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Classification of Elements and Periodicity in Properties Chapter - Chemistry JEE Main

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Concepts of Classification of Elements and Periodicity in Properties for JEE Main Chemistry

The periodic table is a tabular arrangement of chemical elements that is arranged by electron configurations, atomic number (number of protons in the nucleus), and recurrent chemical properties. s-block, p-block, d-block, and f-block are four rectangular blocks shown in this table. Metals fall on the left-hand side of one row or period whereas non-metals fall on the right-hand side.

This chapter covers all of the trends in the physical and chemical properties of the elements in detail. Electron affinity, Electropositivity, Electronegativity, Aufbau Principle, and Ionisation Energy are some of the significant topics covered in this chapter. This is a crucial chapter for students who want to learn about chemistry and how to identify the properties of elements. The periodic properties have a low exam weighting, although it is an important chapter in the chemical topic.


JEE Main Chemistry Chapters 2024 


Important Topics of Classification of Elements and Periodicity in Properties Chapter

  • Modern Periodic Table

  • Periodic Properties of Elements


Classification of Elements and Periodicity in Properties Important Concept for JEE Main

Periodic Table

  • The periodic table is a grouping of elements that have similar properties.

  • It's a representation of the periodic law in graphic form.
    Periodic Law: It claims that the properties of chemical elements are related to their atomic numbers in a predictable manner.

  • The table is divided into four blocks, each of which is about rectangular in shape. 

  • The table's rows are known as periods, and the columns are known as groups. 

  • Chemical characteristics of elements in the same column of the periodic table are comparable.

  • Nonmetallic character (holding their own electrons) increases from left to right throughout a period and from down to up across a group.

  • While the metallic character (surrendering electrons to other atoms) increases in the opposite way across the periodic table.


Groups

  • Groups refer to the 18 vertical columns.

  • A family is made up of elements that belong to the same group and is usually titled after the initial number.


Periods

  • Periods refer to the horizontal rows. 

  • In the long version of the periodic table, there are seven periods.


Periodic Properties of Elements

Shielding or Screening Effect

  • In a multi-electron atom, the nucleus attracts valence electrons whereas inner-shell electrons repel them.

  • The valence-electron perceives less attraction from the nucleus as a result of the combined effect of these attractive and repulsive forces acting on it. The effect is known as shielding or screening.


Effective Nuclear Charge

  • Denoted by z*.

  • The following formula relates it to actual nuclear charge (Z):
    Z* = (Z - 𝜎), where 𝜎 is the given screening constant.

  • When we proceed from left to right, we notice that the size of effective nuclear charge grows.


Variation of Ionisation Energy in Periodic Table

  • As the atomic number of a group grows, the ionisation energy drops. With an increase in atomic number, the ionisation energy increases over time.


Electron Affinity and Electronegativity 

  • Generally, the definition of Electron affinity is given as "the energy produced when an additional electron is added to a neutral gaseous atom." 

  • The electron affinity increases from left to right in the period. 

  • The ability of an atom in a compound to attract a pair of bound electrons to itself is referred to as electronegativity.

  • Electronegativity grows from left to right over time. Because of the decrease in size and increase in nuclear charge, this is the case.

  • As a result, alkali metals have the lowest value and halogens have the highest.

  • The electronegativity of inert gases is 0.

  • Electronegativity falls from top to bottom in a group. This is due to an increase in the size of the atoms.


Mulliken's Scale

  • The average value of an atom's ionisation potential and electron affinity was defined by Mulliken as electronegativity.

  • Electronegativity = (Ionisation Potential + Electron Affinity)/2.


Pauling Scale

  • The Pauling electronegativity scale is the most extensively used. It is predicated on the presence of surplus bond energy.

  • ΔE = Actual Bond Energy = -√(EA-A X EB-B).


Atomic Volume

  • It is the volume occupied by one gramme of an element's atom.

  • Atomic Volume = (Gram Atomic Weight)/(Density in Solid State).

  • c.c./mole is the unit of atomic volume.


Density

  • The density of elements in solid form fluctuates with their atomic numbers on a regular basis.

