class 11 chemistry CBSE,JEE,NEET

1. Classification of Elements

2. Why Do We Need Classification?
• With the passage of time, new elements are being
discovered.
To study them in an orderly manner and to predict their
properties easily, we require classification.
The following scientists have made some significant
contributions in the achieving of the same: • Johann
Wolfgang Dobereiner • John Newlands • Mendeleev •
Henry Moseley

3. Classification of Elements
With the discovery of a large number of elements,
it became difficult to study the elements
individually, so classification of elements was
done to make the study easier.
Earlier Attempts to Classify Elements
Many attempts were made to classify the known
elements from time to time. The earlier attempts
are as follows:

4. Dobereiner’s Triads (1829)
Dobereiner classified the elements into groups of three elements with similar
properties in such a manner so that the atomic weight of the middle element
was the arithmetic mean of the other two, e.g.,
Element Li NA K
Atomic
weight
7 23 39
Mean of atomic masses = (7 + 39) / 2 = 23
Similarly CI, Br, I; Ca, Sr, Ba are two more examples of such triads.
Limitations Dobereiner could not arrange all the elements known at that time
into triads. He could identify only three such triads that have been mentioned.

6. Newland’s Octaves (1864)
Newland states that when elements are arranged in order of increasing atomic masses,
every eighth element has properties similar to the first just like in the musical note
[Every eighth musical note 1S the same as the first mentioned note]. This can be
illustrated as given below
Sa re ga ma pa dha ni
Li Be B C N O F
Na Mg AI Si P S CI
Limitations
1. This classification was successful up to the element calcium.
2. When noble gas elements were discovered at a later stage, their inclusion in
these octaves disturbed the entire arrangement.

7. Mendeleef’s Periodic Table
Mendeleefs Periodic Table is based upon Mendeleefs periodic law which
states ‘The physical and chemical properties of the elements are a
periodic function of their atomic masses.”
At the time of Mendeleef, only 63 elements were known.
This Periodic Table is divided into seven horizontal rows (periods) and
eight vertical columns (groups). Zero group was added later on in the
modified Mendeleefs Periodic Table.

8. Importance of Mendeleers Periodic Table
Few important achievements of Periodic Table are
1 Systematic study of the elements.
2. Prediction of new elements and their properties. he left
space for the elements yet to be discovered. e.g., he left
spaces for Ga and Ge and named these elements as ERa-
aluminium (Ga) and EKa-silicon (Ge) respectively
3. Atomic mass correction of doubtful elements on the basis of
their expected positions and properties.

11. Defects in the Mendeleev's Periodic Table
(i) Position of hydrogen Hydrogen has been placed in group IA (alkali
metals). but it also resembles with halogens of group VIlA. Thus. its position in
the Mendeleev's Periodic Table is controversial.
(ii) Position of isotopes As Mendeleev's classification is based on atomic
weight, Isotopes would have to be placed in different positions due to their
different atomic weights, e.g., 1H1 2H1 3H1 would occupy different positions.
(iii) Anomalous positions of some elements Without any proper justification.
in some cases the element with higher atomic mass precedes the element
with lower atomic mass. For example, AI (atomic weight = 39.9) precedes K
(atomic weight = 39.1) and similarly Co (atomic weight. = 58.9) has been
placed ahead of Ni (atomic weight = 58.7).
(iv) Position of Lanthanoids and actinoids Lanthanoids and actinoids were
not placed in the main Periodic Table.

12. Modern Periodic Table (1913)
Moseley modified Mendeleefs periodic law. He stated “Physical
and chemical properties of elements are the periodic function of
their atomic numbers.” It is known as modern periodic law and
considered as the basis of Modern Periodic Table.
When the elements were arranged in increasing order of atomic
numbers, it was observed that the properties of elements were
repeated after certain regular intervals 01 2, 8, 8, 18, 18 and 32.
These numbers are called magic numbers and cause of
periodicity in properties due to repetition of similar electronic
configuration.

13. Structural Features of Long Form of Periodic Table
1. Long form of Periodic Table is called Bohr’s Periodic Table. There arc 18
groups and seven periods in this Periodic Table
2. The horizontal rows are called periods.
•First period (1H – 2He) contains 2 elements. It is the shortest period
•Second period (3Li – 10Ne) and third period (11 Na – 18Ar) contain is
elements each. These are short periods.
•Fourth period (19K – 36Kr) and fifth period (37Rb – 54Xe) contain 18
elements each. These are long periods.
•Sixth period (55Cs – 86 Rn) consists of 32 elements and is the longest period.
•Seventh period starting with 87Fr is incomplete and consists of 19 elements.

15. 3. The 18 vertical columns are known as groups.
•Elements ot group 1arc called alkali metals.
•Elements of group 2 are called alkaline earth metals.
•Elements of group 16 are called chalcogens [ore
forming elements].
•Elements of group 17 are called halogens. [sea salt
forming]
•Elements of group 18 are called noble gases.

