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CBSE NOTES CLASS 11 CHEMISTRY CHAPTER 3

Periodic Classification Of Elements

Need for Classification

Dobereiner’s Law of Triads

Newland’s Law of Octaves

Mendeleev’s Periodic Table

Achievements of Mendeleev’s Periodic Table

Limitations of Mendeleev’s Classification

Modern Periodic Law

Modern Periodic Table/Long Form of Periodic Table or Bohr’s Table

Cause of periodicity

Nomenclature of Element with Z > 100

Trends in Modern Periodic Table

Electronic Configurations of Elements and The Periodic Table

Metallic and non-metallic character

Atomic size or radius

Ionisation Enthalpy

Shielding Effect

Electron Gain Enthalpy

Valence Electrons
Periodicity In Chemical Properties

Valency or Oxidation State

Anomalous Properties of Second Period Elements

Periodic Trends and Chemical Reactivity

CBSE NOTES CLASS 11 CHEMISTRY CHAPTER 3

Periodic Classification Of Elements

Classification means identifying similar species and grouping them together.

Need for Classification – There are 118 elements and each of them has innumerable compounds. It is very difficult to study and understand all the elements and their compounds individually. Hence there is a need for classification or grouping of elements based on some common properties, so that the study of elements and their compounds could be made easy.

Lavoisier divided elements into two main types known as metals and non-metals.

Dobereiner’s Law of Triads

For some groups of three elements,

For example atomic masses of Li, Na and K are respectively 7, 23 and 39, thus the mean of atomic masses of 1st and 3rd element is equal to the mass of the second one.

Group A Elements

Atomic Masses

Group B Elements

Atomic Masses

Group C Elements

Atomic Masses

Li

7

Ca

40

Cl

35.5

Na

23

Sr

88

Br

80

K

39

Ba

137

I

127

Limitations of Doberiner’s Triads

He could identify only a few such triads and so the law could not gain importance. In the triad of Fe, Co, Ni, all the three elements have a nearly equal atomic mass and thus do not follow the above law.

Newland’s Law of Octaves

According to this law

sa (do)

re (re)

ga (mi)

ma (fa)

pa (so)

da (la)

ni (ti)

H

Li

Be

B

C

N

O

F

Na

Mg

Al

Si

P

S

Cl

K

Ca

Cr

Ti

Mn

Fe

Co, Ni

Cu

Zn

Y

In

As

Se

Br

Rb

Sr

Ce, La

Zr

Drawback of Newland’s Law of Octaves:

Thus, Newland’s Law of Octaves worked well with lighter elements only.

Mendeleev’s Periodic Table

Mendeleev arranged 63 elements, known at that time, in the periodic table. According to Mendeleev

Basis of Classification

The formulae of the hydrides and oxides formed by an element were treated as the basic properties of an element for its classification.

He selected hydrogen and oxygen because they are very reactive and formed compounds with most elements. If the formula for hydride was RHx, then the element R will be placed in xth group.

Structure of Mendeleev’s Periodic Table

Achievements of Mendeleev’s Periodic Table

  1. The arrangement of elements in groups and periods made the study of elements quite systematic. That is, if properties of one element in a particular group are known, those of the others in the group can be easily predicted.

  2. Prediction of new elements and their properties - Many gaps were left in this table for undiscovered elements and properties of these elements could be predicted in advance from their expected position. Mendeleev named them by prefixing a Sanskrit numeral, Eka (one) to the name of preceding element in the same group. For instance, scandium, gallium and germanium, discovered later, have properties similar to Eka–boron, Eka–aluminium and Eka–silicon, respectively.

  3. Correction of doubtful atomic masses: Mendeleev corrected the atomic masses of certain elements with the help of their expected positions and properties.

  4. When inert gases were discovered they were placed in a new group without disturbing the order.

Limitations of Mendeleev’s Classification

  1. He could not assign a correct position to hydrogen in his periodic table, as the properties of hydrogen resemble both with alkali metals as well as with halogens.

  2. The isotopes of the same element will be given different position if atomic mass is taken as basis. But the properties of isotopes of the same element are same.

  3. The atomic masses do not increase in a regular manner in going from one element to the next. At some places the order of atomic masses in the table was changed in order to justify the chemical and physical nature. For example, cobalt (atomic mass 58.9) appeared before nickel (atomic mass 58.7).

  4. Some dissimilar elements were grouped together while some similar elements were placed in different groups. For example Li, Na, K are grouped with Cu, Ag, Au and Cu and Hg are placed in different groups.

  5. Group viii elements have been put in three triads without any justification.

  6. Lanthanoids and Actinoids have not been provided any position in the table and put separately without any justification.

Modern Periodic Law

This law was given by Henry Moseley in 1913.

Henry Moseley observed regularities in the characteristic X-ray spectra of the elements. A plot of ν (where ν is frequency of X-rays emitted) against atomic number (Z) gave a straight line and not the plot of ν vs atomic mass. He thereby concluded that the atomic number is a more fundamental property of an element than its atomic mass.

