PCM
Periodic Table
DOBEREIGNER'S LAW OF TRIADS
In the history of the periodic table, Döbereiner's triads were an early attempt to sort the elements into some logical order by their physical properties.
Dobereiner noticed a pattern with certain elements that had similar chemical and physical properties. He called these elements triads.
If you put these elements in order of their atomic masses, the average of the molar mass of the first and third elements in the triad is the molar mass of the second element.
Newland's law of octaves
Law of octaves, in chemistry, the generalization made by the English chemist Newlands that, if the chemical elements are arranged according to increasing atomic weight, those with similar physical and chemical properties occur after each interval of seven elements.
Mendeleef's periodic table
Mendeleev realized that the physical and chemical properties of elements were related to their atomic mass in a 'periodic' way, and arranged them so that groups of elements with similar properties fell into vertical columns in his table.
Gaps and predictions Sometimes this method of arranging elements meant there were gaps in his horizontal rows or 'periods'. But instead of seeing this as a problem, Mendeleev thought it simply meant that the elements which belonged in the gaps had not yet been discovered. He was also able to work out the atomic mass of the missing elements, and so predict their properties. And when they were discovered, Mendeleev turned out to be right. For example, he predicted the properties of an undiscovered element that should fit below aluminum in his table. When this element, called gallium, was discovered in 1875 its properties were found to be close to Mendeleev's predictions. Two other predicted elements were later discovered, lending further credit to Mendeleev's table.
MODERN PERIODIC TABLE
(+ve)
In it, elements are arranged according to thier atomic number and classified according to electronic configuration.
It has 7 periods(horizontal), 18 groups(vertical) and 4 blocks.
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The first period is the shortest period, that contains two elements.
-
2nd and 3rd periods are short and contain 8 elements.
-
the 4th and 5th periods are long and contain 18 elements.
-
the 6th period is the longest period, which contains 32 elements.
-
The 7th period is incomplete.
BLOCKS OF THE PERIODIC TABLE
There are 4 blocks in the modern periodic table:
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The S Block
-
The P Block
-
The D Block
-
The F Block
S BLOCK.
In S Block Elements, the last electron is present in the S subshell.
The general electronic configuration is:
(ns) i.e group 1,2
1,2
ns
1
Alkali Metals
ns
2
Alkaline
Earth
Metals
NOTE
HELIUM WILL BE WRITTEN WITH P BLOCK ELEMENTS AS IT HAS PROPERTIES SIMILAR TO THE NOBLE GASSES.
S Block is present in the leftmost section of the periodic table.
Note: They're called alkaline earth metals because they cant easily form an alkali but thier oxides can.
p BLOCK.
In P Block elements, the last electron is present in the P Subshell. The general electronic configuration of P-Block elements is:
ns np
2
1-6
boron family
carbon family
Nitrogen family
Chalkogens
Halogens
Noble Gasses
NS2NP1
NS2NP2
NS2NP3
NS2NP4
NS2NP5
NS2NP6
P Block is present in the leftmost section of the periodic table.
D BLOCK.
In D Block elements, the last electron is present in the penultimate shell. The general electronic configuration of D-Block elements is:
ns (n-1)d
0-10
0,1,2
-
It starts from 4th period of periodic table.
-
It's placed in between S&P shell.
-
They're called transition elements.
-
They are all metals with thier last 2 shells incompletely filled.
3 b
4b
NS2(n-1)D1
5 b
6b
7b
8b
9b
10b
11b
NS1(n-1)D10
12b
NS2(n-1)D10
There are exceptions to the rule, that are:Cromium, Cobalt, Palladium, etc.
The S,D,F block elements are mostly metals.
F BLOCK.
The configuration for F block elements is:
ns (n-1)d (n-2)f
2
0,1
1-14
-
Present in the 6th and 7th period of the periodic table
-
placed in group III B and written seperately at the bottom of the table
6s2 5d1 4f1
Lanthanides
6s2 5d1 4f14
Actanides
7s2 5d1 4f1
7s2 5d1 4f14
s BLOCK.
Gp1- Alkali metals
Gp2- Alkaline earth metals
d BLOCK.
Gp(3-11)-transition
Gp12-metals
f BLOCK.
Gp12-rare earth metals
*p block contains metals, non metals and mettaloids*
Valency & oxidation state
s&p block
-
For 2nd period, valency with respect to Hydrogen first increases then decreases.
LiH BeH BH CH
NH
0H
FH
2
3
4
3
2
-
For 3nd period, the pattern of valency with respect to Hydrogen remains the same as the second period. However, oxides increase from left to right.
NaH MgH AlH SiH
PH
SH
ClH
2
3
4
3
2
Na2O Al2O3 Si2O4 P2O5
S2O6 Cl2O7
S block elements show a fixed oxidation state.
-
Alkali Metals always show +1 state in thier compound whereas alkaline earth metals always show +2 state in thier compound.
