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Classification of Elements & Periodicity in Properties Class 11 Notes Chemistry Chapter 3

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 Page 1


The first systematic classification of elements was provided by Russian chemist 
D.I. Mendeleev.
1. Mendeleev's periodic law
“The physical and chemical properties of elements are periodic functions
of their atomic weight.”
2. It was modified to Modern Periodic law :
“The physical and chemical properties of elements are periodic functions
of their atomic numbers.”
It is the long form of periodic table :
7 Horizontal rows are called Periods and 18 Vertical columns are called
Group
Group-1 are called Alkali metals Group-2 are called Alkaline earth metals.
Group-15 are called Pnicogens Group-16 are called Chalcogens
Group-17 are called Halogens Group-18 are called Noble gases
3. 1
st
 period—2 elements 2
nd
 and 3
rd
 period—8 elements
4
th
 and 5
th
 period—18 elements 6
th
 period—32 elements
  7
th
 period—Incomplete (32 elements)
4.Groups
1 and 2 — ‘ s’ block elements last electron entered in ‘ s’ subshell [ s
1
, s
2
]
3 to 12 — ‘d’ block elements last electrons entered in ‘d’ subshell [d
1
 to d
10
].
13 to 18 — ‘p’ block elements last electrons enter in ‘p ’ subshell [p
1
 to p
6
].
Two f-block series lanthanoids and actinoids are placed in the bottom of
periodic table.
5. (A) In ‘s’ and ‘p’ block elements the electrons enters in outer most shell.
In ‘d’ block elements the electron enters in the penultimate shell ( n – 1).
Page 2


The first systematic classification of elements was provided by Russian chemist 
D.I. Mendeleev.
1. Mendeleev's periodic law
“The physical and chemical properties of elements are periodic functions
of their atomic weight.”
2. It was modified to Modern Periodic law :
“The physical and chemical properties of elements are periodic functions
of their atomic numbers.”
It is the long form of periodic table :
7 Horizontal rows are called Periods and 18 Vertical columns are called
Group
Group-1 are called Alkali metals Group-2 are called Alkaline earth metals.
Group-15 are called Pnicogens Group-16 are called Chalcogens
Group-17 are called Halogens Group-18 are called Noble gases
3. 1
st
 period—2 elements 2
nd
 and 3
rd
 period—8 elements
4
th
 and 5
th
 period—18 elements 6
th
 period—32 elements
  7
th
 period—Incomplete (32 elements)
4.Groups
1 and 2 — ‘ s’ block elements last electron entered in ‘ s’ subshell [ s
1
, s
2
]
3 to 12 — ‘d’ block elements last electrons entered in ‘d’ subshell [d
1
 to d
10
].
13 to 18 — ‘p’ block elements last electrons enter in ‘p ’ subshell [p
1
 to p
6
].
Two f-block series lanthanoids and actinoids are placed in the bottom of
periodic table.
5. (A) In ‘s’ and ‘p’ block elements the electrons enters in outer most shell.
In ‘d’ block elements the electron enters in the penultimate shell ( n – 1).
‘f ’ block elements last electron enter the antepenultimate shell (n – 2).
(B) ‘f ’ block elements are placed in between ‘d’ block elements.
‘f ’ block elements in 2 rows [4f lanthanoids, 5f actinoids]
6. General outer electronic configuration
‘s’ block :  ns
1
, ns
2
 [Group 1 to 2]
‘p’ block :  ns
1
np
1
 to ns
2 
np
6
 Group 13 to 18
‘d’ block :  ns
0–2
 (n – 1) d
1 to 10 
Group 3 to 12
‘f ’ block :  (n – 2)f 
1 to 14 
(n – 1)d
0, 1 
ns
2
7. General periodic trends in properties of elements
Atomic Radius
(A) Left to right decreases due to effect of successive increasing nuclear change 
without addition of a new shell.
(B) From top to bottom atomic radius increases due to successive addition of 
shell.
(C) Noble gases have large radius than group 17 due to complete filling of 
electron in outer shell electron-electron repulsion mildy increases.
Covalent radius
It is half of the distance between the centre of nuclei of two adjacent similiar 
atoms which are bonded to each other by single covalent bond.
van der waal's radius
van der waal's radius is defined as one-half the distance be tween the centres 
of nuclei of two nearest like atoms belonging to two adjacent molecules of the 
element in the solid state.