16. 4. The Periodic Table is divided into foul’ main blocks (s, p, d and f)
depending upon the subshell to which the valence electron enters into.
1.S-block Elements:
Elements in which the last electron enters the s-orbitals.
General outer shell electronic configuration of s-block elements – ns1-2 where n = 2-7
General characteristics of s-block elements
1.They are soft metals with low melting and boiling points
2.They have low ionization enthalpies (energies) and are highly electro positive.
3.They lose the valence (outermost) electrons readily to form +1 (in case of alkali
metals) and +2 ions (in case of alkaline earth metals).
4.They are very reactive metals. The metallic character and the reactivity increase as
we move down the group. because of high reactivity, they are never found pure in
nature.
5.The compounds of s-block elements with the exception of those of beryllium are
predominantly ionic.

17. 2.P-Block Elements:
Elements in which the last electron enters any one of three p-orbitals of their
respective outermost shells are called P-Block elements.
General outer shell electronic configuration of p-block elements –
ns2 mp1-6 where n = 2-7
General characteristics of p-block elements
1.P-block elements include both metal and non-metals but the number of non-
metals is much higher than that of metals. Further, the metallic character increases
from top to bottom within a group and non-metallic character increase from left to
right along a period in this block.
2.Their ionization enthalpies are relatively higher as compared to those s-block
elements.
3.They mostly form covalent compounds.
4.Some of the show more than one (variable) oxidation states in their compounds.
5.Their oxidizing character increase from left to right in a period and reducing
character increase from top to bottom in a group.

18. 3.D-block Elements:
Elements in which the last electron enters any one of the five d-orbitals
of their respective penultimate shells are called d-block elements.
General outer shell electronic configuration of d-block elements –
(n-1) d1-10 ns0-2, where n = 4 – 7.
General Characteristics of d-block elements –
1.They are hard, malleable (i.e. can be converted into sheets) and
ductile (i.e. can be drawn into wires) metals with high melting ad boiling
points.
2.They are good conductor of heat and electricity.
3.They ionization enthalpies are between s- and p- block elements.
4.They show variables oxidation states.
5.They form both ionic and covalent compounds.

19. 4.F-block elements:
Elements in which the last electron entre any one of the
seven f-orbitals of their respective ante-penultimate shells are
called f-block elements.
General outer shells electronic configuration of f-block
elements –
(n – 2) f0-14 (n – 1) d0-2 ns2
General Characteristics of f-block elements
1.They have generally high melting and boiling points
2.They show variable oxidation states
3.Their compounds are generally coloured
4.They have heavy metals
5.They have a high tendency to form complexes.

20. 4.F-block elements:
Elements in which the last electron entre any one of the
seven f-orbitals of their respective ante-penultimate shells
are called f-block elements.
General outer shells electronic configuration of f-block
elements –
(n – 2) f0-14 (n – 1) d0-2 ns2
General Characteristics of f-block elements
1.They have generally high melting and boiling points
2.They show variable oxidation states
3.Their compounds are generally coloured
4.They have heavy metals
5.They have a high tendency to form complexes.

21. Atomic Properties
Atomic radii: It is the distance from the centre of nucleus to outermost shell of an atom.
’
Covalent radii:
we can say covalent radii is one half of the inter nuclear
distance between the bonded atoms.
For example in H2 molecule the bond length is 74pm
therefore covalent radii is 37pm.
Metallic radii: It is considered only for metals.
It is defined as one half the intern nuclear distances
between two neighboring metals in lattice.
For example, in Cu metal, the bond length is 256pm
therefore radii are128pm

22. Variation along period:
Along period the size keeps on decreasing
due to increase in nuclear charge.
Nuclear charge is the attraction of positive
charged protons towards electron.
•As we move towards right the atomic no.
increase so as no. of protons increase,
attraction increase. So the order is that it
keeps on decreasing
For example, in 3rdperiod: the order is
Li>Be>B>C>N>O>F
Variation along group:
As we move down there is an addition of shell along group it increases
•It is because every time a new shell is added so the effect of addition of shell is more than the
pronounced nuclear charge therefore the order as we discussed is increasing down the group.
For example: in first group the order is Li<Na<k<Rb<Cs
•In the whole periodic table alkali group is the largest group according to size

23. ionic radii: it is defined as one half of the internuclear distance between the bonded atoms.
Or
it is the measure of distance between cation and anion in ionic crystals
For example: in NaCl the bond length is 276pm but if we take exactly the half then there will be
error so, we calculate distance separately for sodium ion and chloride ion and define the ionic radii
as: the effective distance from the centre of nucleus of ion to up to the distance where it has
influence in ionic bond.
Cation is always smaller than parent atom as it is formed by removal of electron due to which the
same nuclear charge acts on lesser number of electrons therefore the nuclear charge increases
and the size decreases

24. Anion is always bigger than parent atom as it is formed by gaining electrons due to which the
same nuclear charge acts on more number of electrons and moreover the magnitude
of repulsions increase ,therefore the nuclear charge decrease and size increase
Isoelectronic species: they are those that have same number of electrons
For example: O2-,F-,Na+etc. are isoelectronic
Vander wall radii: it is the half of internuclear distance between the adjacent atoms of substance
belonging to two different molecules.
It is present in inert gases
Ionization energy
It is the amount of energy required to remove electron from valence shell of isolated gaseous atom.