-
P block elements can show variable oxidation states, depending on the element with which they are combined with.
Oxidation State
How many oxygens can 2 atoms of that substance take, also number of valence electrons.
INERT PAIR EFFECT
is observed in 4th, 5th and 6th periods of group 13, 14 and 15.
ns np ----> ns np
2
2
1
3
excited
Group 14 elements show a valency of 4. Give Reasons.
Covalent bonds are made by unpaired electrons. Grp14 elements have a general configuration of ns2np2 but before bonding, due to excitation one electron jumps to np into an empty orbital therefore 4 unpaired electrons will now exist.
However, as we go down the group, this excitation becomes more and more difficult and so the tendency of the ns electron pair to act as the inert pair increases.
This is called the inert pair effect due to which variable valencies are observed in the same group.
Ga < Ga Ge < Ge As < As
Sn -- Sn In --- In Sb -- Sb
Pb > Pb Tl > Tl Bi > Bi
2+
4+
2+
4+
2+
4+
+
+
+
3+
3+
3+
3+
3+
3+
5+
5+
5+
for d block
all the ns electrons and some of the (n-1)d electrons act as Valence Electrons.
The number of D electrons participating in bonding depends on bonding conditions. It could vary from 0 - all, therefore d block elements show VARIABLE VALENCIES.
All lanthanides show +3 as thier mosts stable oxidation state, with the ecxceptions of Eu and Yb which show 1 2
Effective Nuclear Charge (Zeffective)
Shielding effect
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The repulsion caused by inner electrons.
-
The net repulsion is called as shielding constant and is calculated by slater's rule.
Zeffective= Z-
shielding constant
Z
actual nuclear charge
Calculate
Step 1 write the electronic configuration
Step 2 group the ns,np shells according to principal quantum number (n) and nd and nf will be a seperate group.
Step 3 all the electrons to the right give 0 contibution to it.
Step 4 on the left, in its own group, all electrons except itself will contribute 0.35 each and in 1s will contribute 0.30 each
Step 5 the (n-1)th group will contribute 0.85 each
Step 6 the (n-2)th group will contribute 1.0 each
Step 7 In D,F shells, all lower shells contribute 1.0 electrons
ATOMIC RADIUS
.
.
nucleus
outermost electron
As electrons are waves and distance cannot be measured, we make 3 types of radiuses and make assumptions based on them:
covalent radius
Covalent radius is defined as half of internuclear distance between 2 singly bonded homonuclear diatomic molecules.
shortest
metallic radius
Half of the internuclear distance between two adjacent atoms in a closely packed solid structure
between van der waal and covalent
van der waal radius
Half of the internuclear distance between two adjacent atoms in a solid state
largest
FACTORS AFFECTING ATOMIC RADIUS
-
Z R
-
R
DETAILED:
S&P BLOCK
As we move from left to right in a period, atomic radius decreases due to increase in effective nuclear charge. Exceptions are observed in group 18 - the reported value of noble gasses is more than that of halogens.
This exception occurs as instead of the covalent radius, the van der waal's radius is taken instead. Otherwise there is hypothetically no exception.
Top to Bottom Radius increases due to increase in n . Exceptions are observed in Group 13:
B < Ga -- Al < In -- Tl
Before Gallium, 3rd subshell is filled. These D electrons provide a poor shielding effect to p electrons of gallium and hence Zeff increases causing a decrease in radius which almost counterbalances the increase in radius caused by extra shell.
Therefore, the radius of Al, Ga are almost the same.
Similarly, before thallium, 4s is filled which provides a poor shielding which counterbalances increase in radius due to extra shell hence radius of thallium is almost equal to Indium.
D&F block
3D Series
As we move from Sc to Mn, the configurations of d1 to d5 provides poor shielding to 4s electrons and hence radius decreases.
From Fe to Co configurations of d6 to d8 provide a better shielding and hence radius remains almost the same .
From Cu to Zn configuration of d10 provides an even better shielding and radius increases.
Lanthanide Contraction
In F block, in Lanthanide Series, radius continuously decreases from C to Lu due to poor shielding of S electron. This is called Lanthanide contraction.
*Radius of Eu and Yb are exceptionally high*
As we move from 3d to 4d series, size increases due to number of shells. However, as we move from 4d to 5d radius of Lanthanium is more than Yithium, but radius of all the other 5d elements is almost the same as 4d elements due to lanthanide contraction.
ELECTRONEGATIVITY
The ability of an atom to pull a covalently bonded electron pair towards itself.
It is a comparitive property.
Due to electronegativity difference between the atoms having a covalent bond, a partial positive and partial negative charge is developed in its covalent bond, making it partially ionic, called polar covalent bond.
PAULING SCALE
|ENX-ENH|=0.208 HX
hx= beh-x -beh-hBEx-x
