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❖ Uniqueness principle
❖Diagonal effect
❖Inert-pair effect
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PERIODIC TRENDS OF THE ELEMENTS
❖The way certain properties (chemical and physical) of elements vary
according to their location on the periodic table
❖Some variations in:
Electronegativity increases
Ionization energy increases
Atomic radius decreases
Ionic radius decreases
Ionic radius increase
Effective nuclear charge increases
Atomic radius increase
Effective nuclear charge decreases
Ionization energy decrease
Electron affinity decreases
Electronegativity decreases
Electron affinity increases
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Uniqueness principle
Li
Be
B
C
N
O
F
Ne
The chemistry of the second-period elements
(Li, Be, B, C, N, O, F, Ne) are significantly different from
other elements in their respective groups
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Uniqueness principle
Li
Be
B
C
N
O
F
Ne
Na
Mg
Al
Si
P
S
Cl
Ar
Second element in each group
(Na, Mg, Al, Si, P, S, Cl, Ar) are more representative
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Uniqueness principle
Li
Be
B
C
N
O
F
Ne
Why the first elements of the groups are different from
their congeners (elements of the same group)?
a) The small size of the elements leading to a high polarizing power
and a high degree of covalent character in their compounds
b) The greater probability of π bonds (pπ-pπ)
c) The lack of availability of the d orbitals
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•
⮚ small size of the first
elements
Size effect leads to :
X
-
e
⑧
1) Smaller electron affinities - ability to accept one or more electrons
Electron affinity is the energy change that occurs when an electron
is accepted by an atom in the gaseous state to form an anion
e
X (g) +
=
⑧
X
↓
move
repulsion
will
ē
X-(g)
Electrons added to these small atom,
experience more electron repulsions
form
2) Larger charge densities
3) Enhanced degrees of covalent character in their compounds
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The first electron affinities
of the halogens
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What about fluorine?
> It is a very small atom.
> Incoming electron quite close to the nucleus.
> The existing electron density is very high.
> The extra repulsion is particularly great and
lessens the attraction from the nucleus
enough to lower the electron affinity below that
of chlorine.
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GROUP 1 (ALKALINE METALS) : Li, Na, K, Rb, Cs, Fr
❖ lithium behaves differently than others
▪ small size
▪ high charge density of cation allow it to polarize nearby
anion – allows a large degree of covalency in its bond and
less ionic
▪ less soluble in water and more soluble in polar organic
solvents
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1) Small lithium cation get very close to larger, more diffuse or
filled electron cloud of chloride ion.
↓
ionic,tefat"nonmetal)
Bond4
covalentshare i
2) Electron cloud of chloride ion is distorted or polarized by lithium ion
3) This distortion makes overlap between two ions. Orbital overlap and
sharing electron between two species – characteristic of a covalent bond
overlap between the valence orbitals in Li+ (empty 2s) and Cl- (filled 3p)
is increased
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Bond energies increase
Weakness of the fluorine – fluorine bond
F2
x
X
Cl2
x-x
between
Taran
nucleus
->
-owing to the small atomic size – closeness/ nearness in F2 compared to Cl2
-lone pairs of electrons on adjacent fluorine atoms repel each other (increase
repulsion ) -> macam magnet, clantara nucleus dig ada nucleus dia adajaran antara
each
other
-weakening the F-F bond in F2 molecule
leman
->
sebab
repulsion
banyan
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•
⮚ increased π bonding in the first
element
Due
to small size, increase π bond formation among themselves and with
other elements (capable of forming strong double and triple bond)
•
π bonds involve parallel overlap between for example two p orbitals
(π bonding can occur using d orbitals and antibonding molecular orbitals
of some molecules)
•
Utilizing π bond than σ bond :
•
Parallel orbital overlap or π bonding is more effective in smaller first
elements for example carbon than larger congeners, silicon
neil
o
"
4xde
suiz
bond because
besar
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⮚ lack of availability of d orbitals in the first
elements
•
The lighter elements lack availability of d orbitals and therefore
cannot form compounds with expanded octets
CF4
ClF3
SiF62-
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Diagonal effect
DIAGONAL RELATIONSHIPS exists between the chemistry of the
first member of a group and the second member of the next group
•
Similarities between pairs of
elements in different groups
and periods of the periodic
table
1
2
13
Li
Be
B
Mg
Al
14
Phenomenon happens :
1) small ionic size
2) closeness of the charge
densities of their cations
(charge density = charge of
an ion divided by its radius :
charge nm-1)
3) electronegativity
Si
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Diagonal effect
1
2
properties
13
14
B
C
antara
unsur
satu
menghampin
nilai dia.
