ICSE Revision Notes for Periodic Properties, Periodic Table & Variations of Properties Class 10 Chemistry


Chapter Name

Periodic Properties, Periodic Table & Variations of Properties

Topics Covered

  • Modern Periodic Table 
  • Electronic Configuration in Periods
  • Electronic Configurations and Types of Elements
  • Periodic Trends in Physical Properties 

Related Study

Modern Periodic Table 

Mendeleev made a successful effort in grouping elements in the form of his periodic table. He had many achievements, but there were many limitations in his Periodic Table as well.

Some limitation of Mendeleev's periodic table are listed below:

  1. The position of hydrogen was not justified in Mendeleev's Periodic Table.
  2. The discovery of isotopes revealed another limitation of Mendeleev's periodic table.
  3. Although Mendeleev arranged the elements in the increasing order of their atomic masses, there were instances where he had placed an element with a slightly higher atomic mass before an element with a slightly lower atomic mass.

The limitations of Mendeleev’s periodic table forced scientists to believe that atomic mass could not be the basis for the classification of elements. 

In 1913, Henry Moseley demonstrated that atomic number (instead of atomic mass) is a more fundamental property for classifying elements. The atomic number of an element is equal to the number of protons present in an atom of that element. Since the number of protons and electrons in an atom of an element is equal, the atomic number of an element is equal to the number of electrons present in a neutral atom. 

Atomic number = Number of protons = Number of electrons 

The number of protons or electrons in an element is fixed. No two elements can have the same atomic number. Hence, elements can be easily classified in the increasing order of their atomic numbers. In the light of this fact, Mendeleev’s Periodic Law was done away with. As a result, the modern periodic law came into the picture. 

The table that is obtained when elements are arranged in the increasing order of their atomic numbers is called the Modern Periodic Table or Long Form of the Periodic  Table as shown in the figure.


The Modern periodic table 

In the modern periodic table, the elements are arranged in rows and columns. These rows and columns are known as periods and groups respectively. The table consists of 7 periods and 18 groups. 

Do You Know: 

  • In the modern periodic table, hydrogen is placed above alkali metals because of resemblance with their electronic configurations. However, it is never regarded as an alkali metal. This makes hydrogen a unique element. 
  • If you look at the modern periodic table, you will find that all elements in the same group contain the same number of valence electrons. Let us see the following activity to understand better. 

Activity 1: Look at group two of the modern periodic table. Write the name of the first three elements followed by their electronic configurations.

What similarity do you observe in their electronic configurations? How many  valence electrons are present in these elements? 

  • The first three elements of group two are beryllium, magnesium, and calcium. All these elements contain the same number of valence electrons. The number of valence electrons present in these elements is 2. On the other hand, the number of shells increases as we go down the group. 
  • Again, if you look at periods in the modern periodic table, you will find that all elements in the same period contain the same valence shell. Let us see the following activity to understand better. 

Activity 2: Look at the elements of the third period of the modern periodic table. Write the electronic configuration of each element and calculate the number of valence electrons present in these elements. 

What do you observe from the given activity? Do these elements contain the  same number of shells? How many valence electrons are present in these elements? 

  • You will find that elements such as sodium, magnesium, aluminium, silicon, phosphorus, sulphur, chlorine, and argon are present in that period. The valence shell in all these elements is the same, but they do not have the same number of valence electrons. 

Name of the element

Electronic configuration (K, L, M)

Sodium

2, 8, 1

Magnesium

2, 8, 2

Aluminium

2, 8, 3

Silicon

2, 8, 4

Phosphorous

2, 8, 5

Sulphur

2, 8, 6

Chlorine

2, 8, 7

Argon

2, 8, 8

Thus, the number of electrons in the valence shell increases by one unit as the atomic number increases by one unit on moving from left to right in a period. 

Let us calculate the number of elements that are present in the first, second, third, and fourth periods.

The maximum number of electrons that a shell can hold can be calculated using the formula 2n2. Here, n represents the number of shells from the nucleus. For example, n is equal to 1, 2, and 3 for K, L, and M shells respectively. Hence, the maximum number of electrons that each of these shells can hold can be calculated by substituting the value of n in the given formula. 

Number of electrons that K shell can accommodate = 2n2 

= 2×12

= 2 

Hence, K shell can accommodate only 2 electrons and only two elements are present in the first period. 

Similarly, the second and third shell (L and M respectively) can accommodate 8 and 18 electrons respectively. Since the outermost shell can contain only 8 electrons, there are only 8 elements in both the periods. 

Important Note: 

The position of an element in the Modern Periodic Table tells us about its chemical reactivity. The valence electrons determine the kind and the number of bonds formed by an element. 

IUPAC Nomenclature for Elements with Atomic Number > 100

Latin word roots for various digits are listed in the given table.