Metallic radius
Half of the distance between the centres of the nuclei of two adjacent atoms in the 
metallic crystal. A comparision of the three atomic radii show that van der waal's 
radius is maximum while the covalent radius has the least value.
van der waal's radius > Metallic radius > Covalent radius
Ionic radius (A) Cation radius < Atomic radius—due to more no. of protons than 
Page 3


The first systematic classification of elements was provided by Russian chemist 
D.I. Mendeleev.
1. Mendeleev's periodic law
“The physical and chemical properties of elements are periodic functions
of their atomic weight.”
2. It was modified to Modern Periodic law :
“The physical and chemical properties of elements are periodic functions
of their atomic numbers.”
It is the long form of periodic table :
7 Horizontal rows are called Periods and 18 Vertical columns are called
Group
Group-1 are called Alkali metals Group-2 are called Alkaline earth metals.
Group-15 are called Pnicogens Group-16 are called Chalcogens
Group-17 are called Halogens Group-18 are called Noble gases
3. 1
st
 period—2 elements 2
nd
 and 3
rd
 period—8 elements
4
th
 and 5
th
 period—18 elements 6
th
 period—32 elements
  7
th
 period—Incomplete (32 elements)
4.Groups
1 and 2 — ‘ s’ block elements last electron entered in ‘ s’ subshell [ s
1
, s
2
]
3 to 12 — ‘d’ block elements last electrons entered in ‘d’ subshell [d
1
 to d
10
].
13 to 18 — ‘p’ block elements last electrons enter in ‘p ’ subshell [p
1
 to p
6
].
Two f-block series lanthanoids and actinoids are placed in the bottom of
periodic table.
5. (A) In ‘s’ and ‘p’ block elements the electrons enters in outer most shell.
In ‘d’ block elements the electron enters in the penultimate shell ( n – 1).
‘f ’ block elements last electron enter the antepenultimate shell (n – 2).
(B) ‘f ’ block elements are placed in between ‘d’ block elements.
‘f ’ block elements in 2 rows [4f lanthanoids, 5f actinoids]
6. General outer electronic configuration
‘s’ block :  ns
1
, ns
2
 [Group 1 to 2]
‘p’ block :  ns
1
np
1
 to ns
2 
np
6
 Group 13 to 18
‘d’ block :  ns
0–2
 (n – 1) d
1 to 10 
Group 3 to 12
‘f ’ block :  (n – 2)f 
1 to 14 
(n – 1)d
0, 1 
ns
2
7. General periodic trends in properties of elements
Atomic Radius
(A) Left to right decreases due to effect of successive increasing nuclear change 
without addition of a new shell.
(B) From top to bottom atomic radius increases due to successive addition of 
shell.
(C) Noble gases have large radius than group 17 due to complete filling of 
electron in outer shell electron-electron repulsion mildy increases.
Covalent radius
It is half of the distance between the centre of nuclei of two adjacent similiar 
atoms which are bonded to each other by single covalent bond.
van der waal's radius
van der waal's radius is defined as one-half the distance be tween the centres 
of nuclei of two nearest like atoms belonging to two adjacent molecules of the 
element in the solid state.
Metallic radius
Half of the distance between the centres of the nuclei of two adjacent atoms in the 
metallic crystal. A comparision of the three atomic radii show that van der waal's 
radius is maximum while the covalent radius has the least value.
van der waal's radius > Metallic radius > Covalent radius
Ionic radius (A) Cation radius < Atomic radius—due to more no. of protons than 
number of electron coloumbic force increases, size decreases.
[Mg
2+
    <      Mg
+
   <    Mg]
 (B) Anion radius > Atomic radius—Due to more number of electron than 
number of protons 
  [N
3–
    >    O
2–  
 >   F
–
]
 Electron-Electron repulsion increase, coloumbic force of attraction 
decreases.
 (C) For Isoelectronic species—More is the charge of cation lesser the size.
 More is the charge of anion, more is the size.
(D) Size — O
2–
  > F
–
 > Na  > Na
+
  > Mg
2+
8. (A) Ionization energy :
  The minimum amount of energy which is required to remove the most 
loosely bound electron from an isolated atom in the gaseous state is called 
Ionization enthalpy.
