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The Periodic Table
Chapter 8
Periodic Relationships Among the Elements
Dr. Sa’ib Khouri
AUM- JORDAN
Chemistry
By Raymond Chang
2
The discovery of the elements chronologically
3
ns1
ns2
ns2np1
ns2np2
ns2np3
ns2np4
ns2np5
ns2np6
d1
d5
d10
4f
5f
Ground State Electron Configurations of the Elements
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.
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
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
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
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
+1
+2
+3
-1
-2
-3
Cations and Anions of Representative Elements
11
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
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
Effective Nuclear Charge (Zeff)
increasing Zeff
increasingZeff
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
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
plots the atomic radii of the elements
against their atomic numbers
19
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
Comparison of Atomic Radii with Ionic Radii
22
Cation is always smaller than atom from which it is
formed.
Anion is always larger than atom from which it is
formed.
23
The Radii (in pm) of Ions of Familiar Elements
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
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
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
General Trends in First Ionization Energies
Increasing First Ionization Energy
IncreasingFirstIonizationEnergy
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
31
Variation of Electron Affinity With Atomic Number (H – Ba)
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.
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
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
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
Group 3A Elements (ns2np1, n  2)
4Al(s) + 3O2(g) 2Al2O3(s)
2Al(s) + 6H+
(aq) 2Al3+
(aq) + 3H2(g)
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
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)
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
Group 5A Elements (ns2np3, n  2)
Nitrogen and Phosphorus: nonmetals
Arsenic and Antimony: metalloids
Bismuth: metal
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
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
All the halogens are nonmetals with the general formula X2,
where X denotes a halogen element.
51
Group 7A Elements (ns2np5, n  2)
Increasingreactivity
X + 1e- X-1
X2(g) + H2(g) 2HX(g)
Noble gases
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.

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Ch8 the periodic table

  • 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
  • 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.
  • 34.
  • 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
  • 36.
  • 37.
  • 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
  • 39.
  • 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.
  • 42.
  • 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.
  • 45.
  • 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.
  • 48.
  • 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)
  • 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.