, Mg
Ca, Be, Ba, Mg, Sr
Ca, F, As, Rb, O, K, S, Ga
1. Na < Al < Mg < S < P < Cl
2. Ba < Sr < Ca < Mg < Be
3. Rb < K < Ga < Ca < As < S < O < F

42. Electron Affinity
It measures the ease with which and atom gains an electron

43. Electron Affinity
Energy change accompanying addition of electron to gaseous atom:
Cl + e−  Cl−
For most atoms energy is released when an electron is added. The change is exothermic. The more exothermic the most stable the product.

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

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

46. 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.

47. 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.

48. November 11 Do now Find the Oxidation Number for Oxygen in KO2 K2O2
Name the compounds

49. Electronegativity increases
Electronegativity: The ability of an atom in a molecule to attract electrons to itself.
Pauling set electronegativities on a scale from 0.7 (Cs) to 4.0 (F). Values are calculated from ionization energies and electron affinities.
Electronegativity increases
across a period and
Up a group.

50. Electronegativity

51. Examples – put each set in order by increasing electronegativity:
Na, Li, Rb, K, Fr
Cl, Ca, F, P, Mg, S, K

52. Examples – put each set in order by increasing electronegativity:
Na, Li, Rb, K, Fr
Cl, Ca, F, P, Mg, S, K
Fr < Rb < K < Na < Li
K < Ca < Mg < P < S < Cl < F

53. Properties of Metal, Nonmetals, and Metalloids

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

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

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

57. Metallic Trend
Metals react by losing electrons and forming cations. Metallic character increases with the easiness of losing electrons. The trend is like the ionization energy. Decreases across a period and increases down a group.

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

59. 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.

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

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

62. Group Trends

63. Most metal oxides are basic: (NOT A REDOX RX!!!)
Metals
Most metal oxides are basic: (NOT A REDOX RX!!!)
Metal oxide + water  metal hydroxide
Na2O(s) + H2O(l)  2NaOH(aq)
Nonmetals
Nonmetals are more diverse in their behavior than metals.
When nonmetals react with metals, nonmetals tend to gain electrons (reduced):
metal + nonmetal  salt
2Al(s) + 3Br2(l)  2AlBr3(s)

64. nonmetal oxide + water  acid (NON A REDOX RX)
Nonmetals
Most nonmetal oxides are acidic:
nonmetal oxide + water  acid (NON A REDOX RX)
P4O10(s) + H2O(l)  4H3PO4(aq)
Metalloids
Metalloids have properties that are intermediate between metals and nonmetals.
Example: Si has a metallic luster but it is brittle.
Metalloids have found application in the semiconductor industry.

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

66. moviealkalimetals.mp4 Alkali Metals
Found only as compounds in nature. Very reactive!!!
Have low densities and melting points.
Also have low ionization energies.

67. Alkali Metals Their reactions with water are famously exothermic.
They react vigorously with water and have to be stored in mineral oil or kerosene.
2M (s) + 2 H2 O (l)  2MOH (aq) + H2 (g)

68. Alkali Metals Reactions
Alkali metals combine directly with most nonmetals Alkali metals (except Li) react with oxygen to form peroxides. (O2 2- remember ON -1 )
K, Rb, and Cs also form superoxides:
K + O2  KO2 (ON -1/2!!)
Alkali metals produce different oxides when reacting with O2:
4Li(s) + O2(g)  2Li2O(s) (oxide)
2Na(s) + O2(g)  Na2O2(s) (peroxide)
K(s) + O2(g)  KO2(s) (superoxide)

69. With Hydrogen they form HYDRIDES (H-) LiH, KH
With Sulfur they form Sulfides

70. Group Trends for the Active Metals
Group 1: The Alkali Metals
Produce bright colors when placed in flame.
Na line (589 nm): 3p  3s transition
Li line: 2p  2s transition
K line: 4p  4s transition

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

72. 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. They are less reactive than Alkali metals.

73. Flame Test Colors Ca2+ Orange-red Sr2+ Crimson Red Ba2+ Pale Green
They are used in firework!

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

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

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

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

78. 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

79. Group 17: The Halogens
The chemistry of the halogens is dominated by gaining an electron to form an anion:
X2 + 2e-  2X-.
Fluorine is one of the most reactive substances known:
2F2(g) + 2H2O(l)  4HF(aq) + O2(g) H = kJ.
SiO2 (s) + 2F2(g)  SiF4 (g) + O2 (g) H = kJ
Why is it so reactive?
All halogens consist of diatomic molecules, X2.

80. 2NaCl(aq) + 2H2O(l)  2NaOH(aq) + H2(g) + Cl2(g).
Chlorine is the most industrially useful halogen. It is produced by the electrolysis of brine (NaCl):
2NaCl(aq) + 2H2O(l)  2NaOH(aq) + H2(g) + Cl2(g).
The reaction between chorine and water produces hypochlorous acid (HOCl) which disinfects pool water:
Cl2(g) + H2O(l)  HCl(aq) + HOCl(aq).
Halogens react directly with metals to form ionic halides
H2 (g) + X2  2 HX (g)
Hydrogen compounds of the halogens are all strong acids with the exception of HF.
LEARN THESE REACTIONS!!!

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

82. 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.