M
(g)   
+  Energy    — ? M
+
    +   e
–
IE
3
  >   IE
2
    >    IE
1
Variation of I.E along a period:
  Ionization energy increase along the period because atomic radii decrease 
and nuclear charge increase along the period.
I  ionization enthalpy  Li < B < Be < C < O < N < F < Ar
 II ionization enthalpy  Be < C < B < N < F < O < Ne
Variation down the group:
Inozation energy decrease down the group because atomic radius increase 
down the group.
  Metallic behaviour : Decrease from left to right due to increase in 
ionization enthalpy.
  Non metallic  behaviour : Increase from left to right due to more number 
of electron in outershell and added electron goes towards nucleus.
(9)   Screening effect or shielding effect:-
 It is the decrease in the force of attraction between nucleus and outermost 
electron due to presence of inner shell electrons. As a result, the outer most 
Page 4


The first systematic classification of elements was provided by Russian chemist 
D.I. Mendeleev.
1. Mendeleev's periodic law
“The physical and chemical properties of elements are periodic functions
of their atomic weight.”
2. It was modified to Modern Periodic law :
“The physical and chemical properties of elements are periodic functions
of their atomic numbers.”
It is the long form of periodic table :
7 Horizontal rows are called Periods and 18 Vertical columns are called
Group
Group-1 are called Alkali metals Group-2 are called Alkaline earth metals.
Group-15 are called Pnicogens Group-16 are called Chalcogens
Group-17 are called Halogens Group-18 are called Noble gases
3. 1
st
 period—2 elements 2
nd
 and 3
rd
 period—8 elements
4
th
 and 5
th
 period—18 elements 6
th
 period—32 elements
  7
th
 period—Incomplete (32 elements)
4.Groups
1 and 2 — ‘ s’ block elements last electron entered in ‘ s’ subshell [ s
1
, s
2
]
3 to 12 — ‘d’ block elements last electrons entered in ‘d’ subshell [d
1
 to d
10
].
13 to 18 — ‘p’ block elements last electrons enter in ‘p ’ subshell [p
1
 to p
6
].
Two f-block series lanthanoids and actinoids are placed in the bottom of
periodic table.
5. (A) In ‘s’ and ‘p’ block elements the electrons enters in outer most shell.
In ‘d’ block elements the electron enters in the penultimate shell ( n – 1).
‘f ’ block elements last electron enter the antepenultimate shell (n – 2).
(B) ‘f ’ block elements are placed in between ‘d’ block elements.
‘f ’ block elements in 2 rows [4f lanthanoids, 5f actinoids]
6. General outer electronic configuration
‘s’ block :  ns
1
, ns
2
 [Group 1 to 2]
‘p’ block :  ns
1
np
1
 to ns
2 
np
6
 Group 13 to 18
‘d’ block :  ns
0–2
 (n – 1) d
1 to 10 
Group 3 to 12
‘f ’ block :  (n – 2)f 
1 to 14 
(n – 1)d
0, 1 
ns
2
7. General periodic trends in properties of elements
Atomic Radius
(A) Left to right decreases due to effect of successive increasing nuclear change 
without addition of a new shell.
(B) From top to bottom atomic radius increases due to successive addition of 
shell.
(C) Noble gases have large radius than group 17 due to complete filling of 
electron in outer shell electron-electron repulsion mildy increases.
Covalent radius
It is half of the distance between the centre of nuclei of two adjacent similiar 
atoms which are bonded to each other by single covalent bond.
van der waal's radius
van der waal's radius is defined as one-half the distance be tween the centres 
of nuclei of two nearest like atoms belonging to two adjacent molecules of the 
element in the solid state.
Metallic radius
Half of the distance between the centres of the nuclei of two adjacent atoms in the 
metallic crystal. A comparision of the three atomic radii show that van der waal's 
radius is maximum while the covalent radius has the least value.
van der waal's radius > Metallic radius > Covalent radius
Ionic radius (A) Cation radius < Atomic radius—due to more no. of protons than 
number of electron coloumbic force increases, size decreases.
