1. PowerPoint to accompany
Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
PowerPoint to accompany
General Chemistry
Third Edition
Chapter 8
Raymond Chang

2. Problems from pages
1-8, 10-14, 18, 20, 22-24, 26, 28, 30, 32-35, 38, 40, 42, 44, 46, 49, 50, 52, 53, 54, 66, 70, 72, 84,

3. EH Assignment
Due April 1
Unit 9
Minimum score
Sec 1
already done
Sec 2 90
Sec 3
Sec 4
Sec 5
Sec 6

4. Test on Chapters 7 and 8 April 6
EH Assignment
Due April 4
Unit 11
Minimum score
Sec 1 90
Sec 2
Sec 3
Sec 4
Sec 5 75
Test on Chapters 7 and 8 April 6

5. Development of the Periodic Table
In 1870 the Russian chemistry Dmitri Mendeleev arranged
the 65 known elements into a periodic table.
The periodic law: when the elements are arranged by atomic
mass similar properties recur periodically.
Mendeleev placed elements with similar properties in the
same column. But he had to make the rows of unequal
length and leave gaps to make similar properties line up.

7. 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

8. Fig. 8.2

9. Valence Electrons 1. Inner (core) electrons are those in the previous
noble gas and any completed transition series.
2. Outer electrons are those in the highest energy
level (highest n value).
3. Valence electrons are the outer electrons for the
representative elements. The number of valence
electrons can easily be determined from the
position in the periodic table.
4. For the transition and inner transition elements the concept of valence electron is not useful.
Examples: sodium, chlorine

10. Electron Configurations of
Metal Ions
All metals lose electrons to form cations. Many of the
metals which are representative elements lose all their
valence electrons to have noble gas configurations.
Transition metal ions rarely attain a noble gas configuration.
They always lose the outer s electrons first. Usually they
lose one or more d electrons as well.
Example: Fe (how many unpaired electrons?)

11. Electron Configurations of
Transition Metal Ions
Predict if the following are paramagnetic or diamagnetic?
Sc Mn Zn2+

12. Electron Configurations of Cations and Anions
Of Representative Elements
Na [Ne]3s1
Na+ [Ne]
Metals lose electrons.
Ca [Ar]4s2
Ca2+ [Ar]
Al [Ne]3s23p1
Al3+ [Ne]
H 1s1
H- 1s2 or [He]
Many nonmetals gain electrons so the 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

13. 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- ?
8.2

14. Trends in Key Periodic Atomic Properties
Atomic Size: There is no real limit to how far an electron
can be away from the nucleus. But, atomic size is a
measurable quantity. In practice we measure the distance
between two identical, adjacent atomic nuclei and divide
distance in half.
Many of the properties of elements are related to the size
of their atoms.

15. Fig. 8.11

16. Fig. 8.4

17. Fig. 8.5

18. Ranking Elements by Size
Problem: Rank the following elements in each
up according to decreasing size
( largest first!):
a) Na, K, Rb b) Sr, In, Rb
c) Cl, Ar, K d) Sr, Ca, Rb

19. 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

20. TA p237

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

22. 8.3

23. Fig. 8.6

24. There is a rough inverse correlation between I1 and atomic radius.
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
There is a rough inverse correlation between I1 and atomic radius.
8.4

25. 8.4 Filled n=1 shell Filled n=2 shell Filled n=3 shell

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

27. Fig. 8.15

28. Problem: Using the Periodic table only, rank the
Ranking Elements by First Ionization Energy
Problem: Using the Periodic table only, rank the
following elements in each of the
following sets in order of increasing I1
a) Ar, Ne, Rn b) At, Bi, Po
c) Be, Na, Mg d) Cl, K, Ar

30. Fig. 8.16

31. Identifying Elements by Its Successive Ionization Energies
Problem: Given the following series of ionization
energies (in kJ/mol) for an element
in period 3, name the element and write
its electron configuration:
I I I I4
, , ,600

32. 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

33. Trends in Metallic Behavior
Fig. 8.19

34. Group 1A Elements (ns1, n  2)
M M+1 + 1e-
2M(s) + 2H2O(l) MOH(aq) + H2(g)
4M(s) + O2(g) M2O(s)
Increasing reactivity
8.6

35. 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

36. Metal oxides tend to be basic.
Non metal oxides tend to be acidic.
Amphoteric oxides exhibit both acidic
and basic properties.

37. The Trends in Acid-Base Behavior of Elemental Oxides