1. Chapter 7 Periodic Properties of the Elements
Lecture Presentation
Chapter 7 Periodic Properties of the Elements
James F. Kirby
Quinnipiac University
Hamden, CT
[Modified by E.Schneider 11/2015

2. Development of the Periodic Table
Dmitri Mendeleev and Lothar Meyer independently:
Ordered elements by atomic weight
Observed that similar chemical & physical properties recur periodically
However….

3. Mendeleev and the Periodic Table
Mendeleev gets the glory, because he:
Put elements into columns based on similar properties
Recognized that “blank spots” in the table were for as-yet-undiscovered elements
Predicted – with good accuracy – the properties of the missing elements

4. Atomic Number
Mendeleev’s table based on atomic masses, the most fundamental property of elements known at the time.
Ernest Rutherford discovered the nuclear atom about 35 years later.
Henry Moseley experimentally developed the concept of atomic number.
The number of protons is the basis for the organization of the periodic property of elements.

5. Periodicity…
The repetitive pattern of a property for elements based on atomic number.
The following properties are discussed in this chapter:
Sizes of atoms and ions
Ionization energy
Electron affinity
Some group chemical property trends
First talk about effective nuclear charge – a fundamental property that leads to may of the trends

6. Effective Nuclear Charge
Many properties depend on attractions between valence electrons and the nucleus.
Electrons are both attracted to the nucleus AND repelled by other electrons.
The forces an electron experiences depend on both factors.

7. Effective Nuclear Charge
The effective nuclear charge: Zeff = Z − S
Z is the atomic number
S is a screening constant, roughly = number of inner electrons.
Effective nuclear charge is a periodic property:
It increases across a period.
It decreases down a group.
Analogy:

8. Effective Nuclear Charge Increases across a Period
As # of protons (atomic #) increases, Zeff increases.
As # of shells (n) increases, Zeff does a “step change decrease”, but then increases again across the period.
BUT…Zeff increases slightly as you move down the column because more diffuse e- cloud ↓ e- to e- repulsions

9. What Is the Size of an Atom?
Nonbonding atomic radius (van der Waals radius):
½ the shortest distance separating two nuclei during a collision of atoms (i.e. when atoms do NOT bond, so no “overlap” of orbitals)
Bonding atomic radius
½ the internuclear distance when atoms are bonded (i.e. orbitals DO “overlap”)

10. Give It Some Thought
Zeff ↑ as you move down a column of the periodic table
“Size” of an orbital (s, d, f, etc) ↑ as you move down a column of the periodic table
When considering the size of the atomic radii, do these trends work together or against each other?
Which effect is larger?

11. Use bonding atomic radii to calculate bond lengths
Sizes of Atoms
The bonding atomic radius tends to
Decrease from left to right across a period (Zeff ↑, so electrons pulled in closer to nucleus)
Increase from top to bottom of a group (n ↑, so atom gets bigger as # shells increase)
Use bonding atomic radii to calculate bond lengths

12. Predicting Relative Sizes of Atomic Radii
Sample Exercise 7.2
Predicting Relative Sizes of Atomic Radii
Referring to the periodic table, arrange (as much as possible) the atoms B, C, Al, and Si in order of increasing size.
Solution
C and B are in the same period, with C to the right of B. Therefore, we expect the radius of C to be smaller than that of B because radii usually decrease as we move across a period.
Likewise, the radius of Si is expected to be smaller than that of Al. Al is directly below B, and Si is directly below C. We expect, therefore, the radius of B to be smaller than that of Al and the radius of C to be smaller than that of Si.
Thus, so far we can say C < B, B < Al, C < Si, and Si < Al. We can therefore conclude that C has the smallest radius and Al has the largest radius, and so can write C < ? < ? < Al.
Discover that for the s- and p-block elements the increase in radius moving down a column tends to be the greater effect. There are exceptions, however.
C(0.76 Å) < B(0.84 Å) < Si(1.11 Å) < Al(1.21 Å).

