Classification of Elements and Periodicity in Properties Class 11

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classification of elements and periodicity in properties

In the third chapter of Class 11 Chemistry, students learn the classification of elements and periodicity in properties. The periodic table is one of the most fundamental topics among all Class 11 Chemistry chapters. The first half of this unit introduces the history of the periodic table and the modern periodic law. In the second half of the chapter, students learn periodicity in the physical and chemical properties of the elements. Checkout these Classification of Elements and Periodicity in Properties study notes to know more about the topic!

Why was there a need for the classification of elements? 

As a large number of elements were discovered, it became challenging to study each element individually. Hence, there was a need for classification to make the study of elements easier.

Historical Development of the Periodic Table – Early Attempts

There were several attempts to classify elements in the 18th and 19th centuries. While the history of the development in Classification of Elements and Periodicity in Properties is quite detailed, in our notes, we highlight some of the famous early classification attempts:

Prout’s Hypothesis (1815) – It is also known as unitary theory. It is based on Hydrogen as the fundamental unit from which all other atoms are created.

Dobereiner’s Tridas (1829) – In this method, the central element’s atomic weight is the arithmetic mean of the other two. Some examples of triads are Cl, Br, and I; Li, Na, and K; Ca, Sr, and Ba.

Newland’s Octaves (1864) – According to this theory, when elements are arranged in increasing atomic masses, the eighth element’s property is the same as the first one, just like in the musical notes.

Lother Meyer’s Atomic Volume Curve (1869) – Meyer arranged elements in the form of a curve and stated that an element’s properties are based on its atomic volume.

Limitations of Early Classification Laws: These laws couldn’t account for new elements. Hence, they became irrelevant as new elements were discovered. 

Mendeleev’s Periodic Table – It is based on the law, “Physical and chemical properties of an element is a periodic function of its atomic mass.” Only 63 elements were known when Mendeleev designed this periodic table. 

Key Features of Mendeleev’s Periodic Table:

  • It was divided into seven rows (periods) and eight columns (groups).
  • It helped in the systematic study of elements.
  • He left spaces for new elements that would be discovered. 
  • Zero group elements were later added to the modified Mendeleev’s periodic table. 

Limitations of Mendeleev’s Periodic Table 

  • The position of Hydrogen is disputed
  • Position of isotopes 
  • Anomalous positions of certain elements 
  • Lanthanoids and Actinoids were not placed in the main table 

Must Read: Class 11 Chemistry Syllabus

The Modern Periodic Table (1913)  – Moseley introduced the modern periodic table, which was a modified version of Mendeleev’s Periodic Law. Moseley stated that “physical and chemical properties of an element is a periodic function of its atomic number.” The basis for the modern form of the periodic table is that when elements are arranged in increasing orders of their atomic numbers, it was observed that specific properties repeated at certain regular intervals. The intervals at which the properties are repeated are 0, 1, 2, 8, 18, and 31. These numbers are referred to as magic numbers.

Key Features of the Long Form of Periodic Table 

  • The long form of the modern periodic table is known as Bohr’s Periodic Table. It has 18 groups (vertical columns) and 7 periods (horizontal rows). 
  • Periods 
    • The first period is the shortest and contains just 2 elements – 1H and 2He
    • The second and third period contains 8 elements each and is short. The second period is from 3Li to 10Ne, and the third period is from 11Na to 18Ar
    • The fourth period (19K to 36Kr) and the fifth period (37Rb to 54Xe) are long periods and contain 18 elements.
    • The sixth period has 32 elements (55Cs to 86Rn) and is the longest period. 
    • The seventh period starts from 87Fr and is incomplete. It has 19 elements. 
  • Groups 
    • 1st group – alkali metal 
    • 2nd group – alkaline earth metals
    • 16th group – ore-forming elements known as chalcogens 
    • 17th group – sea-salt forming elements known as halogens 
    • 18th group – noble gases 
  • Anomalous behaviour of the 1st element in a group – The property of the first element in a group differs considerably from the rest of the elements in the group. This is due to several factors:
    • High electronegativity 
    • Small size
    • Non-availability of d-orbitals for bonding 
  • The periodic table is divided into four main blocks – s, p, d, and n – depending on the sub-shell into which the valence electron enters. s and p-block elements are jointly known as representative elements
    • Electronic configuration of s-block elements: ns1-2
    • Electronic configuration of p-block elements: ns2np1-6
    • Electronic configuration of d-block elements: ns1-2(n-1)d1-10
    • Electronic configuration of n-block elements: (n – 2)f1-14(n-1)d0-1ns2

Limitations of Long Form of the Periodic Table

  • The position of Hydrogen is still uncertain.
  • The inner-transition elements are not placed in the main body of the table but are placed separately.