[Mg
2+
    <      Mg
+
   <    Mg]
 (B) Anion radius > Atomic radius—Due to more number of electron than 
number of protons 
  [N
3–
    >    O
2–  
 >   F
–
]
 Electron-Electron repulsion increase, coloumbic force of attraction 
decreases.
 (C) For Isoelectronic species—More is the charge of cation lesser the size.
 More is the charge of anion, more is the size.
(D) Size — O
2–
  > F
–
 > Na  > Na
+
  > Mg
2+
8. (A) Ionization energy :
  The minimum amount of energy which is required to remove the most 
loosely bound electron from an isolated atom in the gaseous state is called 
Ionization enthalpy.
M
(g)   
+  Energy    — ? M
+
    +   e
–
IE
3
  >   IE
2
    >    IE
1
Variation of I.E along a period:
  Ionization energy increase along the period because atomic radii decrease 
and nuclear charge increase along the period.
I  ionization enthalpy  Li < B < Be < C < O < N < F < Ar
 II ionization enthalpy  Be < C < B < N < F < O < Ne
Variation down the group:
Inozation energy decrease down the group because atomic radius increase 
down the group.
  Metallic behaviour : Decrease from left to right due to increase in 
ionization enthalpy.
  Non metallic  behaviour : Increase from left to right due to more number 
of electron in outershell and added electron goes towards nucleus.
(9)   Screening effect or shielding effect:-
 It is the decrease in the force of attraction between nucleus and outermost 
electron due to presence of inner shell electrons. As a result, the outer most 
33
electrons does not feel full charge of the nucleus. The actual charge felt by 
an electron is called effective Nuclear charge.
Shielding effect is in the following order
  s   >    p   >   d   >   f
d & f subshell show weak sheilding effect because their orbital size are large 
and are more diffused.
(10) Isoelectronic species:
Ions of different elements which have the same number of electons but different 
no. of protons are called isoelectronic ions.
Na
+
 Mg
2+
 Al
3+
 N
3–
 O
2–
 F
–
 
No. of Protons 11 12 13 7 8 9
No. of electrons 10 10 10 10 10 10
Ionic Radii Al
3+
  < Mg
2+
 < Na
+
  < F
–
   < O
2–
  < N
3–
(11) Electron gain enthalpy:
The enthalpy change when an extra electron is added to neutral gaseous atom to 
form anion. 
E(g) +  e
–
  ???   E
– 
(g)
¦ Trends—From left to right—Increase due to decrease in size, more attraction
of added electron by nucleus.
¦ From top to bottom—Decreases as the added electron is away from nucleus
due to increase in size.
¦ Cl has more negative electron gain enthalpy than fluorine— Due to small
size of fluorine extra added electron has more inter electronic repulsion than
chlorine which has large size.
¦ Similarly Phosphorus and Sulphur have negative electron gain enthalpy than
nitrogen and oxygen respectively.
¦ Maximum electron gain enthalpy—Chlorine (in periodic table)
¦ Electron gain enthalpy—
Halogen > Oxygen > Nitrogen > Metal of group 1 and 13 and non metal of
group 14 > metal of group 2.
¦ 2nd electron gain enthalpy is always positive.
Page 5


The first systematic classification of elements was provided by Russian chemist 
D.I. Mendeleev.
1. Mendeleev's periodic law
“The physical and chemical properties of elements are periodic functions
of their atomic weight.”
2. It was modified to Modern Periodic law :
“The physical and chemical properties of elements are periodic functions
of their atomic numbers.”
It is the long form of periodic table :
7 Horizontal rows are called Periods and 18 Vertical columns are called
Group
Group-1 are called Alkali metals Group-2 are called Alkaline earth metals.
Group-15 are called Pnicogens Group-16 are called Chalcogens
Group-17 are called Halogens Group-18 are called Noble gases
3. 1
st
 period—2 elements 2
nd
 and 3
rd
 period—8 elements
4
th
 and 5
th
 period—18 elements 6
th
 period—32 elements
  7
th
 period—Incomplete (32 elements)
4.Groups
1 and 2 — ‘ s’ block elements last electron entered in ‘ s’ subshell [ s
1
, s
2
]
3 to 12 — ‘d’ block elements last electrons entered in ‘d’ subshell [d
1
 to d
10
].