13. Predicting Relative Sizes of Atomic Radii
Sample Exercise 7.2
Predicting Relative Sizes of Atomic Radii
Comment Note that the trends we have just discussed are for the s- and p-block elements. Figure 7.7 shows that the transition elements do not show a regular decrease moving across a period.
Practice Exercise 1
By referring to the periodic table, but not to Figure 7.7, place the following atoms in order of increasing bonding atomic radius: N, O, P, Ge.
(a) N < O < P < Ge (b) P < N < O < Ge (c) O < N < Ge < P (d) O < N < P < Ge
(e) N < P < Ge < O

14. Predicting Relative Sizes of Atomic Radii
Sample Exercise 7.2
Predicting Relative Sizes of Atomic Radii
Comment Note that the trends we have just discussed are for the s- and p-block elements. Figure 7.7 shows that the transition elements do not show a regular decrease moving across a period.
Practice Exercise 1
By referring to the periodic table, but not to Figure 7.7, place the following atoms in order of increasing bonding atomic radius: N, O, P, Ge.
(a) N < O < P < Ge (b) P < N < O < Ge (c) O < N < Ge < P (d) O < N < P < Ge
(e) N < P < Ge < O
Practice Exercise 2
Arrange Be, C, K, and Ca in order of increasing atomic radius.
C < Be < Ca < K

15. Sizes of Ions Ionic size depends on the nuclear charge.
the number of electrons.
the orbitals in which electrons reside.

16. Sizes of Ions 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.

17. Predicting Relative Sizes of Atomic and Ionic Radii
Sample Exercise 7.3
Predicting Relative Sizes of Atomic and Ionic Radii
Arrange Mg2+, Ca2+, and Ca in order of decreasing radius.
Solution
Cations are smaller than their parent atoms, and so Ca2+ < Ca. Because Ca is below Mg in group 2A, Ca2+ is larger than Mg2+. Consequently, Ca > Ca2+ > Mg2+.
Practice Exercise 1
Arrange the following atoms and ions in order of increasing ionic radius: F, S2–, Cl, and Se2–.
(a) F < S2– < Cl < Se2– (b) F < Cl < S2– < Se2– (c) F < S2– < Se2– < Cl
(d) Cl < F < Se2– < S2– (e) S2– < F < Se2– < Cl

18. Predicting Relative Sizes of Atomic and Ionic Radii
Sample Exercise 7.3
Predicting Relative Sizes of Atomic and Ionic Radii
Arrange Mg2+, Ca2+, and Ca in order of decreasing radius.
Solution
Cations are smaller than their parent atoms, and so Ca2+ < Ca. Because Ca is below Mg in group 2A, Ca2+ is larger than Mg2+. Consequently, Ca > Ca2+ > Mg2+.
Practice Exercise 1
Arrange the following atoms and ions in order of increasing ionic radius: F, S2–, Cl, and Se2–.
(a) F < S2– < Cl < Se2– (b) F < Cl < S2– < Se2– (c) F < S2– < Se2– < Cl
(d) Cl < F < Se2– < S2– (e) S2– < F < Se2– < Cl
Practice Exercise 2
Which of the following atoms and ions is largest: S2–, S, O2–?
S2–

19. Size of Ions— Isoelectronic Series
Isoelectronic series, ions have the same number of electrons.
Ionic size ↓with an ↑ nuclear charge
An Isoelectronic Series (10 electrons)
Note increasing nuclear charge with decreasing ionic radius as atomic number increases
Charge
O2– F–
Na+
Mg2+
Al3+
Ionic Radius
1.26 Å
1.19 Å
1.16 Å
0.86 Å
0.68 Å

20. Ionic Radii in an Isoelectronic Series
Sample Exercise 7.4
Ionic Radii in an Isoelectronic Series
Arrange the ions K+, Cl–, Ca2+, and S2– in order of decreasing size.
Solution
This is an isoelectronic series, with all ions having 18 electrons. In such a series, size decreases as nuclear charge (atomic number) increases. The atomic numbers of the ions are S 16, Cl 17, K 19, Ca 20. Thus, the ions decrease in size in the order S2– > Cl– > K+ > Ca2+.
Practice Exercise
In the isoelectronic series Ca2+, Cs+, Y3+, which ion is largest?
Cs+

21. Ionization Energy (I)
Ionization energy - minimum energy required to remove an electron from the ground state of a gaseous atom or ion.
The first ionization energy is that energy required to remove the first electron.
The second ionization energy is that energy required to remove the second electron, etc.
Note: the higher the ionization energy, the more difficult it is to remove an electron!

