General Chemistry 1

1. The Periodic Table
Chapter 8
Periodic Relationships Among the Elements
Dr. Sa’ib Khouri
AUM- JORDAN
Chemistry
By Raymond Chang

2. 2
The discovery of the elements chronologically

3. 3
ns1
ns2
ns2np1
ns2np2
ns2np3
ns2np4
ns2np5
ns2np6
d1
d5
d10
4f
5f
Ground State Electron Configurations of the Elements

4. 4
Classification of the Elements
Note: Group 2B elements are often classified as transition metals even
though they do not exhibit the characteristics of the transition metals.

5. Representative elements (main group elements): Groups 1A
through 7A, all of which have incompletely filled s or p subshells
of the highest principal quantum number.
Noble gases: all have a completely filled p subshell except
helium.
Transition metals: Groups 1B and 3B through 8B; the elements
which have incompletely filled d subshells, or readily produce
cations with incompletely filled d subshells (d-block transition
elements).
Lanthanides and actinides: incompletely filled f subshells.
The group 2B elements are Zn, Cd, Hg, which are neither
representative elements nor transition metals.

6. 6
The chemical reactivity of the elements is largely
determined by their valence electrons, which are the
outermost electrons.
The International Union of Pure and Applied Chemistry
(IUPAC) has recommended numbering the columns
sequentially with Arabic numerals 1 through 18.
Notes
For the representative elements, the valence electrons
are those in the highest occupied n shell.
All nonvalence electrons in an atom are referred to as
core electrons.

7. 7

8. 8
Electron Configurations of Cations and Anions
Na: [Ne]3s1
Na+
: [Ne]
Ca: [Ar]4s2
Ca2+
: [Ar]
Al : [Ne]3s2
3p1
Al3+
: [Ne]
Atoms lose electrons so that
cation has a stable noble-gas
outer electron configuration
(ns
2
np
6
)
H :1s1
H- : 1s2
or [He]
F : 1s2
2s2
2p5
F- : 1s2
2s2
2p6
or [Ne]
O : 1s2
2s2
2p4
O2- : 1s2
2s2
2p6
or [Ne]
N : 1s2
2s2
2p3
N3- : 1s2
2s2
2p6
or [Ne]
Atoms gain electrons
so that anion has a
stable noble-gas outer
electron configuration
ns
2
np
6
)
Pauli exclusion principle and Hund’s rule are used in writing the
ground-state electron configurations of cations and anions.
Ions Derived from Representative Elements

9. 9
Notice that F-
, Na+
, and Ne (and Al3+
, O2-
, and N3-
) have the
same number of electrons, and hence the same ground-state
electron configuration. They are said to be isoelectronic.
H
-
(1s
2
) and He are also isoelectronic)
Cations Derived from 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, which is more stable than the
ns orbital.
Fe: [Ar]4s23d6
Fe2+: [Ar]4s03d6 or [Ar]3d6
Fe3+: [Ar]4s03d5 or [Ar]3d5
Mn: [Ar]4s23d5
Mn2+: [Ar]4s03d5 or [Ar]3d5
Mn3+: [Ar]4s03d4 or [Ar]3d4

10. 10
+1
+2
+3
-1
-2
-3
Cations and Anions of Representative Elements

11. 11

12. 12
Effective nuclear charge (Zeff)
Na
Mg
Al
Si
11
12
13
14
10
10
10
10
1
2
3
4
186
160
143
132
ZeffCoreZ Radius (pm)
Zeff = Z - s 0 < s < Z
Zeff  Z – number of inner or core electrons
Periodic Variation in Physical Properties
s: the shielding constant (or screening constant) and Z: nuclear charge
is the “positive nuclear charge” felt by an electron when both
the actual nuclear charge (Z) and the repulsive effects
(shielding) of the other electrons are taken into account.

13. 13
In a period the added electron is a valence electron and
valence electrons do not shield each other well, the net effect
of moving across the period is a greater effective nuclear
charge felt by the valence electrons.
Example on the second period:
The effective nuclear charge also increases as we go down a
particular periodic group, because the valence electrons are
now added to increasingly large shells as n increases, the
electrostatic attraction between the nucleus and the valence
electrons actually decreases. Example on the first group:

14. 14
Effective Nuclear Charge (Zeff)
increasing Zeff
increasingZeff

15. 15
Atomic Radius
metal diatomic molecule
Density, melting point, and boiling point are related to the sizes
of atoms. The size of an atom defined in terms of its atomic
radius.
Atomic radius: one-half the distance between the two nuclei
in two adjacent metal atoms or in a diatomic molecule.

16. 16

17. 17
because the effective nuclear charge increases
from left to right, the added valence electron at
each step is more strongly attracted by the
nucleus than the one before, therefore we find
the atomic radius decreases from left to right.
orbital size increases with the increasing
principal quantum number n. the size of the
atomic radius increases even though the
effective nuclear charge also increases.
In a period
In a group

18. 18
plots the atomic radii of the elements
against their atomic numbers

19. 19

20. 20
Ionic Radius
is the radius of a cation or an anion
Ionic radius affects the physical and chemical
properties of an ionic compound.
If the atom forms an anion, its size (or radius)
increases because the repulsion resulting from the
additional electron(s) enlarges the domain of the
electron cloud.
If the atom forms an cation, removing one or more
electrons from an atom reduces electron-electron
repulsion, so the electron cloud shrinks.
Note: in both cases the nuclear charge remains the
same.

