1. Periodic Relationships Among the Elements
Chapter 8
8.1 Development of the Periodic Table
8.2 Periodic Classification of the Elements
8.3 Periodic Variation in Physical Properties
8.4 Ionization Energy
8.5 Electron Affinity
8.6 Variation in Chemical Properties of the Representative Elements
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2. 8.1 Development of the Periodic Table
J. W. Dobereiner grouped elements into triads:
Three elements with similar properties;
Properties followed a pattern;
The same element was in the middle of all trends
BUT not all elements belongs to a triad.
By mid 19th century, the atomic masses of many of the elements had been determined.
John Newlands:
hypothesized that the chemistry of the elements might be related to their masses;
arranged the known elements in order of increasing atomic masses;
discovered that every 7th element had similar properties (noble gases were still unknown during this time);
Proposed the arrangement of elements into ‘octaves’.
BUT Newland’s ‘Law of Octaves’ did not work for some elements.

3. Julius Lothar Meyer & Dimitri Mendeleev focused on the relationship between atomic mass & various physical and chemical properties of elements.
Meyer aligned the elements according to ‘atomic volumes’.
Mendeleev:
Credited with the origin of the modern periodic table;
Assumed that some elements have not been discovered yet, and deliberately left blanks in his table.
H.G.J. Moseley:
Discovered that the foundation to the order of the elements was the atomic number, not atomic mass.

4. The Modern Periodic Table:
Elements are arranged in order of increasing atomic number;
There are 7 horizontal rows called Period;
There are 18 vertical columns called Group (or families);
Classification of elements: metals, non-metals & metalloids;
Electron configurations help explain the recurrence of physical & chemical properties;
Very important: understanding of general properties and trends within a group or a period helps in predicting the properties of any element.
Non-Metal: Elements in the upper right corner of the periodic Table. Their chemical and physical properties are different from metals.
Metal: Elements that are usually solids at room temperature. Most elements are metals.
Metalloid: Elements that lie on a diagonal line between the Metals and non-metals. Their chemical and physical properties are intermediate between the two.

6. Classification of the Elements

7. Periodic Patterns:
The chemical behavior of elements is determined by its electron configuration.
Energy levels are quantized so roughly correspond to layers of electrons around the nucleus.
A shell is all the electrons with the same value of n. n is a row in the periodic table.
Each period begins with a new outer electron shell.
Each period ends with a completely filled outer shell that has the maximum number of electrons for that shell.
The number identifying the families identifies the number of electrons in the outer shell, except helium.
The outer shell electrons are responsible for chemical reactions.

8. 8.2 Periodic Classification of the Elements
According to the type of subshells being filled, the elements can be divided into:
Representative Elements (Main Group Elements)
Group 1A – 7A (Group 1, 2, 13, 14, 15, 16, 17);
All elements have incompletely filled s or p subshells of the highest principal quantum number;
Noble Gases – Group 8A or Group 18; all have completely filled p subshell.
Transition Elements (d block transition elements)
Groups 1B, 3B – 8B (Group 3 – 11), with incompletely filled d subshells.
Special metals – Zn, Cd, Hg.
Lanthanides – f block transition elements, have incompletely filled f subshells.
Actinides – also called f block transition elements.

9. 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

10. Valence and Core Electrons
Valence Electrons - Those electrons outside of a closed electron shell. These electrons take part in chemical reactions.
Core Electrons - The electrons in the closed shells. They cannot take part in chemical reactions.
Sodium 11 electrons:
Valence electrons [Ne] 3s one
Core electrons s 2 2s 2 2p ten
Chlorine 17 electrons:
Valence electrons [Ne] 3s 2 3p seven
Core electrons 1s2 2s 2 2p ten

11. 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]

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

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
Isoelectronic is the condition where ions and/or atoms possess the same number of electrons and hence the same ground state electronic configurations.
What neutral atom is isoelectronic with H− ?
8.2

14. Electron Configurations of Cations of Transition Metals
When a cation is formed from an atom of a transition metal, electrons are always removed first from the ns orbital and then from the (n – 1)d orbitals.
Fe: [Ar]4s23d6
Fe2+: [Ar]4s03d6 or [Ar]3d6
Fe3+: [Ar]4s03d5 or [Ar]3d5
Mn: [Ar]4s23d5
Mn2+: [Ar]4s03d5 or [Ar]3d5
8.2

15. 8.3 Periodic Variation in Physical Properties
Atomic size/radius – how big the atoms are;
Ionic size/radius – how big ions are.
Ionization energy – how much energy to remove an electron;
Electron afinity – the ability to accept one or more electrons.
Periodic trends – How does the four properties above vary as you go across a period?
Group trends – How does the four properties above vary as you go down a group
Why? Explain why they vary…
8.3

16. 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
8.3

17. Inner shell electrons shielded the nuclear charge effectively from the outer shell electrons.
Example: Li 1s2 2s1
Electrons in 1s shield the nuclear charge and the 2s electron feels or sees the 3+ nucleus as having 1+ charge. The net charge or Zeff is 1+.
Electrons in the same shell do not shield each other effectively because, in average, they are at the same distance.
Example: Be 1s2 2s2
The two electrons in 2s see the nucleus as having a 2+ charge.

