1. Chapter 7 Periodic Properties of the Elements
熊同銘

2. Contents Development of the PeriodicTable Effective Nuclear Charge
Sizes of Atoms and Ions
Ionization Energy
Electron Affinity
Metals, Nonmetals, and Metalloids
Trends for Group 1A and Group 2A Metals
Trends for Selected Nonmetals

3. 1. Development of the Periodic Table
Mendeleev
Ordered elements by atomic mass and saw a repeating pattern of properties
Put elements with similar properties in the same column and used pattern to predict properties of undiscovered elements such as eka-aluminum (Gallium) and eka-silicon (Germanium).

4. Modern periodic table
Ordered elements by atomic number. The number of protons was considered the basis for the periodic property of elements.
For example, the atomic weight of Ar is greater than that of K.
Discovering the elements:

5. 2. Effective Nuclear Charge
Effective nuclear charge (Zeff): The net positive charge experienced by an electron in a many-electron atom; this charge is not the full nuclear charge because there is some shielding of the nucleus by the other electrons in the atom.
Electrons are both attracted to the nucleus and repelled by other electrons.

6. The simplest method for Estimating Zeff of the valence electron
Zeff = Z – S [7.1]
Zeff: effective nuclear charge
Z: actual nuclear charge (atomic number)
S: screening constant (numbers of core electrons)
Examples for representative elements:
11Na: 1s22s22p63s1 Zeff = 11 – 10 = 1
12Mg: 1s22s22p63s2 Zeff = 12 – 10 = 2
**For same period: Z increase Zeff increase
(B) Examples for transition elements:
26Fe: [Ar]3d64s2 Zeff = 26 – 24 = 2
27Co: [Ar]3d74s2 Zeff = 27 – 25 = 2
28Ni: [Ar]3d84s2 Zeff = 28 – 26 = 2
**For 1st transition elements, Zeff are about the same

7. The advanced method for Estimating Zeff of the valence electron
11Na for example, because there is a small probability that the 3s electron is close to the nucleus, the value of S in Equation 7.1 changes from 10 to 8.5. The actual value of Zeff for the 3s electron in Na is Zeff = 2.5+.

8. *Zeff increases across a period. V
Variations in effective nuclear charge for period 2 and period 3 elements.
*Zeff increases across a period. V V

9. 3. Sizes of Atoms and Ions
Atomic Radius: The average radius of an atom based on measuring large numbers of elements and compounds.

10. Trends in bonding atomic radii

11. *****
Continued
Within each group, bonding atomic radius tends to increase from top to bottom because the principal quantum number (n) of the electrons in the outermost principal energy level increases.
Within each period, bonding atomic radius tends to decrease from left to right because the effective nuclear charge (Zeff) experienced by the electrons in the outermost principal energy level increases.
The radii of transition elements vary slightly across each row because the Zeff are about the same.

12. Sizes of Ions Determined by interatomic distances in ionic compounds
Ionic size depends on
the nuclear charge.
the number of electrons.
the orbitals in which electrons reside.

13. Continued
Cations are smaller than their parent atoms:
The outermost electron is removed and repulsions between electrons are reduced.
Anions are larger than their parent atoms:
Electrons are added and repulsions between electrons are increased.

14. Isoelectronic Series
In an isoelectronic series, ions have the same number of electrons.
Ionic size decreases with an increasing nuclear charge.
An Isoelectronic Series (10 electrons)
Note increasing nuclear charge with decreasing ionic radius as atomic number increases
O2– F–
Na+
Mg2+
Al3+
1.26 Å
1.19 Å
1.16 Å
0.86 Å
0.68 Å

15. 4. Ionization Energy
*****
Ionization Energy (IE): The minimum energy required to remove an electron from the ground state of a gaseous atom or ion.
First ionization energy (I1)
X(g) + I1  X+(g) + e–
Second ionization energy (I2)
X+(g) + I2  X2+(g) + e–
Endothermic process
The easier to remove valence electron, the lower IE
For an particular element: I1 < I2 < I3.....
Removing a core electron takes much more energy than removing a valence electron

16. Trends in First Ionization Energy (IE1)

17. *****
Continued
I1 generally decreases down a group because electrons in the outermost valence electron are farther away from the nucleus.
I1 generally increases across a period because electrons in the outermost valence electron experience a greater effective nuclear charge (Zeff).
The s- and p-block elements show a larger range of values for I1. (The d-block generally increases slowly across the period; the f-block elements show only small variations.)

18. Irregularities in the General Trend
Why actually the I1 of Be (2A) > I1 of B (3A)?
Ans: The 2p electron (of B) is at the higher energy level than 2s electron (of Be), thus it is relatively easy to be removed, therefore, I1 of Be > I1 of B.    Be
Be+ 1s 2s 2p 1s 2s 2p
To ionize Be, you must break up a full sublevel, which costs extra energy. B  1s 2s 2p  B+  1s 2s 2p
When you ionize B, you get a full sublevel, which costs less energy.

19. Continued Why actually the I1 of N (5A) > I1 of O (6A)?
Ans: Repulsion between paired electrons in a 2p orbital (of oxygen), lead the electron relatively easy to be removed therefore, IE1 of N > IE1 of O. N  1s 2s 2p  N+  1s 2s 2p 
To ionize N, you must break up a half-full sublevel, which costs extra energy.  O 1s 2s 2p   O+  1s 2s 2p
When you ionize O, you get a half-full sublevel, which costs less energy.

