4. Atomic Radii
• If a cloud of electrons rotate around the
nucleus of an atom, giving it a spherical
shape, how do we measure accurately the
length of the radius of an atom?
5. Atomic Radii
• If a cloud of electrons rotate around the
nucleus of an atom, giving it a spherical
shape, how do we measure accurately the
length of the radius of an atom?
8. • Well we can’t!
• Using one atom renders the process
impossible.
9. • Well we can’t!
• Using one atom renders the process
impossible.
• We must look at 2 atoms that are
covalently bonded together, e.g. H2.
10. • Well we can’t!
• Using one atom renders the process
impossible.
• We must look at 2 atoms that are
covalently bonded together, e.g. H2.
11. • Well we can’t!
• Using one atom renders the process
impossible.
• We must look at 2 atoms that are
covalently bonded together, e.g. H2.
12. • Well we can’t!
• Using one atom renders the process
impossible.
• We must look at 2 atoms that are
covalently bonded together, e.g. H2.
13. • Well we can’t!
• Using one atom renders the process
impossible.
• We must look at 2 atoms that are
covalently bonded together, e.g. H2.
14. • Well we can’t!
• Using one atom renders the process
impossible.
• We must look at 2 atoms that are
covalently bonded together, e.g. H2.
15. • Well we can’t!
• Using one atom renders the process
impossible.
• We must look at 2 atoms that are
covalently bonded together, e.g. H2.
• The radius of a H atom is calculated by
16. • Well we can’t!
• Using one atom renders the process
impossible.
• We must look at 2 atoms that are
covalently bonded together, e.g. H2.
• The radius of a H atom is calculated by
17. • Well we can’t!
• Using one atom renders the process
impossible.
• We must look at 2 atoms that are
covalently bonded together, e.g. H2.
• The radius of a H atom is calculated by
21. • DEFINITION:
• The atomic radius of an atom is defined as
½ the distance between two nuclei, of two
identical atoms, linked together by a single
covalent bond.
22. • DEFINITION:
• The atomic radius of an atom is defined as
½ the distance between two nuclei, of two
identical atoms, linked together by a single
covalent bond.
23.
24. • The size of the radius of an atom depends on
many factors:
25. • The size of the radius of an atom depends on
many factors:
26. • The size of the radius of an atom depends on
many factors:
1. The nuclear charge (no of p+’s in the nucleus)
27. • The size of the radius of an atom depends on
many factors:
1. The nuclear charge (no of p+’s in the nucleus)
2. The screening effect of inner e-’s
28. • The size of the radius of an atom depends on
many factors:
1. The nuclear charge (no of p+’s in the nucleus)
2. The screening effect of inner e-’s
3. The number of energy levels occupied
29. • The size of the radius of an atom depends on
many factors:
1. The nuclear charge (no of p+’s in the nucleus)
2. The screening effect of inner e-’s
3. The number of energy levels occupied
30. • The size of the radius of an atom depends on
many factors:
1. The nuclear charge (no of p+’s in the nucleus)
2. The screening effect of inner e-’s
3. The number of energy levels occupied
• We can work out the trend in the size of the
radii of atoms of the Periodic Table from our
existing knowledge.
31. • The size of the radius of an atom depends on
many factors:
1. The nuclear charge (no of p+’s in the nucleus)
2. The screening effect of inner e-’s
3. The number of energy levels occupied
• We can work out the trend in the size of the
radii of atoms of the Periodic Table from our
existing knowledge.
• Lets move across a period and down a group.
32.
33. • As a general rule moving across a period on
the table results in a decrease in the atomic
radius of atoms.
34. • As a general rule moving across a period on
the table results in a decrease in the atomic
radius of atoms.
35. • As a general rule moving across a period on
the table results in a decrease in the atomic
radius of atoms.
• As we move across from element to element:
36. • As a general rule moving across a period on
the table results in a decrease in the atomic
radius of atoms.
• As we move across from element to element:
1. An e- is gained in the same energy level
37. • As a general rule moving across a period on
the table results in a decrease in the atomic
radius of atoms.
• As we move across from element to element:
1. An e- is gained in the same energy level
2. The nuclear charge increased by 1p+
38. • As a general rule moving across a period on
the table results in a decrease in the atomic
radius of atoms.
• As we move across from element to element:
1. An e- is gained in the same energy level
2. The nuclear charge increased by 1p+
39. • As a general rule moving across a period on
the table results in a decrease in the atomic
radius of atoms.
• As we move across from element to element:
1. An e- is gained in the same energy level
2. The nuclear charge increased by 1p+
• Increased nuclear charge tightens its grip on
the outer energy level, pulls it inwards and so
decreases the atomic radius of the atom.
40.
41. • As a general rule moving down a group on the
table results in an increase in the atomic radius
of atoms.
42. • As a general rule moving down a group on the
table results in an increase in the atomic radius
of atoms.
43. • As a general rule moving down a group on the
table results in an increase in the atomic radius
of atoms.