13 to 18 — ‘p’ block elements last electrons enter in ‘p ’ subshell [p
1
 to p
6
].
Two f-block series lanthanoids and actinoids are placed in the bottom of
periodic table.
5. (A) In ‘s’ and ‘p’ block elements the electrons enters in outer most shell.
In ‘d’ block elements the electron enters in the penultimate shell ( n – 1).
‘f ’ block elements last electron enter the antepenultimate shell (n – 2).
(B) ‘f ’ block elements are placed in between ‘d’ block elements.
‘f ’ block elements in 2 rows [4f lanthanoids, 5f actinoids]
6. General outer electronic configuration
‘s’ block :  ns
1
, ns
2
 [Group 1 to 2]
‘p’ block :  ns
1
np
1
 to ns
2 
np
6
 Group 13 to 18
‘d’ block :  ns
0–2
 (n – 1) d
1 to 10 
Group 3 to 12
‘f ’ block :  (n – 2)f 
1 to 14 
(n – 1)d
0, 1 
ns
2
7. General periodic trends in properties of elements
Atomic Radius
(A) Left to right decreases due to effect of successive increasing nuclear change 
without addition of a new shell.
(B) From top to bottom atomic radius increases due to successive addition of 
shell.
(C) Noble gases have large radius than group 17 due to complete filling of 
electron in outer shell electron-electron repulsion mildy increases.
Covalent radius
It is half of the distance between the centre of nuclei of two adjacent similiar 
atoms which are bonded to each other by single covalent bond.
van der waal's radius
van der waal's radius is defined as one-half the distance be tween the centres 
of nuclei of two nearest like atoms belonging to two adjacent molecules of the 
element in the solid state.
Metallic radius
Half of the distance between the centres of the nuclei of two adjacent atoms in the 
metallic crystal. A comparision of the three atomic radii show that van der waal's 
radius is maximum while the covalent radius has the least value.
van der waal's radius > Metallic radius > Covalent radius
Ionic radius (A) Cation radius < Atomic radius—due to more no. of protons than 
number of electron coloumbic force increases, size decreases.
[Mg
2+
    <      Mg
+
   <    Mg]
 (B) Anion radius > Atomic radius—Due to more number of electron than 
number of protons 
  [N
3–
    >    O
2–  
 >   F
–
]
 Electron-Electron repulsion increase, coloumbic force of attraction 
decreases.
 (C) For Isoelectronic species—More is the charge of cation lesser the size.
 More is the charge of anion, more is the size.
(D) Size — O
2–
  > F
–
 > Na  > Na
+
  > Mg
2+
8. (A) Ionization energy :
  The minimum amount of energy which is required to remove the most 
loosely bound electron from an isolated atom in the gaseous state is called 
Ionization enthalpy.
M
(g)   
+  Energy    — ? M
+
    +   e
–
IE
3
  >   IE
2
    >    IE
1
Variation of I.E along a period:
  Ionization energy increase along the period because atomic radii decrease 
and nuclear charge increase along the period.
I  ionization enthalpy  Li < B < Be < C < O < N < F < Ar
 II ionization enthalpy  Be < C < B < N < F < O < Ne
Variation down the group:
Inozation energy decrease down the group because atomic radius increase 
down the group.
  Metallic behaviour : Decrease from left to right due to increase in 
ionization enthalpy.
  Non metallic  behaviour : Increase from left to right due to more number 
of electron in outershell and added electron goes towards nucleus.
(9)   Screening effect or shielding effect:-
 It is the decrease in the force of attraction between nucleus and outermost 
electron due to presence of inner shell electrons. As a result, the outer most 
33
electrons does not feel full charge of the nucleus. The actual charge felt by 
an electron is called effective Nuclear charge.
Shielding effect is in the following order
  s   >    p   >   d   >   f
d & f subshell show weak sheilding effect because their orbital size are large 
and are more diffused.
(10) Isoelectronic species:
Ions of different elements which have the same number of electons but different 
no. of protons are called isoelectronic ions.
Na
+
 Mg
2+
 Al
3+
 N
3–
 O
2–
 F
–
 
No. of Protons 11 12 13 7 8 9
No. of electrons 10 10 10 10 10 10
Ionic Radii Al
3+
  < Mg
2+
 < Na
+
  < F
–
   < O
2–
  < N
3–
(11) Electron gain enthalpy:
The enthalpy change when an extra electron is added to neutral gaseous atom to 
form anion. 