20. Successive Values of Ionization Energies
*****
Successive Values of Ionization Energies
Removal of each successive electron costs more energy. due to having more protons than electrons.
When all valence electrons have been removed, it takes a great deal more energy to remove the next electron.
Mg through Ar, Zeff increase, I2 decrease

21. Electron Configurations of Ions Anions examples:
*****
Electron Configurations of Ions
Anions examples:
Br(Z=35) [Ar]4s23d104p5 Br–: [Ar]4s23d104p6 or [Kr]
N(Z=7) [He]2s22p3 N3–: [He]2s22p6 or [Ne]
(B) Cations examples (2nd and 3rd period metals):
Na(Z=11) [Ne]3s1 Na+: [Ne] or 1s22s22p6
Al(Z=13) [Ne]3s23p1 Al3+: [Ne]

22. (As the (n-1)d orbitals begin to fill in the first transition series)
*****
(C) Cations examples (4th period and beyond metals)
ELevel
(As the (n-1)d orbitals begin to fill in the first transition series)
Ga(Z=31): [Ar]3d104s24p1
Ga3+: [Ar]3d10
Sn(Z=50): [Kr]4d105s25p2
Sn2+: [Kr]4d105s2
Sn4+: [Kr]4d10
Cr(Z=24):[Ar]4s23d4
Cr2+:[Ar]3d4
Cr3+:[Ar]3d3

23. *****
5. Electron Affinities
Electron affinity (EA): The energy change associated with the gaining of an electron by an atom in its gaseous state.
M(g) + e– → M–(g) ΔE = EA
The more negative (exothermic) the value, the larger the EA.
Trend of EA is much more irregular than IE

24. Trends in Electron Affinity
*****
Trends in Electron Affinity
Alkali metals decrease electron affinity down the column.
“Generally” increases across period (becomes more negative from left to right).
Group 5A generally has lower EA than expected because extra electron must pair.
Groups 2A and 8A generally have very low EA (many of these elements is positive) because added electron goes into higher energy level or sublevel.
Halogen have highest EA in any period.

25. 6. Metal, Nonmetals, and Metalloids

26. Metals Differ from Nonmetals
Metals tend to form cations.
Nonmetals tend to form anions.

27. Metals Properties of metals: Shiny luster Conduct heat and electricity
Malleable and ductile
Solids at room temperature (except mercury)
Low ionization energies/form cations easily

28. Metal Chemistry
Compounds formed between metals and nonmetals tend to be ionic.
Metal oxides tend to be basic.

29. Nonmetals
Nonmetals are found on the right hand side of the periodic table.
Properties of nonmetals include the following: Solid, liquid, or gas (depends on element)
Solids are dull, brittle, poor conductors
Large negative electronegativity/form anions readily
Dull 無光澤

30. Nonmetal Chemistry
Substances containing only nonmetals are molecular compounds.
Most nonmetal oxides are acidic.

31. Recap of a Comparison of the Properties of Metals and Nonmetals

32. Metalloids
Metalloids have some characteristics of metals and some of nonmetals.
Several metalloids are electrical semiconductors (computer chips).

33. Group Trends Elements in a group have similar properties.
Trends also exist within groups.
Groups Compared:
Group 1A: The Alkali Metals
Group 2A: The Alkaline Earth Metals
Group 6A: The Oxygen Group
Group 7A: The Halogens
Group 8A: The Noble Gases

34. Alkali Metals Alkali metals are soft, metallic solids.
They are found only in compounds in nature, not in their elemental forms.
Typical metallic properties (luster, conductivity) are seen in them.

35. Alkali Metal Properties
They have low densities and melting points.
They also have low ionization energies.

36. Alkali Metal Chemistry
Their reactions with water are famously exothermic.

37. Differences in Alkali Metal Chemistry
*****
Differences in Alkali Metal Chemistry
Lithium reacts with oxygen to make an oxide:
4 Li + O2  2 Li2O
Sodium reacts with oxygen to form a peroxide:
2 Na + O2  Na2O2
K, Rb, and Cs also form superoxides:
M + O2  MO2

38. Flame Tests
Qualitative tests for alkali metals include their characteristic colors in flames.

39. Alkaline Earth Metals—Compare to Alkali Metals
Alkaline earth metals have higher densities and melting points than alkali metals.
Their ionization energies are low, but not as low as those of alkali metals.

40. Alkaline Earth Metals
Beryllium does not react with water, and magnesium reacts only with steam, but the other alkaline earth metals react readily with water.
Reactivity tends to increase as you go down the group.

41. Group 6A—Increasing in Metallic Character down the Group
Oxygen, sulfur, and selenium are nonmetals.
Tellurium is a metalloid.
The radioactive polonium is a metal.

42. Group 7A—Halogens The halogens are typical nonmetals.
They have highly negative electron affinities, so they exist as anions in nature.
They react directly with metals to form metal halides.

43. Group 8A—Noble Gases
The noble gases have very large ionization energies.
Their electron affinities are positive (can’t form stable anions).
Therefore, they are relatively unreactive.
They are found as monatomic gases.

44. End of Chapter 07