• As we move down from element to element:
44. • As a general rule moving down a group on the
table results in an increase in the atomic radius
of atoms.
• As we move down from element to element:
1. An new energy level is gained on the atom
45. • As a general rule moving down a group on the
table results in an increase in the atomic radius
of atoms.
• As we move down from element to element:
1. An new energy level is gained on the atom
2. The screening effect of the inner energy levels
decrease the pull of the nucleus on the outer
energy level as the e-’s in these levels repel
each other.
46. • As a general rule moving down a group on the
table results in an increase in the atomic radius
of atoms.
• As we move down from element to element:
1. An new energy level is gained on the atom
2. The screening effect of the inner energy levels
decrease the pull of the nucleus on the outer
energy level as the e-’s in these levels repel
each other.
47. • As a general rule moving down a group on the
table results in an increase in the atomic radius
of atoms.
• As we move down from element to element:
1. An new energy level is gained on the atom
2. The screening effect of the inner energy levels
decrease the pull of the nucleus on the outer
energy level as the e-’s in these levels repel
each other.
• Extra energy levels and increased screening
effect lessens the grip of the nucleus on the
outer energy level and therefore allows it to
move outwards increasing the atomic radius.
52. Ionisation Energies
• An ion is a charged atom!
• Atoms become charged ions when the gain or lose
e-’s.
53. Ionisation Energies
• An ion is a charged atom!
• Atoms become charged ions when the gain or lose
e-’s.
• e-’s can be removed from atoms by using energy.
54. Ionisation Energies
• An ion is a charged atom!
• Atoms become charged ions when the gain or lose
e-’s.
• e-’s can be removed from atoms by using energy.
• DEFINITION: The minimum energy required to
remove the most loosely held e- from the outer
energy level of an atom of an element is called the
ionisation energy of that atom.
55. Ionisation Energies
• An ion is a charged atom!
• Atoms become charged ions when the gain or lose
e-’s.
• e-’s can be removed from atoms by using energy.
• DEFINITION: The minimum energy required to
remove the most loosely held e- from the outer
energy level of an atom of an element is called the
ionisation energy of that atom.
• Na has one e- in its outer energy level. It wants to
lose this e- so the energy required to remove it
would be very low. Therefore a small ionisation
energy would be required to ionise an atom of Na.
56. Na to Na +
Ionisation Energy required to remove an electron from an atom of Na
57. •The general trend as we move across the Periodic
Table is an increase in ionisation energy.
There are,
however
some
anomolies!!
58.
59. • Notice that there is a drop in ionisation energy as we
move across the second period from Be to B and
from N to O.
60. • Notice that there is a drop in ionisation energy as we
move across the second period from Be to B and
from N to O.
• To see why this is the case we need to look at their
full electronic configuration.
61. • Notice that there is a drop in ionisation energy as we
move across the second period from Be to B and
from N to O.
• To see why this is the case we need to look at their
full electronic configuration.
• Be - 1s2, 2s2
62. • Notice that there is a drop in ionisation energy as we
move across the second period from Be to B and
from N to O.
• To see why this is the case we need to look at their
full electronic configuration.
• Be - 1s2, 2s2
• B - 1s2, 2s2, 2px1
63. • Notice that there is a drop in ionisation energy as we
move across the second period from Be to B and
from N to O.
• To see why this is the case we need to look at their
full electronic configuration.
• Be - 1s2, 2s2
• B - 1s2, 2s2, 2px1
• N - 1s2, 2s2, 2px1, 2py1, 2pz1
64. • Notice that there is a drop in ionisation energy as we
move across the second period from Be to B and
from N to O.
• To see why this is the case we need to look at their
full electronic configuration.
• Be - 1s2, 2s2
• B - 1s2, 2s2, 2px1
• N - 1s2, 2s2, 2px1, 2py1, 2pz1
• O - 1s2, 2s2, 2px2, 2py1, 2pz1
65. • Notice that there is a drop in ionisation energy as we
move across the second period from Be to B and
from N to O.
• To see why this is the case we need to look at their
full electronic configuration.
• Be - 1s2, 2s2
• B - 1s2, 2s2, 2px1
• N - 1s2, 2s2, 2px1, 2py1, 2pz1
• O - 1s2, 2s2, 2px2, 2py1, 2pz1
• Be is more stable than B because fully filled orbitals
are more stable than ½ filled orbitals.
66. • Notice that there is a drop in ionisation energy as we
move across the second period from Be to B and
from N to O.
• To see why this is the case we need to look at their
full electronic configuration.
• Be - 1s2, 2s2
• B - 1s2, 2s2, 2px1
• N - 1s2, 2s2, 2px1, 2py1, 2pz1
• O - 1s2, 2s2, 2px2, 2py1, 2pz1
• Be is more stable than B because fully filled orbitals
are more stable than ½ filled orbitals.
• N is more stable than O since its orbitals of equal
energy px, py and pz are all ½ filled. This is a stable
configuration. O has an unequal distribution of
electrons in px, py and pz and is therefore less
stable.