18. Effective Nuclear Charge (Zeff)
increasing Zeff
increasing Zeff

19. Atomic radius – variation in size.
Atomic size is usually represented by atomic radius. Since an atom are usually bonded to another atom, atomic radius is taken as half the distance between two closest identical atoms.
One-half the distance between the nuclei centers of two adjacent atoms.
(covalent radius)
(metallic radius)

20. Fig. 8.12 Periodic variation in atomic size:
The atomic size become larger going from top to bottom in a group and become smaller from left to right across a period.
Within a group, the Zeff felt by the outer electrons is the same. However, the core size is increasing due to increasing value of n, thus, increasing the atomic size.
Within a period, electrons are added into the same shell. As Zeff is increasing from left to right, the attraction force is stronger and the electrons are drawn inward.
The size variation is less pronounce for the transition element.
Fig. 8.12

22. Trends in Atomic Radii

23. Question #1: Ranking Elements by Size
Problem: Rank the following elements in each group according to decreasing size ( largest first!):
Na, K, Rb
Sr, In, Rb
Cl, Ar, K
Sr, Ca, Rb

24. Question #2: Ranking Elements by Size
Using only a periodic Table, arrange these atoms in order of decreasing atomic size:
Br , Cl , Ge , K , S

25. Comparison of Atomic Radii with Ionic Radii

26. Ionic radius – variation in size.
Anion is always larger than their neutral atom. When electrons are added, mutual repulsions increase which cause the electrons to occupy larger volume.
Cation is always smaller than their neutral atom. When electrons are remove, mutual repulsion decrease, electrons are pulled closer together around the nucleus.
Periodic variation in ionic size is similar to periodic variation in atomic size.

27. For isoelectronic cations: M3+ < M2+ < M+
Example: Al3+ < Mg2+ < Na+
Explanation: all cations have the same number of electrons (10 electrons) but Al has 13 protons, Mg has 12 protons whereas Na has only 11 protons. Thus, the electron cloud in Al will be pulled closer to the nucleus as compared to Mg or Na.
For isoelectronic anions: X3– > X2– > X–
Example: N3– > O2– > F–
Explanation: all anions have the same number of electrons (10 electrons) but the repulsion is much stronger in N3– than in O2– or F–.
8.3

28. The Radii (in pm) of Ions of Familiar Elements

29. Question #3: Ranking Ions According to Size
Rank each set of Ions in order of increasing size:
a) K+, Rb+, Na b) Na+, O2-, F c) Fe+2, Fe+3

30. 8.4 Ionization Energy
Ionization energy (IE) is the minimum energy (kJ/mol) required to remove an electron from a gaseous atom in its ground state.
IE1 + X (g) X+(g) + e-
IE2 + X+(g) X2+(g) + e-
IE3 + X2+(g) X3+(g) + e-
IE1 first ionization energy
IE2 second ionization energy
IE3 third ionization energy
IE1 < IE2 < IE3
Periodic trend of IE:
IE generally decrease from top to bottom within a Group and increase from left to right across a Period.
Within a Group, n increases going down a Group, shells become larger and outer electrons are farther away from the nucleus and less tightly bound.
Going across a Period, atomic size become smaller as Zeff increase. The electrons are held more tightly and more difficult to be removed.
8.4

31. Example: Li  Li+ + e IE1 = 520 kJ/mol
IE1 < IE2 < IE3
The 1st electron removed is from 2s orbital and the product gains stable noble gas configuration.
More energy is required to remove electron from an ion that has reached noble gas configuration. Therefore, a sudden increase (a large jump) in value of IE can be seen.
Ionization is always endothermic, because atom (or ion) need to absorb energy to remove the electron(s).
8.4

33. Variation of the First Ionization Energy with Atomic Number
Filled n=1 shell
Filled n=2 shell
Filled n=3 shell
Filled n=4 shell
Filled n=5 shell

34. General Trends in First Ionization Energies
Increasing First Ionization Energy
Increasing First Ionization Energy

35. Successive Ionization Energies
# Valence
Z Element Electrons IE1 IE2 IE IE IE IE6 IE7 Li Be B C N O

36. Question #4:
Using the Periodic table only, rank the following elements in each of the following sets in order of increasing 1st IE.
a) Ar, Ne, Rn c) Be, Na, Mg
b) At, Bi, Po d) Cl, K, Ar

37. Question #5:
Given the following series of ionization energies (in kJ/mol) for an element in period 3, name the element and write its electron configuration:
IE IE IE IE4
, , ,600

38. 8.5 Electron Affinity
Electron affinity(EA) is the negative value of the energy change that occurs when an electron is accepted by an atom in the gaseous state to form an anion.
EA is a measure of the attraction the atom has for an additional electron.
X (g) + e X-(g)
F (g) + e X-(g)
DH = -328 kJ/mol
EA = +328 kJ/mol
O (g) + e O-(g)
DH = -141 kJ/mol
EA = +141 kJ/mol
The large negative value for Fluorine indicates that Fluorine atoms readily accept an electron.
8.5

39. Trend in Electron Affinity
EA decreases when going down a Group (less exothermic) and increases when going across a period from left to right (more exothermic).