Notation for IUPAC Nomenclature of Elements

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
  • Latin words for various digits of the atomic number are written together in the order of digits, which make up the atomic number, and at the end, ‘ium’ is added.
  • Nomenclature of elements with the atomic number above 100 is listed below.

Nomenclature of Elements with Atomic Number Above 100

Atomic number

Name

Symbol

IUPAC official Name

IUPAC Symbol

101

Unilunium

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

Hassnium

Hs

109

Unnilennium

Une

Meitnerium

Mt

110

Ununnilium

Uun

Darmstadtium

Ds

111

Unununnium

Uuu

Rontgenium

Rg

112

Ununbium

Uub

 

 

113

Ununtrium

Uut

 

 

114

Ununquadium

Uuq

 

 

115

Ununpentium

Uup

 

 

116

Ununhexium

Uuh

 

 

117

Ununseptium

Uus

 

 

118

Ununoctium

Uuo

 

 

Electronic Configuration in Periods

Electronic Configuration and the Periodic Table 

  • Period indicates the value of ‘n’ (principal quantum number) for the outermost or valence shell. 
  • Successive periods in the periodic table are associated with the filling of the next higher principal energy level (n = 2, n = 3, etc). 
  • First period (n = 1) → hydrogen (1s1) and helium (1s2) [2 elements] 
  • Second period (n = 2) → Li (1s2 2s1), Be (1s2 2s2), B (1s2 2s2 2p1) to Ne (2s2 2p6) [8 elements] 
  • Third period (n = 3) → filling to 3s and 3p orbitals gives rise to 8 elements (Na to Ar) • Fourth period (n = 4) → 18 elements (K to Kr) − filling of the 4s and 4p orbitals 3d orbital is filled up before 4p orbitals (3d orbitals → energetically favourable) • 3d-transition series → Sc (3d1 4s2) to Zn (3d10 4s2) 
  • Fifth period (n = 5) → 18 elements (Rb to Xe) 
  • 4d-transition series starts at Ytterbium and ends at Cadmium. 
  • Sixth period (n = 6) → 32 elements; electrons enter 6s, 4f, 5d, and 6p orbitals successively. Elements from Z = 58 to Z = 71 are called 4f-inner transition series or lanthanoid series (filling up of the 4f orbitals).
  • Seventh period (n = 7) → electrons enter at 7s, 5f, 6d, and 7p orbitals successively. Filling up of 5f orbitals after Ac (Z = 89) gives 5f-inner transition series or the actinoid series. 

Electronic Configuration in Groups

  • Same number of electrons is present in the outer orbitals (that is, similar valence shell electronic configuration). 
  • Electronic configuration of group 1 elements is given in the following table.

Atomic number

Symbol

Electronic configuration

3

Li

1s22s1(or) [He]2s1

11

Na

1s2 2s2 2p6 3s1 (or) [Ne]3s1

19

K

1s2 2s2 2p6 3s2 3p6 4s1 (or) [Ar] 4s1

37

Rb

1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 5s1 (or) [kr] 5s1

55

Cs

1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 4d10 5s2 5p6 6s1 (or) [Xe]6s1

87

Fr

[Rn]7s1

Electronic Configurations and Types of Elements

• s- Block Elements 

  • Group 1 (alkali metals) − ns1(outermost electronic configuration)
  • Group 2 (alkaline earth metals) − ns2(outermost electronic configuration) 
  • Alkali metals form +1 ion and alkaline earth metals form +2 ion.
  • Reactivity increases as we move down the group.
  • They are never found in the pure state in nature. (Reason − they are highly reactive) 

• p - Block Elements 

  • Elements belonging to Groups 13 to 18
  • Outermost electronic configuration varies from ns2np1to ns2np6in each period. 
  • Group 18 (ns2np6) − noble gases
  • Group 17 − halogen
  • Group 16 − chalcogens
  • Non-metallic character increases from left to right across a period. 

• d- Block Elements (Transition Elements) 

  • Elements of group 3 to group 12
  • General electronic configuration is (n − 1) d1−10 ns0-2 
  • Called transition elements
  • Zn, Cd, and Hg with (n − 1) d10 ns2 configuration do not show properties of transition elements.
  • All are metals. They form coloured ions, exhibit variable oxidation states, paramagnetism, and are used as catalysts.

•  f- Block Elements 

  • Lanthanoids → Ce (Z = 58) to Lu (Z = 71)
  • Actinoids → Th (Z = 90) to Lr (Z = 103)
  • Outer electronic configuration → (n − 2) f1−14 (n −1) d0−1 ns2
  • They are called inner-transition elements
  • All are metals.
  • Actinoid elements are radioactive.
  • Elements after uranium are called Transuranium elements.