69. • Notice in the previous diagram, the drop in
energy as we move from Mg to Al and again
as we move from P to S.
70. • Notice in the previous diagram, the drop in
energy as we move from Mg to Al and again
as we move from P to S.
71. • Notice in the previous diagram, the drop in
energy as we move from Mg to Al and again
as we move from P to S.
• To understand why this is happening, we
need to look at the full electronic
configuration of both elements.
72. • Notice in the previous diagram, the drop in
energy as we move from Mg to Al and again
as we move from P to S.
• To understand why this is happening, we
need to look at the full electronic
configuration of both elements.
• Mg-1s2, 2s2, 2px2, 2py2, 2pz2, 3s2
73. • Notice in the previous diagram, the drop in
energy as we move from Mg to Al and again
as we move from P to S.
• To understand why this is happening, we
need to look at the full electronic
configuration of both elements.
• Mg-1s2, 2s2, 2px2, 2py2, 2pz2, 3s2
• Al -1s2, 2s2, 2px2, 2py2, 2pz2, 3s2, 3px1
74. • Notice in the previous diagram, the drop in
energy as we move from Mg to Al and again
as we move from P to S.
• To understand why this is happening, we
need to look at the full electronic
configuration of both elements.
• Mg-1s2, 2s2, 2px2, 2py2, 2pz2, 3s2
• Al -1s2, 2s2, 2px2, 2py2, 2pz2, 3s2, 3px1
75. • Notice in the previous diagram, the drop in
energy as we move from Mg to Al and again
as we move from P to S.
• To understand why this is happening, we
need to look at the full electronic
configuration of both elements.
• Mg-1s2, 2s2, 2px2, 2py2, 2pz2, 3s2
• Al -1s2, 2s2, 2px2, 2py2, 2pz2, 3s2, 3px1
• Since the ionisation energy of Al is lower that
Mg we can conclude that electrons are harder
to remove from full orbitals than from ½ filled
orbitals. Full orbitals are the most stable.
78. • The previous diagram shows that as we
move down group II on the Periodic
Table the ionisation energies show a
general decrease in value.
79. • The previous diagram shows that as we
move down group II on the Periodic
Table the ionisation energies show a
general decrease in value.
80. • The previous diagram shows that as we
move down group II on the Periodic
Table the ionisation energies show a
general decrease in value.
• This is due to 2 factors:
81. • The previous diagram shows that as we
move down group II on the Periodic
Table the ionisation energies show a
general decrease in value.
• This is due to 2 factors:
1. Increased number of energy levels
82. • The previous diagram shows that as we
move down group II on the Periodic
Table the ionisation energies show a
general decrease in value.
• This is due to 2 factors:
1. Increased number of energy levels
2. Increased screening effect from inner e-’s
83. • The previous diagram shows that as we
move down group II on the Periodic
Table the ionisation energies show a
general decrease in value.
• This is due to 2 factors:
1. Increased number of energy levels
2. Increased screening effect from inner e-’s
84. • The previous diagram shows that as we
move down group II on the Periodic
Table the ionisation energies show a
general decrease in value.
• This is due to 2 factors:
1. Increased number of energy levels
2. Increased screening effect from inner e-’s
• The outer electron is easier to remove as
a result of these changes and so the
ionisation energy required to remove the
outermost electron is smaller.
87. • Al = 1s2,2s2,2px2,2py2,2pz2,3s2,3px1
• If we remove electrons from this atom starting at
3px1, then a low energy is needed as it is unpaired.
Unpaired electrons are easy to remove.
• If we try to remove an electron from 3s2, a higher
energy required as it is in a full orbital. Once
electrons are paired within an orbital, it is difficult to
remove them.
• When we enter the 2pz2 orbital and try to remove an
electron there is a huge increase in the ionisation
energy required as this electron belongs to a
88. • Therefore:
• Ionisation energies with an atom increase
slightly when electrons are being removed
from the same energy level.
• Ionisation energies increase greatly once
we enter different energy levels within an
atom. Hence the reason for the ‘steps’ on
the previous graph.
89. Trends in Electronegativity
• Definition: It is a measure of the pulling
power and atom has over a shared pair of
electrons in a covalent bond.
• Taking into consideration the trends on the
periodic table in the values of (i) atomic
radius and (ii) ionisation energies it is not
difficult to predict the trends with
electronegativity values both across and
down the periodic table.
92. • The increase is for two reasons:
1. The number of protons in the nucleus
increases therefore increasing the nuclear
charge.
2. The attraction between the nucleus and
a pair of electrons in a covalent bond
increases and thus the atomic radius
decreases.
93. • Electronegativity values decrease as we
go down a group on the periodic table.
Look at Group II:
94. • The decrease is for two reasons:
1.Addition of a new energy level to the atom
2.Increased screening effect of the full
energy levels on the electrons in the
covalent bond on the outer energy level.
95. Trends Within Groups
• We learned in Junior Cert Science that
families of elements on the Periodic Table
show similar chemical properties.
• With group II and VII especially the trends
in the reactivity of the elements in these
families is worth studying.