The law states,

Modern Periodic Table/Long Form of Periodic Table or Bohr’s Table

Moseley proposed the modern periodic table also known as long form of the periodic table.

  1. There are 18 groups and 7 periods in the table. The groups are numbered 1, 2, 3 …. 18 from left to right and the periods from 1 to 7 from top to bottom

  2. The elements belonging to a particular group make a family and usually named after the first member.

  3. In a group all the elements contain the same number of valence electrons.

  4. In a period all the elements contain the same number of shells, but as we move from left to right the number of valence shell electrons increases by one unit.

  5. The maximum number of electrons that can be accommodated in a shell is 2n2 where ‘n’ is the number of the given shell from the nucleus.

Cause of periodicity

Periodicity may be defined as the repetition of the properties of the elements separated by certain definite gap of atomic numbers.

Chemical bonding depends on the number of electrons in the valence shell of an element. Due to same number of electrons in the outermost shell of elements in the same group, the chemical properties are similar.

Since nucleus (mass) resides deep inside the atom, it does not take part in chemical reactions. On the other hand the electrons are outside, in the orbits, hence they interact with other atoms.

Nomenclature of Element with Z > 100

Digit

Name

Abbreviation

0

nil

n

1

un

u

2

bi

b

3

tri

t

4

quad

q

5

pent

p

6

hex

h

7

sept

s

8

oct

o

9

enn

e

Atomic Number

Name according

to IUPAC

Symbol

IUPAC Official Name

IUPAC Symbol

101

Unnilunium

Unu

Mendelevium

Md

102

Unnilbium

Unb

Nobelium

No

103

Unniltrium

Unt

Lawrencium

Lr

104

Unnilquadium

Unq

Rutherfordium

Rf

105

Unnilpentium

Unp

Dubnium

Db

106

Unnilhexium

Unh

Seaborgium

Sg

107

Unnilseptium

Uns

Bohrium

Bh

108

Unniloctium

Uno

Hassium

Hs

109

Unnilennium

Une

Meitnerium

Mt

110

Ununnillium

Uun

Darmstadtium

Ds

111

Unununnium

Uuu

Rontgenium

Rg

112

Ununbium

Uub

Copernicium

Cn

113

Ununtrium

Uut

*

114

Ununquadium

Uuq

Flerovium

Fl

115

Ununpentium

Uup

*

116

Ununhexium

Uuh

Livermorium

Lv

117

Ununseptium

Uus

*

118

Ununoctium

Uuo

*

Trends in Modern Periodic Table

The trends observed in some important properties of the elements in moving down the group (from top to bottom of the table) and across a period (from left to right in a period) are discussed below.

Electronic Configurations of Elements and The Periodic Table

Electronic Configurations across Periods

The period indicates the value of n for the outermost or valence shell. (Take examples)

Electronic Configurations in Groups

Elements can be classified into four blocks viz., s-block, p-block, d-block and f-block depending on the type of atomic orbitals that are being filled with electrons. (Take examples)

s-block elements: Group-1 (Alkali metals) and Group-2 elements (Alkaline earth metals) which respectively have ns1 and ns2 outermost electronic configurations

p-block elements: They belongs to group 13 to 18. The outer most electronic configuration is ns2 np1-6.

Exception - He (1s2) is an s-block element but is positioned with the group 18 elements (ns2 np6) because it has completely filled valence shell and as a result, exhibits properties characteristic of other noble gases.

Group 15 – Pnictogen

Group 16 – Chalcogens

Group 17 – Halogens

Group 18 – Noble gases

d-block elements (Transition elements) are the elements of group 3 to 12 having outer electronic configuration (n-1)d1-10 ns1-2. Four transition series are 3d, 4d, 5d and 6d. The 6d series is incomplete.

f-block elements (Inner- transition Series)

Lanthanoids charecterised by the filling of 4f orbitals, are the elements following lanthanum from 58Ce to 71Lu. The are called 4f-inner transition series. Actinoids characterised by filling of 5f orbitals, are the elements following actinium from 70Th to 103Lr. Characteristic outer electronic configuration is (n-2)f1-14(n-1) d0-1ns2. Thesy are called 5f-inner transition series.

Noble Gases: The gaseous elements of group 18 are called noble gases. The general outermost electronic configuration of noble gases (except He) is ns2 np6. He exceptionally has 1s2 configuration. Thus the outermost shell of noble gases is completely filled.

Metallic and non-metallic character

A zig-zag line separates metals from non-metals. The borderline elements – boron, silicon, germanium, arsenic, antimony, tellurium and polonium – are intermediate in properties and are called metalloids or semi-metals.

As the effective nuclear charge acting on the valence shell electrons increases across a period, the tendency to lose electrons will decrease. Down the group, the effective nuclear charge experienced by valence electrons is decreasing because the outermost electrons are farther away from the nucleus. Therefore, these can be lost easily. Hence metallic character decreases across a period and increases down a group.

These trends also help to predict the nature of oxides formed by the elements because it is known that the oxides of metals are basic and that of non-metals are acidic in general.

Atomic size or radius

The distance between the centre of nucleus of an isolated atom to its outermost shell containing electrons, is called its atomic radius.