Metals, Non-metals, and Metalloids

  • Metals → Appear on the left side of the periodic table 
  • Non-metals → Located at the top right-hand side of the periodic table 
  • Elements change from metallic to non-metallic from left to right. 
  • Elements such as Si, Ge, As, Sb, Te show the characteristic properties of both metals and non-metals. They are called semi-metals or metalloids.

Periodic Trends in Physical Properties 

Atomic Radius 

  • Atomic radii decrease with the increase in the atomic number in a period. 
  • For example, atomic radii decrease from Li to F in the second period.

  • Nuclear charge increases progressively by one unit on moving from left to right across the period. As a result, the electron cloud is pulled closer to the nucleus by the increased effective nuclear charge, which causes decrease in atomic size. 
  • Atomic radii increase from top to bottom within a group of the periodic table. 
  • Variation of atomic radii with atomic number among alkali metals and halogen:
Ionic Radius
  • Cation is smaller than its parent atom. 
  • The size of the anion is larger than its parent atom.

Ionization Enthalpy

Defined as the amount of energy required to remove the most loosely bound electron from the isolated gaseous atom in its ground state


  • Decreases with the increase in atomic size 
  • Increases with the increase in nuclear charge 
  • Decreases with the increase in the number of inner electrons 
  • Increases with the increase in penetration power of electrons 
  • Atom having a more stable configuration has high value of enthalpy. 
  • Variation across a period: Increases with the increase in atomic number across the period. 
  • Variation in a group: Decreases regularly with the increase in atomic number within a group.

Electron Gain Enthalpy

  • Defined as the enthalpy change taking place when an isolated gaseous atom accepts an electron to form a monovalent gaseous anion.
  • Larger the value of electron gain enthalpy, greater is the tendency of an atom to accept electron. 
  • Greater the magnitude of nuclear charge, larger will be the negative value of electron gain enthalpy. 
  • Larger the size of the atom, smaller will be the negative value of electron gain enthalpy. 
  • More stable the electronic configuration of the atom, more positive will be the value of its electron gain enthalpy. 
  • Variation across a period − Tends to become more negative as we go from left to right across a period 
  • Variation down a group − Becomes less negative on going down the group 

Electronegativity

  • Defined as the tendency of an atom in a molecule to attract the shared pair of electrons towards itself 
  • Greater the effective nuclear charge, greater is the electronegativity. •
  • Smaller the atomic radius, greater is the electronegativity. 
  • In a period − Increases on moving from left to right 
  • In a group − Decreases on moving down a group

Valency

  • It is defined as the number of univalent atoms which can combine with an atom of the given element.
  • Valency is given by the number of electrons in outermost shell.
  • If the number of valence electrons ≤4: valency = number of valence electrons • If the number of valence electrons >4: valency = (8 - number of valence electrons)
  • In a period − Increases from 1 to 4 and then decreases from 4 to zero on moving from left to right
  • In a group − No change in the valency of elements on moving down a group. All elements belonging to a particular group exhibit same valency.

Non −Metallic (and Metallic Character) of an Element

  • Non-metallic elements have strong tendency to gain electrons. 
  • Non-metallic character is directly related to electronegativity and metallic character is inversely related to electronegativity. 
  • Across a period, electronegativity increases. Hence, non-metallic character increases (and metallic character decreases). 
  • Down a group, electronegativity decreases. Hence, non-metallic character decreases (and metallic character increases).
  • The periodic trends of various properties of elements in the periodic table are shown in figure. 

Atomic Number and Mass Number 

  • In the 1830s, representation of elements and compounds was a major concern for chemists. 
  • Many symbolic notations for elements were devised during this period. Gradually, the representations became standardized. Currently, the general symbolic notation for an element is: 
  • Now, take for example the specific symbolic notations for oxygen and nitrogen.
  • You know that the symbolic notation of oxygen is. In this notation, the letter ‘O’ symbolises the element ‘oxygen’; the number ‘16’ represents the mass number of oxygen; and the number ‘8’ indicates the atomic number of oxygen. 
  • Thus, in the general symbolic notation of an element, the letter ‘E’ is the symbol of the element, the letter ‘A’ is its mass number, and the letter ‘Z’ is its atomic number. 
  • The atomic number is the number of protons present in the nucleus of an atom. It is denoted by Z. 
  • The total number of the protons and the neutrons present in the nucleus of an atom is known as mass number. It is denoted by A. 

Symbolic Notations of Some Elements


Symbolic Notations of Some Elements


Relation between Atomic number and Mass Number 

Mass number (A) of an atom = Number of protons + Number of neutrons 

Therefore, Mass number (A) = Atomic number (Z) + Number of neutrons 

Therefore, Number of neutrons = A - Z

Hence, the number of neutrons can be calculated if the atomic number and mass number of an element are known. 