21. 21
Comparison of Atomic Radii with Ionic Radii

22. 22
Cation is always smaller than atom from which it is
formed.
Anion is always larger than atom from which it is
formed.

23. 23
The Radii (in pm) of Ions of Familiar Elements

24. 24
is the minimum energy (kJ/mol) required to remove an
electron from a gaseous atom (no forces between atoms) 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
For many-electron atom
When removing the first electron, the repulsion among the
remaining electrons decreases, and more energy is needed to
remove another electron.
Ionization Energy (IE)
The atom of the higher ionization energy, the more difficult it
is to remove its electron, i.e it is more stable.

25. 25

26. 26
Ionization is always an endothermic process. Energy
absorbed by atoms (or ions) in the ionization process
has a positive value.
The first ionization energies of elements in a period
increase with increasing atomic number due to the
increase in effective nuclear charge from left to right.
The high ionization energies of the noble gases,
stemming from their large effective nuclear charge,
comprise of the reasons for this stability.

27. 27
Filled n=1 shell
Filled n=2 shell
Filled n=3 shell
Filled n=4 shell
Filled n=5 shell
Variation of the First Ionization Energy with Atomic Number

28. 28
General Trends in First Ionization Energies
Increasing First Ionization Energy
IncreasingFirstIonizationEnergy

29. 29
Electron affinity (EA)
X (g) + e- X
-
(g)
F (g) + e- F
-
(g)
O (g) + e- O
-
(g)
DH = -328 kJ/mol EA = +328 kJ/mol
DH = -141 kJ/mol EA = +141 kJ/mol
More positive EA , greater affinity to accept an electron
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.

30. 30

31. 31
Variation of Electron Affinity With Atomic Number (H – Ba)

32. 32
There is a general correlation between electron affinity
and effective nuclear charge, which also increases
from left to right in a given period.
The variation in electron affinities from top to bottom
within a group is much less regular.

33. Variation in Chemical Properties
of the Representative Elements
Ionization energy and electron affinity help
chemists understand the types of reactions
that elements undergo and the nature of the
elements’ compounds.

35. 35
Group 1A Elements (ns1, n  2)
M M+1 + 1e-
2M(s) + 2H2O(l) 2MOH(aq) + H2(g)
4M(s) + O2(g) 2M2O(s)
Increasingreactivity

38. 38
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
Increasingreactivity

40. 40
Group 3A Elements (ns2np1, n  2)
4Al(s) + 3O2(g) 2Al2O3(s)
2Al(s) + 6H+
(aq) 2Al3+
(aq) + 3H2(g)

41. The other elements:
Form both unipositive (+1) and tripositive ions (+3).
The metallic elements in Group 3A also form many
molecular compounds.
Moving down the group: the unipositive ion (+1)
becomes more stable than the tripositive ion (+3).
Example: AlH3, which resembles BeH2 in its
properties.

43. 43
Group 4A Elements (ns2np2, n  2)
Carbon: nonmetal
Silicon and Germanium: metalloid
Tin and Lead: metallic (do not react with water,
but reacts with acids)

44. Sn(s) + 2H+
(aq) Sn2+
(aq) + H2 (g)
Pb(s) + 2H+
(aq) Pb2+
(aq) + H2 (g)
The elements of this Group form compounds in both the +2
and +4 oxidation states.
For carbon and silicon, the +4 oxidation state is the more
stable one.
For example:
CO2 is more stable than CO, and SiO2 is a stable compound,
but SiO does not exist under normal conditions.

46. 46
Group 5A Elements (ns2np3, n  2)
Nitrogen and Phosphorus: nonmetals
Arsenic and Antimony: metalloids
Bismuth: metal

47. N2O5(s) + H2O(l) 2HNO3(aq)
P4O10(s) + 6H2O(l) 4H3PO4(aq)
Nitrogen forms a number of oxides (NO, N2O, NO2 , N2O4 , and
N2O5), of which only N2O5 is a solid; the others are gases.
Most metallic nitrides (like Li3N and Mg3N2) are ionic
compounds.
Phosphorus exists as P4 molecules, it forms two solid oxides
with the formulas P4O6 and P4O10
Arsenic, antimony, and bismuth have extensive three-
dimensional structures. Bismuth is a far less reactive metal
than those in the preceding groups.

49. 49
Group 6A Elements (ns2np4, n  2)
SO3(g) + H2O(l) H2SO4(aq)
Oxygen, sulfur, and selenium: are nonmetals
Tellurium and polonium: are metalloids
The important compounds of sulfur are SO2, SO3,
and H2S

50. All the halogens are nonmetals with the general formula X2,
where X denotes a halogen element.

51. 51
Group 7A Elements (ns2np5, n  2)
Increasingreactivity
X + 1e- X-1
X2(g) + H2(g) 2HX(g)

52. Noble gases

53. 53
Group 8A Elements (ns2np6, n  2)
Completely filled ns and np subshells. Highest
ionization energy of all elements. No tendency
to accept extra electrons.