Covalent Radius

Half the inter-nuclear distance between two similar atoms of any element which are covalently bonded to each other by a single covalent bond is called covalent radius.

Van Der Waals’ Radius

Half the inter-nuclear separation between two similar adjacent atoms belonging to the two neighbouring molecules of the same substance in the solid state is called the van der waals’radius of that atom.

Metallic Radius

Half the distance between the nuclei of the two adjacent metal atoms in a close packed lattice of the metal is called its metallic radius.

Van der Waals’ radius  >   Metallic radius  >  Covalent radius

Trend

Exceptions

N

O

F

70pm

74pm

72pm

  1. In case of N to O, inter electron repulsions outweigh the effect of increased nuclear charge, hence r(N) < r(O).

  2. In case of O to F, nuclear charge increases by 1. 2 of the p-orbitals now have 2 electrons each, which repel each other. However the increased nuclear charge outweighs the repulsion, hence r(F) < r(O).
Ionic Radius

Cation is formed by removal of an electron from an atom.

Anion is formed by gain of an electron by an atom.

Iso-Electronic Species

Atoms or ions which have same number of electrons are called iso-electronic species.

For example

Species

O2-

F-

Na+

Mg2+

No. of Protons

8

9

11

12

No. of Electrons

10

10

10

10

Size of Mg2 < Na+ < F- < O2-

Ionisation Enthalpy

The energy required to remove an electron from an isolated gaseous atom (X) in its ground state, is called ionization enthalpy.

M(g) + ΔiH1 → M+(g) + e-

The successive ionization enthalpies of loosing electrons are called 1st iH1), 2ndiH2) and 3rdiH3) ionization enthalpies.

M+(g) + ΔiH2 → M2+(g) + e-

The term ionization enthalpy is taken for the first ionization enthalpy.

SI unit is kJ mol-1.

Ionisation Enthalpy is always +ve.

Shielding Effect

Effective nuclear charge (Zeff) experienced by valance electrons will be less than the actual charge due to shielding or screening.

Shielding is more if the inner shells are completely filled. Hence Alkali (1st group) experience greater shielding and can give away valance electrons easily.

Shielding effect of s-orbital is more than that of p-orbital.

Trend

Exceptions

Be-B

Boron (Z=5) has slightly less electron enthalpy than that of Beryllium because shielding by s-orbital electrons is more than by p-orbital electrons due the shape and size.

N-O

In case nitrogen the 3 p electrons reside in different orbitals whereas in case of oxygen 2 of them share one p-orbital, resulting in inter electron repulsion. Hence ionization enthalpy of oxygen is smaller than that of nitrogen.

Electron Gain Enthalpy

The electron gain enthalpy ((∆egH) is the molar enthalpy change when an isolated gaseous atom or ion in its ground state adds an electron to form the corresponding anion.

X(g)  +  e- → X-(g) + ∆egH

The ∆egH may be positive or negative.

The successive values for the addition of second, third etc. electron, are called second (∆egH2), third (∆egH3) etc. electron gain enthalpies. For example,

X(g) + e- → X-(g),  ∆H = ∆egH1 is called first electron gain enthalpy

X(g) + e- → X2-(g),   ∆H = ∆egH2 is called second electron gain enthalpy

X2-(g) + e- → X3- (g)   ∆H = ∆egH3 is called third electron gain enthalpy

Exceptions

egH(O) is less negative than ∆egH(S) and ∆egH(F) is less negative than ∆egH(Cl).

Explantion – The electrons added to O and F occupy 2p orbital which is smaller orbit and inter electron repulsion is higher, whereas in case of S and Cl the orbitals are 3p, the size of orbitals is much larger and inter electron repulsion is lesser.

Factors Affecting Electron Gain Enthalpy

Electronegativity

The relative tendency of an atom to attract the shared electron pair of electrons towards itself is called electronegativity.

Electronegativity is not a measureable quantity. However there are numerical scales developed, e.g. Paling, Mullikan etc.)

Electronegativity of an element is not fixed. It varies depending on the element to chich it is bound.

It provides the nature of force between bonding atoms.

Electronegativity is directly related to that non-metallic properties of elements and inversely related to the metallic properties (or electropositivity) of elements.

Trend

Valence Electrons

The electrons present in outermost shell are called valence electron.

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Summary of Trends In Physical Properties
Periodicity In Chemical Properties

Valency or Oxidation State

The number of electrons which an atom loses or gains or shares with other atom to attain the noble gas configuration is termed as its valency.

Anomalous Properties of Second Period Elements

The first element of each of the groups 1 (lithium) and 2 (beryllium) and groups 13-17 (boron to fluorine) differs in many respects from the other members of their respective group.

For example, Lithium and Beryllium form covalent compounds, whereas the other members of these groups form ionic compounds.

The behaviour of lithium and beryllium is more similar with the second element of the following group i.e., magnesium and aluminium, respectively. This sort of similarity is commonly referred to as diagonal relationship in the periodic properties.

Reasons of Anomalous Behaviour/Diagonal Relationship

Periodic Trends and Chemical Reactivity