An atom of sodium contains 11 protons and 12 neutrons. Can you calculate the mass  number of a sodium atom? 

Now, mass number (A) = number of protons + number of neutrons Therefore, mass number of sodium atom = 11 + 12 = 23 

Hence, the mass number of sodium is 23.

An atom of carbon is represented as. Can you tell the number of neutrons and protons present in carbon atom? 

It is seen from the symbolic notation of carbon that the atomic number and mass number of carbon atom is 6 and 12 respectively. 

Now, number of neutrons = mass number − atomic number = 12 − 6 = 6 

Since the number of protons is equal to the atomic number of that element. Thus, the number of protons present in a carbon atom is 6. 

Solved Examples

Example 1: What is the symbol of the element sodium?

1. Na

2. N

3. So

4. S

Answer

The correct answer is A. 

The symbol of sodium is Na. It is derived from the Latin name for the element, i.e., ‘natrium’. 

Example 2: What is the atomic number of an element having five protons and six neutrons?

1. 11

2. 9

3. 6

4. 5

Answer

The correct answer is D.

The atomic number of an element is the number of protons or electrons present in an atom of the element. Since an atom of the given element has five protons, its atomic number is 5. 

Example 3: What is the number of neutrons in an element having 39 protons and 89 as its  mass number? 

1. 45

2. 50

3. 55

4. 60

Answer

The correct answer is B. 

We know that: 

Mass number = Number of protons + Number of neutrons 

In case of the given element: 

Mass number = 89 

Number of protons = 39 

So, 

89 = 39 + Number of neutrons 

⇒ Number of neutrons = 89 -39 = 50 

Example 4: What is the symbol of the element having 22 neutrons and 40 as its mass  number? 

1. Al

2. Mg

3. Ar

4. Ca

Answer

The correct answer is C. 

The given element has: 

Mass number = 40 

Number of neutrons = 22 

We know that: 

Mass number = Number of protons + Number of neutrons 

So, 

40 = Number of protons + 22 

⇒ Number of protons = 40 - 22 = 18 

Also, 

Atomic number = Number of protons = 18 

Argon is the element having 18 as its atomic number and 40 as its mass number. The symbol of argon is Ar. 

Did You Know?

Water is the major constituent of the human body. It is made up of two elements: hydrogen and oxygen.

Almost all the mass of our body is made up of the following six elements. 

  1. Oxygen (65%)
  2. Carbon (18%)
  3. Hydrogen (10%)
  4. Nitrogen (3%)
  5. Calcium (1.5%)
  6. Phosphorus (1%) 

Some of the other elements found in our body are:Sulphur (0.25%)

Sodium (0.15%)

Magnesium (0.05%)

Zinc (0.7%)


The Periodic Table 

  • The periodic table is a table classifying all the known elements. 
  • It is divided into 18 columns (called groups) and 7 rows (called periods). 
  • The elements are arranged in the rows or periods by order of increasing atomic number. 
  • The elements in the columns or groups display similar chemical and physical properties. This feature of the periodic table makes it easy to study the vast number of elements. 

The periodic table is shown in the figure.



Comparison of Alkali Metals and Halogens

Parameter

Alkali Metal

Halogens

Element

Lithium (Li)

Sodium (Na)

Potassium (K)

Rubidium (Rb)

Cesium (Cs)

Francium (Fr)

Fluorine (F)

Chlorine (Cl)

Bromine (Br)

Iodine (I)

Astatine (At)

Occurrence

Combined state

Combined state

Physical State

Metal

Silvery white

Soft and light

Non-metal

Coloured

F and Cl are gases

Br is liquid

I is solid

Valence Electrons

Valence shell contains one electron

Valence shell contains seven electrons

Conductivity

Good conductor of electricity

Non-conductor of electricity

Melting and Boiling Point

Decreases down the group

Increases down the group

Atomic Size

Largest (except inert gases) in their respective period increases down the group

Smallest in their respective period

Increases down the group

Ionisation Energy

Lowest in the respective period

Decreases on moving down the group

Highest (lower than noble gases) in the respective period

Decreases down the group due to increase in atomic size

Electron Affinity

Low

Decreases on moving down the group

Low

Decreases on moving down the group

Electronegativity

Lowest in respective period

Decreases on moving down the group

Highest in respective period

Decreases on moving down the group

Reactivity

Highly reactive because of large size

Low ionization enthalpy

Reactivity increases down the group

Halogens are highly reactive.

They react with metals and non-metals to form halides.

Reactivity decreases down the group.

Reaction with water and acid

Vigorous  

Liberate hydrogen reactivity decreases down the group

Generally they do not react

Reducing/Oxidising Nature

Strong reducing agent

Strong oxidising agent

Formation of Compounds

Form electrovalent compounds with non-metals

Form electrovalent compounds with metals


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