41. Variation of Electron Affinity With Atomic Number (H – Ba)

42. Summary of Trend Periodic Table and Periodic Trends
1. Electron Configuration
3. Ionization Energy: Largest toward NE of PT
4. Electron Affinity: Most favorable NE of PT
2. Atomic Radius: Largest toward SW corner of PT

43. Diagonal relationships:
8.6 Variation in Chemical Properties of the Representative Elements
Elements in the same group resembles one another in chemical behaviour because they have similar outer electron configuration.
Diagonal relationships:
Similarities between pairs of elements in different groups and periods in the Periodic Table.
8.3

44. Closeness of the charge densities of their cations;
Li, Be and B (elements form 2nd period) exhibits many similarities to those elements located diagonally below them in the Periodic Table.
The reason:
Closeness of the charge densities of their cations;
Similar measurement of atomic and ionic radii;
Li = 152 pm Li+ = 76 pm
Mg = 160 pm Mg2+ = 72 pm
Both cations will react similarly with the same anion & form the same type of compound:
Li + N2  Li3N, lithium nitride
Mg + N2  Mg3N2 , magnesium nitride
8.6

45. 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

46. Group 1A Elements (ns1, n  2)

47. 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

48. Group 2A Elements (ns2, n  2)

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

50. Group 3A Elements (ns2np1, n  2)

51. Group 4A Elements (ns2np2, n  2)
Sn(s) + 2H+(aq) Sn2+(aq) + H2 (g)
Pb(s) + 2H+(aq) Pb2+(aq) + H2 (g)

52. Group 4A Elements (ns2np2, n  2)

53. Group 5A Elements (ns2np3, n  2)
N2O5(s) + H2O(l) HNO3(aq)
P4O10(s) + 6H2O(l) H3PO4(aq)

54. Group 5A Elements (ns2np3, n  2)

55. Group 6A Elements (ns2np4, n  2)
SO3(g) + H2O(l) H2SO4(aq)

56. Group 6A Elements (ns2np4, n  2)

57. Group 7A Elements (ns2np5, n  2)
X + 1e X-1
X2(g) + H2(g) HX(g)
Increasing reactivity

58. Group 7A Elements (ns2np5, n  2)

59. Group 8A Elements (ns2np6, n  2)
Completely filled ns and np subshells.
Highest ionization energy of all elements.
No tendency to accept extra electrons.

60. Compounds of the Noble Gases
A number of xenon compounds XeF4, XeO3, XeO4, XeOF4 exist.
A few krypton compounds (KrF2, for example) have been prepared.

61. Comparison of Group 1A and 1B
The metals in these two groups have similar outer electron configurations, with one electron in the outermost s orbital.
Chemical properties are quite different due to difference in the ionization energy.
Lower IE1, more reactive

62. Properties of Oxides Across a Period
Metallic characters of the elements decreases from left to right across a Period and increases from top to bottom within a Group.
Most oxides of representative metals are ionic compounds; in water they act as bases, producing OH ions and reacting with acids.
Example: Barium forms the basic oxide BaO:
Reacts with water:
BaO (s) + H2O (ℓ)  Ba2+ (aq) + 2 OH– (aq)
Reacts with acid:
BaO (s) + H+ (aq)  Ba2+ (aq) + H2O (ℓ)
Other examples of basic oxides:
Li2O , Na2O , K2O , MgO , CaO , SrO
8.6

63. Example: Sulfur forms the acidic oxide SO2: Reacts with water:
Non-metal oxides are covalent compounds; in water they act as acids, producing H+ ions and reacting with bases.
Example: Sulfur forms the acidic oxide SO2:
Reacts with water:
SO2 (g) + H2O (ℓ) ⇌ H2SO3 (aq) ⇌ H+ (aq) + HSO3– (aq)
Reacts with base:
SO2 (g) + 2 OH– (aq)  SO32– (aq) + H2O (ℓ)
Other examples of acidic oxides:
Cl2O7 , SO3 , N2O5 , N2O3 , P4O10 , CO2 , B2O3
8.6

64. Example: Al2O3 is an amphoteric oxide: Reacts with acid:
Some metals and many metalloids form oxides that are amphoteric: they can act as acids and as bases.
Example: Al2O3 is an amphoteric oxide:
Reacts with acid:
Al2O3 (s) + 6 H+ (aq)  2 Al3+ (aq) + 3 H2O (ℓ)
Reacts with base:
Al2O3 (s) + 2 OH– (aq) + 3 H2O ()  2 Al(OH)4– (aq)
Other examples of amphoteric oxides:
BeO , Ga2O3 , SnO2 , SnO , PbO2
8.6

65. Properties of Oxides Across a Period
basic
acidic
8.6

66. The Trends in Acid-Base Behavior of Elemental Oxides:
As the elements become more metallic down a Group, their oxides become more basic.
As the elements become less metallic across a period, their oxides become more acidic.
Fig. 8.22