22. Ionization Energy
It requires more energy to remove each successive electron.
When all valence electrons have been removed, it takes a great deal more energy to remove the next electron.

23. Periodic Trends in First Ionization Energy (I1)
I1 generally increases across a period.
I1 generally decreases down a group.
s- and p-block elements show a larger range of values for I1.
d-block generally increases slowly across the period
f-block elements show only small variations.

24. Factors that Influence Ionization Energy
Smaller atoms have higher I1 values.
I1 values depend on - effective nuclear charge and average distance of the electron from the nucleus.

25. Sample Exercise 7.5 Trends in Ionization Energy
Three elements are indicated in the periodic table below. Which one has the largest second ionization energy?
Solution
Solve The red box represents Na, which has one valence electron. The second ionization energy of this element is associated, therefore, with the removal of a core electron. The other elements indicated, S (green) and Ca (blue), have two or more valence electrons. Thus, Na should have the largest second ionization energy.
Check A chemistry handbook gives these I2 values: Ca, 1145 kJ ⁄ mol; S, 2252 kJ ⁄ mol; Na, 4562 kJ ⁄ mol.

26. Sample Exercise 7.5 Trends in Ionization Energy
Continued
Practice Exercise 1
The third ionization energy of bromine is the energy required for which of the following processes?
(a) Br(g) → Br+(g) + e – (b) Br+(g) → Br2+(g) + e – (c) Br(g) → Br2+(g) + 2e –
(d) Br(g)→ Br3+(g) + 3e– (e) Br2+(g) → Br3+(g) + e –
Practice Exercise 2
Which has the greater third ionization energy, Ca or S?

27. Irregularities in the General Trend for I1
The trend is not followed when the added valence electron in the next element
enters a new sublevel (higher energy sublevel);
is the first electron to pair in one orbital of the sublevel (electron repulsions lower energy).

28. Go Figure
The I1 value for astatine, At, is missing in this figure. To the nearest 100 kJ/mol what estimate would you make for the first ionization energy of At?
Anomolies: ex Be B

29. Go Figure
Why is it easier to remove a 2p electron from an oxygen atom than from a nitrogen atom?

30. K < Na < P < Ar < Ne
Sample Exercise 7.6
Periodic Trends in Ionization Energy
Referring to the periodic table, arrange the atoms Ne, Na, P, Ar, K in order of increasing first ionization energy.
Solution
Analyze and Plan We are given the chemical symbols for five elements. To rank them according to increasing first ionization energy, we need to locate each element in the periodic table. We can then use their relative positions and the trends in first ionization energies to predict their order.
Solve Ionization energy increases as we move left to right across a period and decreases as we move down a group. Because Na, P, and Ar are in the same period, we expect I1 to vary in the order Na < P < Ar.
Sample
Because Ne is above Ar in group 8A, we expect Ar < Ne. Similarly, K is directly below Na in group 1A, and so we expect K < Na.
From these observations, we conclude that the ionization energies follow the order
K < Na < P < Ar < Ne

31. Periodic Trends in Ionization Energy
Sample Exercise 7.6
Periodic Trends in Ionization Energy
Check The values shown in Figure 7.10 confirm this prediction.

32. Sample Exercise 7.6 Periodic Trends in Ionization Energy
Continued
Practice Exercise 1
Consider the following statements about first ionization energies:
(i) Because the effective nuclear charge for Mg is greater than that for Be, the first ionization energy of Mg is greater than that of Be.
(ii) The first ionization energy of O is less than that of N because in O we must pair electrons in the 2p orbitals.
(iii) The first ionization energy of Ar is less than that of Ne because a 3p electron in Ar is farther from the nucleus than a 2p electron in Ne.
Which of the statements (i), (ii), and (iii) is or are true?
(a) Only one of the statements is true.
(b) Statements (i) and (ii) are true.
(c) Statements (i) and (iii) are true.
(d) Statements (ii) and (iii) are true.
(e) All three statements are true.
Practice Exercise 2
Which has the lowest first ionization energy, B, Al, C, or Si? Which has the highest?

