NCERT Class 11 Chemistry Chapter 3 Notes - Download PDF

The classification of elements and periodicity in properties is an important chapter in the class 11 Chemistry curriculum, helping us to understand why different elements behave differently. Almost every element around us is linked to the periodic table whether it is salt in our food or iron in large buildings. Periodicity in properties helps us to explain how electronic configuration and periodic trends form the basis of the arrangement of elements in the periodic table. The periodic table is divided into s-block, P- Block, D- And F- Block.

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This Story also Contains

1. NCERT Class 11 Chemistry Chapter 3 Notes
2. Genesis of Periodic Classification
3. Modern Periodic Table
4. IUPAC Nomenclature of Elements with Atomic Number More Than 100
5. Elements in s, p, d, f- Block And Their Electronic Configurations
6. Trends in Physical Properties
7. Trends in Chemical Properties
8. Anomalous Properties of Second Group Elements
9. NCERT Class 11 Notes Chapter-Wise
10. Subject Wise NCERT Exemplar Solutions
11. Subject Wise NCERT Solutions

NCERT class 11 Chemistry notes of Classification of Elements and Periodicity in Properties are designed by our subject experts to offer a systematic and structured approach to these important concepts and help students develop a clear understanding of critical concepts. Classification of elements and periodicity in properties is very important for Class 11 since it explains basics like chemical bonding and periodicity. CBSE Class 11 Chemistry chapter 3 notes go over the various classifications, such as d-block, p-block, s-block, and f-block. Class 11 Classification of Elements and Periodicity in Properties notes contains detailed explanations of each topic with appropriate examples. Class 11 Chemistry notes can help students score good marks in their examinations.

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NCERT Class 11 Chemistry Chapter 3 Notes

Genesis of Periodic Classification

Many elements were found in the nineteenth century, and studying each one separately proved difficult. Many scientists have attempted to classify elements in a variety of ways.

• Dobereiner triads

Johann Wolfgang Dobereiner, a German chemist, attempted to arrange elements with comparable properties into three groups. Triads were the name given to these groups. The atomic mass of the middle element in these triads, according to Dobereiner, should be more or less equal to the mean of the atomic masses of the other two elements in the triad. A trio including lithium, sodium, and potassium is an example of such a triad. Lithium has an atomic mass of 6.94, while potassium has a mass of 39.10. The centre element in this triad, sodium, has an atomic mass of 22.99, which is about similar to the mean of lithium and potassium's atomic masses (which is 23.02).

Limitation

There was no way to classify all of the elements known at the time into triads.

• Newlands law of octaves

In the year 1866, English chemist John Newlands organised the 56 known elements in ascending order of atomic mass. He noticed a pattern in which every eighth element had properties that were comparable to the first.

According to Newland's Law of Octaves, when elements are ordered in ascending order of atomic mass, the periodicity in properties of two elements separated by seven elements will be comparable.

Limitations

The classification of elements using Newland's Octaves was only used up to calcium.

The discovery of noble gases added to the method's limitations because they couldn't be included to the arrangement without completely disrupting it.

• Mendeleev’s periodic table

Dmitri Ivanovich Mendeleev, a Russian chemist, proposed the periodic table in 1869. He saw that the physical and chemical properties of elements were connected to their atomic masses periodically.

The chemical characteristics of elements are a periodic function of their atomic weights, according to the Periodic Law (also known as Mendeleev's Law).

Advantages of Mendeleev’s periodic table

He classified the elements according to their atomic masses.

Undiscovered elements such as Gallium and Scandium germanium were left in the gaps. Even he left the entire group empty in case of undiscovered inert gases.

He was able to anticipate the proportions of certain elements based on their positions in the periodic table, such as Ga and Sc.

He was able to accurately predict inaccuracies in the atomic weights of several elements, such as gold and platinum.

Limitations

Isotope positions could not be explained

The wrong order of atomic masses could not be justified. For example, while Argon has an atomic mass of 40, and K has a low atomic mass of 30, although K should be placed first due to his low atomic mass.

The position of hydrogen was unable to be explained.

Some aspects that are dissimilar are grouped together, while others are divided into different categories. Alkali metals such as Li, Na, K, and others (group AI) are grouped alongside coinage metals such as Cu, Ag, and Au, for example.

The main body excludes lanthanides and actinides.

Modern Periodic Table

Henry Moseley, an English physicist, researched the wavelength of the characteristic X-rays in 1913. Using various metals as anti-cathodes, it was demonstrated that the square root of the line's frequency is proportional to the atomic number. Moseley developed the modern periodic law based on the above data, which states that “Physical and chemical properties of the elements are the periodic function of their atomic numbers”.

As a result, there is periodicity in electronic configuration when elements are organised according to increasing atomic numbers, which leads to periodicity in chemical properties. Bohr's Periodic Table is the long form of the Periodic Table. This Periodic Table has 18 groups and seven periods. Periods refer to the horizontal rows. Groups are made up of the 18 vertical columns.