  • Initially, the density increases gradually over time and reaches a maximum someplace for the centre members, before progressively diminishing.


Melting and Boiling Points

  • The melting temperatures of the elements show considerable regularity as the atomic number increases.


Oxidation State (Oxidation Number, O.N.)

  • The total number of electrons that an element in a compound appears to have gained or lost (negative and positive oxidation states, respectively) during the creation of that compound is known as its oxidation number.


Magnetic Properties

  • The characteristics of individual atoms determine the magnetic properties of matter.


Modern Periodic Law and the Present form of the Periodic Table

The periodic table, a cornerstone of chemistry, has evolved over time with the development of the modern periodic law. This law states that the physical and chemical properties of elements are periodic functions of their atomic numbers, which led to the arrangement of elements in a manner that reflects these periodic trends.


The present form of the periodic table consists of rows (periods) and columns (groups). Elements within the same group share similar chemical properties, while elements within the same period have sequentially increasing atomic numbers. The periodic table is divided into several blocks, including the s, p, d, and f block elements, each having unique characteristics and properties.


S, P, D, and F Block Elements

S Block Elements:

The s-block elements are found in groups 1 and 2 of the periodic table. They include the alkali metals (Group 1) and the alkaline earth metals (Group 2). These elements have one or two valence electrons, making them highly reactive. They are known for their strong reducing properties and the ability to form ionic compounds.


P Block Elements:

The p-block elements are found in groups 13 to 18. These elements include metals, nonmetals, and metalloids. They have a varying number of valence electrons, and their properties range from metallic to non-metallic. Elements in Group 17 are known as the halogens, which are highly reactive nonmetals, while those in Group 18 are noble gases with low reactivity.


D Block Elements:

The d-block elements are also known as the transition elements. They are found in groups 3 to 12 and are characterized by the presence of partially filled d-orbitals. These elements exhibit a wide range of properties and are known for their variable oxidation states and the formation of colorful complexes. They play a vital role in catalysis.


F Block Elements:

The f-block elements are located below the main body of the periodic table and consist of the lanthanides and actinides. Lanthanides are elements with atomic numbers from 57 to 71, while actinides have atomic numbers from 89 to 103. They are known for their radioactive nature and are primarily synthetic. They are important in nuclear reactions and in the development of nuclear energy.


Periodic Trends in Properties of Elements

Atomic and Ionic Radii:

Atomic radius is defined as the distance between the nucleus and the outermost electron in an atom. As you move across a period from left to right, atomic size decreases due to increased nuclear charge. When moving down a group, atomic size increases because of the addition of energy levels. Ionic radii follow a similar trend; cations (positively charged ions) are smaller than their parent atoms, while anions (negatively charged ions) are larger.


Ionization Enthalpy:

Ionization enthalpy is the energy required to remove an electron from an atom. It increases across a period due to the increasing nuclear charge, making it more difficult to remove electrons. Down a group, ionization enthalpy decreases because of the increased atomic size and the shielding effect of inner electrons.


Electron Gain Enthalpy:

Electron gain enthalpy is the energy released when an electron is added to an atom. It becomes more negative across a period as atoms strive to attain a stable electronic configuration. Down a group, electron gain enthalpy becomes less negative due to increased atomic size.


Valence, Oxidation States, and Chemical Reactivity:

Valence electrons are the electrons present in the outermost energy level of an atom. The number of valence electrons determines an element's chemical reactivity. Elements with a full valence shell (Group 18, noble gases) are chemically inert, while those with a few or one valence electron are highly reactive. The periodic table can be used to predict the possible oxidation states (charges) of elements. For example, Group 1 elements tend to form +1 ions, while Group 17 elements often form -1 ions. These oxidation states are essential for understanding chemical reactions and the formation of compounds.


JEE Main Classification of Elements and Periodicity in Properties Solved Examples

Example 1: The third period of the p-block contains an element. Its outermost shell has 5 electrons. Predict the composition of the group. What is the number of unpaired electrons in it?

Solution: It is a member of the 15th group (P). It has three electrons that are not coupled.