Recommended Read: Tricks to Memorise the Periodic Table

IUPAC Nomenclature of Elements for Elements with Z>100

Whenever a new element is discovered, the discoverer names it according to IUPAC (International Union of Pure and Applied Chemistry) nomenclature. The names are derived directly from atomic numbers and involve a mixture of Greek and Latin roots, to avoid duplicity. The name is formed by adding the suffix –ium. 

Roots for IUPAC Nomenclature 

Digit 0 1 2 3 4 5 6 7 8 9
Root Nil Un Bi Tri Quad Pent Hex Sept Oct Enn
Symbol n u b t q p h s o

To give an example, an element with atomic number 123 is written as unbitriium and has the symbol ubt. 

What are Periodic Properties? 

The properties of an element that are directly or indirectly related to their electronic configuration and show gradual change as we move from top to bottom in a group or left to right in a period are known as periodic properties. Periodicity occurs due to the repetition of similar electronic configuration at specific intervals. Let’s take a look at critical periodic properties:

Atomic Radius

It is the distance from the centre of the nucleus in an atom to the outermost shell of electrons. The atomic radius decreases on moving from left to right in a period and increases on moving from top to bottom in a group. 

  • Covalent Radius – The half distance between the nuclei of two identical atoms joined together in a covalent bond is known covalent radius. 
  • Van der Waals’ Radius – It is the one-half distance between the nuclei of two adjacent atoms in the solid state. 
  • Metallic Radius – It is defined as the one-half distance between the nuclei centres of two adjacent atoms that are in the metallic state. 
  • Ionic Radius – An atom changes to a cation by losing electrons and into an anion by gaining electrons. 

Valency 

It is the combining capacity of an element. An element’s valency is determined by the number of electrons present in the valent shell of an atom. The valence increases from 1 to 4 and drops to 0 (for noble gases) while moving from left to right along a period. The valency remains the same while moving from top to bottom in a group. Transition metals have variable valency as they can use electrons from the outer shell and the penultimate shell.  

Ionisation Enthalpy 

Ionisation Enthalpy (IE) is the energy required to remove a loosely bound electron from an isolated gaseous atom. The IE increases on moving left to right in the periodic table and decreases on moving down a group. 

Electron Gain Enthalpy (EGE) 

EGE is the amount of energy released when an electron is added to an isolated gaseous atom. The electron gain enthalpy becomes more and more negative while moving from left to right. It becomes less negative when moving from top to bottom in the periodic table.

Electronegativity 

An atom tends to attract a shared electron pair in a covalent bond. Factors that affect electronegativity are atomic size, charge on the ion, and hybridisation. The electronegativity increases on moving from left to right, and it decreases while moving down a group. The electronegativity value of a noble gas is zero.

Along a Period: 

  1. Metallic character decreases on moving from left to right. It is the maximum on the extreme left of a period. 
  2. Non-metallic character increases on moving from left to right. 
  3. Basic nature of oxides decreases along a period on moving from left to right. 
  4. Acidic nature of oxides increases along a period on moving from left to right. 

Down a Group: 

  1. Metallic character increases on moving down a group. 
  2. Non-metallic character decreases on moving down a group. 
  3. The basic nature of oxides increases moving down a group. 
  4. The acidic nature of oxides decreases moving down a group. 

Important Questions on Classification of Elements and Periodicity in Properties

  • What is the difference between the atomic radius and the covalent radius?
  • What are the limitations of the Mendeleev’s Periodic Table? 
  • Why do elements in the same group have similar physical and chemical properties? 
  • Explain the difference between Electron Gain Enthalpy and Electronegativity. 
  • What is the IUPAC name of an element with atomic number 118? 

Now that you’ve understood the classification of elements and periodicity in properties make sure to check other Class 11 Chemistry chapters on Leverage Edu.

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