33. Electron Configurations of Ions
Cations (+):
Electrons are lost from the highest energy level (n value)
Li Li+ is 1s2 2s1  1s2 (losing a 2s electron).
Fe Fe2+ is [Ar]4s23d6  [Ar] 3d6 (losing two 4s electrons).
If >1 occupied subshell for highest n level, electron is lost from the orbital with the highest l level.
Anions (-):
Electrons added into the empty or partially filled orbital having the lowest value of n.
F F– is [He]2s22p5  [He]2s22p6 (gaining one electron in 2p).
Different than when adding e- in electron config diagrams

34. Give It Some Thought
Do Cr3+ and V2+ have the same or different electron configurations?

35. Sample Exercise 7.7 Electron Configurations of Ions
Write the electron configurations for (a) Ca2+, (b) Co3+, and (c) S2–.
Solution
(a) Calcium (atomic number 20) has the electron configuration [Ar]4s2. To form a 2+ ion, the two outer 4s electrons must be removed, giving an ion that is isoelectronic with Ar:
Ca2+: [Ar]
(b) Cobalt (atomic number 27) has the electron configuration [Ar]4s23d7. To form a 3+ ion, three electrons must be removed. As discussed in the text, the 4s electrons are removed before the 3d electrons. Consequently, we remove the two 4s electrons and one of the 3d electrons, and the electron configuration for Co3+ is
Co3+: [Ar]3d6
(c) Sulfur (atomic number 16) has the electron configuration [Ne]3s23p4. To form a 2– ion, two electrons must be added. There is room for two additional electrons in the 3p orbitals. Thus, the S2– electron configuration is
S2–: [Ne]3s23p6 = [Ar]
Comment Remember that many of the common ions of the s- and p-block elements, such as Ca2+ and S2–, have the same number of electrons as the closest noble gas (Section 2.7)

36. Sample Exercise 7.7 Electron Configurations of Ions
Continued
Practice Exercise 1
The ground electron configuration of a Tc atom is [Kr]5s24d5.
What is the electron configuration of a Tc3+ ion?
(a) [Kr]4d4 (b) [Kr]5s24d2 (c) [Kr]5s14d3 (d) [Kr]5s24d8 (e) [Kr]4d10
Practice Exercise 2
Write the electron configurations for (a) Ga3+, (b) Cr3+, and (c) Br–.

37. Electron Affinity
Electron affinity: energy change that occurs when an electron is added to a gaseous atom:
Measures the “affinity” of the atom for gaining an electron.
Typically exothermic, so, for most elements, it is negative!
Comparison: 1st Ionization Energy is energy change that occurs when an electron is removed from an atom.
Energy needs to be absorbed by the atom to remove an electron, so endothermic.

38. General Trend in Electron Affinity
Little change within a group:
Atomic diameter as move down, so e- attraction to nucleus 
BUT higher energy orbitals are increasingly spread out, so e- to e- repulsions reduced.

39. General Trend in Electron Affinity
Key groups:
Halogens (Group 7A): highest neg values, greatest affinity
Noble Gases (Group 8A) & Alkaline Earth Metals (Group 2A): positive because added e- would go into a higher energy level orbital.
Group 5A: p sublevel is half-full, so added e- would go into orbital that already has 1 e- , so e- to e- replusion 
Group 8A: p sublevel is full!

40. General Trend in Electron Affinity
Across a period, it generally increases. Three notable exceptions include the following:
Group 2A & Group 8A: sublevels are full, so added e- goes into a higher energy sublevel!
Group 5A: p sublevel is half-full, so increase electron to electron repulsion.

42. Metal, Nonmetals, and Metalloids

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

44. Metals Most of the elements in nature are 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

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

46. 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

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

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

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

50. 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

51. 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.

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

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

54. 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

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

56. 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.

57. 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.

58. 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.

59. 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.

60. 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.