Example 2: Recently, an element X with Z = 112 was found. Predict its electrical configuration and recommend which group it belongs to.

Solution: Rn 5f14 6d10 7s2 is the element. It is a member of the 12th group.


Solved Questions From the Previous Year Question Papers

Question 1: The first ionisation potential of Na is 5.1 eV. The value of electron gain enthalpy of Na+ will be?

(1) – 2.55 eV

(2) – 5.1 eV

(3) – 10.2 eV

(4) + 2.55 eV

Solution: 

  • The minimal amount of energy necessary to remove an electron from an atom or molecule in the gaseous state is described by the ionisation energy of an atom or molecule. 

  • However, the amount of energy released when an isolated gaseous atom accepts an electron to become a monovalent gaseous anion is known as electron gain enthalpy. H = -5.1 eV in this case.

  • As a result, option (2) is the proper response.


Question 2: The order of increasing sizes of atomic radii among the elements O, S, Se, and As is?

(1) As < S < O < Se

(2) O < S < As < Se

(3) Se < S < As < O

(4) O < S < Se < As

Solution: 

  • The size of atomic radii grows as you move from top to bottom. 

  • The size reduces as you move from left to right. As a result, O < S < Se < As.

  • Therefore, option (4) is the proper response.


Question 3: The radius of La3+ (atomic number of La = 57) is 1.06 A. Which one of the following given values will be closest to the radius of Lu3+ (atomic number of Lu = 71)?

(1) 40 A

(2) 1.06 A

(3) 0.85 A

(4) 1.60 A

Solution: 

  • ∴ r = 1/Z is the atomic radius. 

  • ∴ Z2/Z1 = r1/r2, i.e., the ratio of two atomic radii.

  • ∴ 71/57 = 1.06/r2

  • ∴ As a result, r2 = 1.0657/71 = 0.85.

  • So, option (3) is the correct answer.


Practice Questions

Question 1: The screening effect of electrons in the s, p, d, and f orbitals of a given shell of an atom on the electrons in its outer shell is in which order:

(i) s > p > d > f

(ii) d > p > s > f

(iii) p < d > s > f

(iv) f < p > s > d

Answer: (i) s > p > d > f


Question 2: Which of the following is the right order of the following species' sizes:

(i) I > I+ > I

(ii) I+ > I > I

(iii) I > I > I+

(iv) I > I > I+

Answer: (iv) I > I > I+


Question 3: Which of the following alternatives has an order of arrangement that does not match the property variation noted against it?

(i) Al3+ < Mg2+ < F > Na+ (increasing ionic size)

(ii) Li > Na < Rb > K (increasing metallic radius)

(iii) I < Br < Cl < F (increasing electron gain enthalpy)

(iv) B < C < N < O (increasing first ionisation enthalpy)

Answer: (iii) I < Br < Cl < F (increasing electron gain enthalpy), and

(iv) Li < Na < K < Rb (increasing metallic radius) and


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Conclusion

In this article, we dive into the essential topic of "Classification of Elements and Periodicity in Properties" for JEE Main. We provide comprehensive insights into the key concepts and effective problem-solving approaches. You'll discover how elements are categorized and understand the patterns in their properties, crucial for acing your exams. We've gathered definitions, questions, and solutions, all in one place. Plus, our PDFs are readily available for free download, making learning convenient. This resource is a valuable asset for students preparing for their exams, helping them grasp the chapter's intricacies and boost their confidence.

FAQs on Classification of Elements and Periodicity in Properties Chapter - Chemistry JEE Main

1. What is periodicity in classification of elements?

Periodic categorization of elements is a way of grouping elements based on their characteristics, such as keeping elements that are similar in one group and the rest of the elements in the other.

2. What are the elements' classifications? What are their distinguishing features?

Chemical and physical features can be used to categorise elements and their compounds. Acids vs. bases, metals vs. nonmetals, and other major classes are probably already recognisable to you.

3. What is the purpose of element classification?

It is vital to classify elements in order to compare their physical and chemical properties and to group similar elements together. There are roughly 114 elements known, and we must investigate each one.