E(g) +  e
–
  ???   E
– 
(g)
¦ Trends—From left to right—Increase due to decrease in size, more attraction
of added electron by nucleus.
¦ From top to bottom—Decreases as the added electron is away from nucleus
due to increase in size.
¦ Cl has more negative electron gain enthalpy than fluorine— Due to small
size of fluorine extra added electron has more inter electronic repulsion than
chlorine which has large size.
¦ Similarly Phosphorus and Sulphur have negative electron gain enthalpy than
nitrogen and oxygen respectively.
¦ Maximum electron gain enthalpy—Chlorine (in periodic table)
¦ Electron gain enthalpy—
Halogen > Oxygen > Nitrogen > Metal of group 1 and 13 and non metal of
group 14 > metal of group 2.
¦ 2nd electron gain enthalpy is always positive.
(12) Electro negativity:
The tendency of an atom to attract the shared pair of electron towards itself in a 
bonded state.
# Fluorine is the most electronegative element in the periodic table.
# Cesium is the least electronegative element in the periodic table.
# Electronegativity is decrease down the group and increase along the period.
Difference between electron gain enthalpy and Electronegativity.
Electron gain enthalpy is the energy, but electronegativity is not the energy, it is 
only the tendency of an atom in a molecule to attract the shared pair of electrons. 
Three highest electronegative atoms F > O > N. 
Maximum electronegative Assign to F.
* Lightest element : — Hydrogen
* Lightest metal : — Lithium
* Heaviest metal (highest density) : — Osmium
* Most reactive metal : — Caesium
* Most reactive nonmetal : — Fluorine
* Most malleable metal : — Gold
* Electrically best conductor : — Silver
* Metals which are relatively volatile : — Zn, Cd, Hg
* Strongest reducing agent in aqueous solution : — Lithium
* Strongest oxidising agent : — Fluorine
* The element of lowest ionization energy : — Caesium
* The element of highest ionization energy : — Helium
* The most electronegative element : — Fluorine
* The element of highest electron gain enthalpy : — Chlorine
* The group containing most electropositive metals : — Group 1
* The group containing most electronegative metals : — Halogens Group 17
* The group containing maximum number of gaseous elements : — Group 18
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FAQs on Classification of Elements & Periodicity in Properties Class 11 Notes Chemistry Chapter 3

1. What is the periodic table and how is it organized?
Ans. The periodic table is a tabular arrangement of chemical elements, organized based on their atomic number, electron configuration, and recurring chemical properties. It is organized in rows called periods and columns called groups. The elements are placed in increasing atomic number from left to right and elements with similar properties are grouped together.
2. What is the significance of the periodic table in chemistry?
Ans. The periodic table is significant in chemistry as it provides a systematic way to classify and organize the elements. It helps in understanding the trends and patterns in the physical and chemical properties of elements, predicting their behavior, and identifying relationships between elements. It is a fundamental tool for studying and understanding the principles of chemistry.
3. What are the main trends observed in the periodic table?
Ans. There are several main trends observed in the periodic table. These include atomic size (decreasing from left to right across a period and increasing down a group), ionization energy (increasing from left to right across a period and decreasing down a group), electronegativity (increasing from left to right across a period and decreasing down a group), and metallic character (decreasing from left to right across a period and increasing down a group).
4. How do elements in the periodic table form chemical compounds?
Ans. Elements in the periodic table form chemical compounds through the sharing, gaining, or losing of electrons to achieve a stable electron configuration. Atoms of different elements combine together to form molecules or ionic compounds. The number and arrangement of electrons in the outermost energy level of an atom, known as the valence electrons, play a crucial role in determining the type of chemical bonds that can be formed.
5. How does the periodic table help in predicting the properties of elements?
Ans. The periodic table helps in predicting the properties of elements by identifying trends and patterns. Elements in the same group exhibit similar chemical properties due to the same number of valence electrons. By examining the position of an element in the periodic table, we can make predictions about its reactivity, atomic size, ionization energy, and other properties based on the trends observed in the elements of similar groups or periods.
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