Charge of ion
Ionic radius, Å
Charge density
Electronegativity
Charge of ion
Ionic radius, Å
Charge density
Electronegativity
Li
+1
0.73
1.4
1.0
Na
+1
1.13
0.88
0.9
Be
+2
0.41
4.9
1.5
+3
0.25
12
2.0
+4
0.29
14
2.5
Mg
Al
Si
+2
0.71
2.8
1.2
+3
0.53
5.7
1.5
+4
0.40
10
1.8
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Diagonal effect
• Ionic radius of Be2+ (0.41Å) is
more similar to Al3+ (0.53Å) than
Mg2+ (0.71Å)
• High charge density of Al3+ ion
(5.7) and Be2+ ion (4.9)
1
2
13
Li
Be
B
14
• Same electronegativity (1.5)
• Be-X and Al-X : covalent character
• Its small size & high charge
density of cation Be2+ and Al3+ allow
them to polarize the electron cloud
of anion (X atom in M-X bond) to
give additional covalent character
▪ AlH3 resembles BeH2 in its
properties (example of the
diagonal relationship)
0.41Å
4.9
Mg
Al
0.71Å
2.8
0.53Å
5.7
Si
13225-2635238645d4p65s4d5p4652
4f145918
6p6
Electron configuration of elements in Group 14
C: 1s2 2s2 2p2
Si: 1s2 2s2 2p6 3s² 3p²
man
o
43
2
meny unardibuang
untur
Ge: 1s2 2s2 2p6 3s2 3p6 4s2 3d¹⁰ 4p²
v
Sn: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s24d¹⁰5p²
Pb:1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p2
Inert-pair effect
Gebuanl
OE
or:
▪ Oxidation state : +4
CCl4, CO2, SiCl4, SiO2, SnO2
▪ CO easily oxidized to CO2 (+2 to +4)
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C
Si
Ge
Sn
Pb
+4
C
Si
Ge
Sn
Pb
+2
▪ Convert tin(II) to tin(IV) – more stable. Sn2+ ions in solution
suitable as reducing agents
▪ Reversed situation for lead
▪ Moving down a group, there are more
and
more examples oxidation state +2 :
CO, SnCl2, PbO, Pb2+
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Inert-pair effect
•
Get closer to the bottom of the group (heavier elements –
increasing tendency s2 pair not to be used in the bonding
(left unchanged / inert pair)
•
Valence electrons in an s orbital are more tightly bound
are of higher energy than electrons in p orbitals
and therefore less likely to be involved in bonding.
•
Electrons closer to the nucleus – difficult to
remove the heavier the element the greater this effect
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Inert-pair effect
ot
• As an example in group 13 the +1 oxidation state of Tl
is the most stable and TlIII compounds comparatively
rare. The stability Increases in the following sequence:
AlI < GaI < InI < TlI
•
geable
All four elements (Al, Ga, In, Tl) give trivalent compounds but the
univalent state becomes increasingly important for Ga, In and Tl
-
•
Valence ns2 electrons of metallic elements : In, Tl, Sn, Pb, Sb, Bi and
Po are less reactive than expected.
•
Inert ns2 pairs mean oxidation state is 2 less than the expected
group valence for the heavier elements of groups 13, 14 and 15
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Inert-pair effect
•
Two major reasons for this effect:
a) larger than normal effective nuclear charges (higher than expected
ionization energies for Ga, In, Tl)
-energy level
L4s, 5s and 6sIelectrons experience larger effective nuclear charge
than expected – they are more difficult to ionize
b) lower bond energies (as expected) due to increase in atomic size and
bond distance
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Ionization energy decrease
Inert-pair effect
F
cl
Br
I
Ionization energy increases
ATOMIC RADIUS
The increase in size of atom is accompanied by a decrease of ionization energy
-going down a group, size of the atom increases
-the outer electrons lie farther away from the nucleus
-attractive pull from the nucleus on the outer electrons decreases
-easier to pull out an electron from the outer shell of the atom
-untuk
nucleus
(+)
IONIZATION ENERGY
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Slight increase between
tin and lead
Large increase between
tin and lead