IUPAC Nomenclature of Elements with Atomic Number More Than 100

Elements in s, p, d, f- Block And Their Electronic Configurations

s-block

• The last electron enters the s-orbitals in these elements.
• ns^1-2, where n = 2-7, is a general outer shell electrical configuration of s-block elements.
• They are metals that are soft and have low melting and boiling points
• They are electro positive and have low ionization enthalpies (energies).
• They rapidly lose their valence (outermost) electrons, forming +1 (alkali metals) and +2 ions (in case of alkaline earth metals).
• They are metals with a lot of reactivity. As we progress through the group, the metallic character and reactivity become more prominent. They are never found pure in nature due to their high reactivity.
• With the exception of beryllium, all s-block element compounds are generally ionic.

p-block

• General electronic configuration ns²np^1-6
• d-block element’s compounds are primarily covalent in nature
• They have a wide range of oxidation states.
• The non-metallic character of the components grows as the period progresses from left to right.
• The reactivity of elements within a group tends to decrease.
• A noble gas element with a closed valence shell ns²np⁶configuration is present at the end of each period.
• As we move through the group, the metallic character becomes more prominent.

d-block

• General electronic configuration is (n-1)d^1-10ns^1-2
• They are all high-melting and boiling-point metals.
• d-block element’s compounds are usually paramagnetic in nature.
• They mostly produce coloured ions and have a varied valence (oxidation states).
• Because they have incompletely filled d-orbitals in their ground state or any of the oxidation states, the d-block elements are known as transition elements
• Compounds of d-block elements are frequently used as catalysts

f-block

• General electronic configuration is n-2f^1-14((n−2)f⁽⁰⁻¹⁴⁾(n−1)d⁽⁰⁻¹⁾ns².
• They are called inner transition elements because in d-block transition elements, electrons are filled in (n – 1) d subshells, whereas in f-block inner transition elements, electrons are filled in (n – 2) f subshells, which is one inner subshell.
• The Lanthanoids Ce (Z = 58) – Lu (Z = 71) and Actinoids Th (Z = 90) – Lr (Z = 103), two rows of elements near the bottom of the Periodic Table.
• They are all made of metal. The characteristics of the elements in each series are relatively similar.
• The actinoid series contains a large number of radioactive elements.

Trends in Physical Properties

• Atomic radii

The atomic radius is the distance between the nucleus's centre and the outermost shell containing electrons. These are classified as covalent radii, van der Waals radii and metallic radii according on the nature of bonding. Covalent radii are one-half of the distance between the centres of the nuclei of two neighbouring identical atoms linked by a single covalent bond. For example, the distance between Cl and Cl is 198 pm, and the covalent radius of Cl is 99 pm. Van der Waals radii are half of the internuclear distance between two comparable nearby atoms belonging to two adjoining molecules of the same material in the solid state. In a metallic crystal, a metallic radius is defined as half the distance between the nuclei of two adjacent atoms. Due to an increase in effective nuclear charge, the atomic radius decreases as we move from left to right in a period. The atomic radius increases as we move down the group due to an increase in the main quantum number, which leads to an increase in the number of shells and a higher shielding effect.

• Ionic radii

When one or more electrons are lost from a neutral atom, the ions created are known as cation (positive ion). When electrons are added to the neutral atom, the ions are known as anion (negative ion) . The ionic radii are the effective distances from the nucleus's centre to which the ion exerts its impact on the electron cloud. Ionic radii follow the same pattern as atomic radii. It reduces from left to right as the period progresses and increases from top to bottom as the group goes. Any natural atom's cation and anion sizes are as follows: cation< neutral atom <anion

• Ionization enthalpy

Ionization Enthalpy is the amount of energy necessary to remove an electron from an isolated gaseous atom's outermost orbit. In general, IE grows as moves from left to right in period, whereas it drops as moves down the group, however half-filled and fully filled orbitals are highly stable and hence have a high IE.

• Electron gain enthalpy

The change in enthalpy that occurs when a gaseous atom gains an extra electron to form a monovalent anion in the gaseous state is known as the electron gain enthalpy. The enthalpy of electron gain increases across the periods but drops as the group progresses.

• Electronegativity

Electronegativity is an atom's tendency to attract the shared pair of electrons in a covalent bond towards itself. The most electronegative element is fluorine, while the least electronegative one is cesium. Electronegativity increases in the periods from left to right. The electronegativity of elements in the groups reduces as go down.

Trends in Chemical Properties

• Valency

It is defined as the element's ability to combine. The electrical arrangement of an element's atom determines its valency, which is usually dictated by electrons in the valence shell. Valency increases from 1 to 4 and eventually declines to 0 (for noble gases) as one moves along a period from left to right, but valency remains constant as one moves along a group. Because they can use electrons from both the outer and penultimate shells, transition metals have variable valency.

• Chemical reactivity

Metal reactivity increases as the IE, electronegativity, and atomic size rise, as well as the metal's electropositive character.Increases in electronegativity, as well as electron gain enthalpy and atomic radii, increase non-metal reactivity.

Anomalous Properties of Second Group Elements

Certain 2nd-period elements have properties that are similar to their 3rd period diagonal elements. As a result, Li is similar to Mg, Be is similar to Al, and B is similar to Si. This is known as a diagonal relationship, and it occurs because both elements have almost identical ionic radii and polarizing power (charge/size ratio). Mg, Al and Si are known as bridge elements.

Significance of NCERT Class 11 Chemistry Chapter 3 Notes

Classification of elements and periodicity in Properties Chemistry class 11 notes discuss the classification of elements and the periodic table. This section is really important in the exam point of view. Students require high-quality study material to crack exams like NEET, JEE MAIN and board exams. NCERT Class 11 Chemistry chapter 3 contain detailed information that help to crack these exams. CBSE Class 11 chemistry ch 3 notes cover the main topics of the Class 12 CBSE Chemistry Syllabus and are designed based on latest exam pattern. NCERT notes for Class 11 Chemistry chapter 3 is very simple for the students to understand. Students can use chemistry class 11 chapter 3 notes pdf to download to study offline.

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