25. There are certain factors on which ionization energy depends:
1.Size of atom: If the size is small, nuclear charge will be high so it will
attract electron more effectively therefore ionization energy will be high.
This means ionization energy is inversely proportion to size of an atom.
2.Charge on nucleus: If nuclear charge is more than attraction for
electron will be more therefore ionization energy will be high or vice
versa
3.Screening effect: Due to filling of inner orbital’s the nuclear charge is
somehow reduced for outermost electron. As a result outermost
electron will be loosely bounded and therefore ionization energy
decreases.
4.Penetration effect: The orbital’s which are closer to nucleus will
experience high nuclear charge so the penetrating effect will be more.
The order it follows is:

26. S>P>d>f
5.5. Electronic configuration: Half-filled and fully filled orbitals are more stable therefore
they have high ionization energies.
Variation along period:
Along period it increases It is because along period the size
decreases, nuclear charge increase therefore ionization
energy increases but the anomalous behavior is seen let us
see that:
Li<Be>B<C<N<O<F<Ne
In this Be has high ionization energy than B because of fully
filled shell N has high ionization energy than O because N is
half filled ,more stable therefore ionization energy is high.
Variation along group:
Ionization energy decreases along the group as the size
increases, nuclear charge decreases as a result less
ionization energy is required to remove the electron.
For example in 1st group the order is:
Li>Na>K>Rb>Cs out of these cesium has largest size and
least ionization energy

27. Electron gain enthalpy
It is the amount of energy released when
electron is added to outermost shell of an atom.
This energy release depends upon the extent of
attraction for an incoming electron.
In case of inert gases they have almost no
attraction Because of stable configuration.
Therefore it is negative for inert gases this
means energy should be supplied in order to
make them accept electron.
This energy can be mathematically shown as:
M + e-à M– (negatively charged ion)

28. •Factors affecting electron gain enthalpies:
1.Atomic radius : more is the size ,less is the nuclear charge therefore less attraction for incoming
electron therefore electron gain enthalpy is less negative for those atoms or vice versa
2.Nuclear charge : less is the nuclear charge therefore less attraction for incoming electron therefore
electron gain enthalpy is less negative for those atoms or vice versa
3.Electronic configuration: atoms with stable electronic configurations have less negative electron gain
enthalpies.
Variation along group
•Electron gain enthalpy becomes less negative as we move down the group:
•As we down the group, the size increases and nuclear charge decreases due to which the attraction for
incoming electron decreases so as electron gain enthalpy decreases.
•But certain anomalies are seen like out of chlorine and fluorine, fluorine has less negative electron
enthalpy because of its extreme small size the incoming electron suffers interelectronic repulsions
therefore electron gain enthalpy is less negative as compared to chlorine
Variation along period
•Along period it increases the reason being the size decreases and nuclear charge increases.
•But still the exceptional behavior is seen like Electron gain enthalpy of inert gases is positive because of
their stable electronic configuration.

29. Electro negativity
It is the tendency of an atom to attract shared pair of
electrons more towards itself. This property depends
upon the size and the nuclear charge.
If the size is less then more will be the nuclear
charge hence more attraction for shared pair of
electrons.
In whole periodic table fluorine is maximum
electronegative due to its smallest size.
Other factors on which electro negativity depends:
State of hybridization:
The order is sp>sp2>sp3Oxidation state of element:
Higher is the oxidation state more it is
electronegative.
Variation along group:
Along group it decreases as size
increases and nuclear charge
decreases

30. Variation along period:
Along period it increases as size decreases and nuclear charge increases therefore
the attraction for shared pair of electrons increaseApplications of electronegativity:
It tells us about the metallic and non metallic character of atom: more is the electro
negativity lesser is the metallic character and more is the non metallic character or
vice versa
Polar or non-polar: if the difference in the electro negativity between the two bonded
atoms is more than the bond is polar that is it has lost but if electro negativity
difference is less or zero than the bonded non polar.

31. Metallic and non-metallic character
Metallic character is tendency to lose electron.
Non-metallic character: It is the tendency to gain electron
Along group – metallic character increase and non metallic character decreases
Size increases i.e. valence electrons become more away from nucleus due to which
nuclear charge decrease and metallic character increases whereas non metallic
character decreases.
For example
Li. Na K Rb Cs Fr
(least (Most
Metallic) metallic)

32. Along periods – Metallic character decreases and non metallic character increases.
As we move the Size decreases along period due to which Nuclear Charge increases
therefore tendency to gain electron increases and to lose electron decreases.
Na Mg Al Si P S Cl
Metallic character decreases
Non-metallic character increases