NCERT Class 11 Chemistry Chapter 3 Notes - Download PDF

NCERT Notes for Class 11 Chemistry Chapter 3 Download PDF

These concise NCERT notes cover all the key concept of chapter 3 to help students in quick revision before exam. You can download NCERT notes pdf from the button given below:

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NCERT Notes for Class 11 Chapter 3

Chapter 3 Class 11 notes covers key concepts, including the layout of the modern periodic table, groups and periods, and periodic trends such as Atomic Size & Atomic Radius, Ionization Enthalpy, and Electronegativity. These notes explain basics like chemical bonding and periodicity. These notes is an excellent resource for quick revision, as it helps build a clear understanding of fundamental principles and their real-life applications.

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 center 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.

• Newland law of octaves

In the year 1866, English chemist John Newlands organized 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 in 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. They even 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, Argon has an atomic mass of 40, and K has a low atomic mass of 30, although K should be placed first due to its low atomic mass.

The position of hydrogen could 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 Law and the Present Form of the 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 Table 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 organized 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 18 vertical columns.

IUPAC Nomenclature of Elements with Atomic Number More Than 100

Electronic Configuration of Elements and the Periodic Table

It explains how the arrangement of electrons in shells or energy levels helps determine an element’s position in the periodic table. It also reveals trends in properties like atomic size and chemical reactivity, etc, across periods and groups.

Electronic Configurations in Periods

The period indicates the value of n for the outermost shell. In other words, successive periods in the Periodic Table are associated with the filling of the next higher principal energy level (n = 1, n = 2, etc.). As we move from left to right across a period, one electron is added to the same principal energy level, while the number of protons in the nucleus also increases. This leads to a change in properties such as atomic size, metallic character, and valency.

Groupwise Electronic Configurations

Elements in the same vertical group have similar valence shell electronic configurations, the same number of electrons in the outer orbitals, and similar chemical reactivity and physical properties. For example, the Group 1 elements (alkali metals) all have an ns¹ valence shell electronic configuration.

Elements in s, p, d, f-block and Their Electronic Configurations

The s-Block Elements

• 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 electropositive and have low ionization enthalpies (energies).
• They rapidly lose their valence (outermost) electrons, forming +1 (alkali metals) and +2 ions (in the 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.

The p-Block Elements

• General electronic configuration ns²np^1-6
• d-block elements’ 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.

The d-Block Elements (Transition Elements)

• The general electronic configuration is (n-1)d^1-10ns^1-2
• They are all high-melting and boiling-point metals.
• d-block elements’ compounds are usually paramagnetic in nature.
• They mostly produce colored ions and have 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

The f-Block Elements (Inner-Transition Elements)

• 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), are 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.

Metals, Non-metals and Metalloids

Based on their physical and chemical properties, elements in the periodic table are broadly classified as metals, non-metals, and metalloids.

Metals, mostly found on the left and centre of the periodic table. They are typically lustrous, good conductors of heat and electricity, malleable, and ductile. They tend to lose electrons to form positive ions.

Non-metals, mostly found on the right side of the periodic table. They are generally poor conductors, brittle in solid form, and tend to gain or share electrons, forming negative ions or covalent bonds.

Metalloids, silicon, germanium, arsenic, antimony and tellurium lie along the stair-step line (between metals and non-metals) and exhibit properties intermediate between those of metals and non-metals. Their classification helps in understanding trends in physical and chemical properties.

Periodic Trends in Properties of Elements

There are many observable patterns in the physical and chemical properties of elements as we descend in a group or move across a period in the Periodic. Based on the physical and chemical properties, trends are explained as below.

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 to 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 cations (positive ions). When electrons are added to the neutral atom, the ions are known as anions (negative ions). 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

Ionisation enthalpy is the amount of energy necessary to remove an electron from an isolated gaseous atom's outermost orbit. In general, IE grows as it moves from left to right in a period, whereas it drops as it 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 they go down.

All the important trends across the period and group are given below in one place:

Trends in Chemical Properties

• Periodicity of Valence or Oxidation States

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.

• Periodic Trends and 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. The atomic and ionic radii, as we know, generally decrease in a period from left to right. As a consequence, the ionization enthalpies generally increase and electron gain enthalpies become more negative across a period. In other words, the ionization enthalpy of the extreme left element in a period is the least and the electron gain enthalpy of the element on the extreme right is the highest negative (note : noble gases having completely filled shells have rather positive electron gain enthalpy values). This results into high chemical reactivity at the two extremes and the lowest in the centre. Thus, the maximum chemical reactivity at the extreme left (among alkali metals) is exhibited by the loss of an electron leading to the formation of a cation and at the extreme right (among halogens) shown by the gain of an electron forming an anion. T

• 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.

Classification of elements and periodicity in Properties: Previous year's Question and Answers

Some important previous years' questions from this chapter are given below:

Question 1. Choose the incorrect trend in the atomic radii ( r ) of the elements :

i) $r_{B r}<r_K$

ii) $\mathrm{r}_{\mathrm{Al}}>\mathrm{r}_{\mathrm{Mg}}$

iii) $r_{\mathrm{Rb}}<\mathrm{r}_{\mathrm{Cs}}$

iv) $r_{\text {At }}<r_{C s}$

Answer:

$\mathrm{r}_{\mathrm{Mg}}>\mathrm{r}_{\mathrm{Al}}$ due to lower effective nuclear charge.

Magnesium has a larger atomic radius than aluminum, measuring about 160 pm for magnesium compared to about 143 pm for aluminum. This is primarily due to the effects of electronic configuration and effective nuclear charge.

Hence, the correct answer is option (2).

Question 2. Which of the following statements is correct?

A. The process of the addition of an electron to a neutral gaseous atom is always exothermic

B. The process of removing an electron from an isolated gaseous atom is always endothermic

C. The $1^{\text {st }}$ ionization energy of the boron is less than that of the beryllium

D. The electronegativity of C is 2.5 in $\mathrm{CH}_4$ and $\mathrm{CCl}_4$

E. Li is the most electropositive among the elements of group I

Choose the correct answer from the options given below

i) B and C only

ii)A, C, and D only

iii) B and D only

iv) B, C, and E only

Answer:

(A) The process of adding an $\mathrm{e}^{-}$to a neutral gaseous atom is not always exothermic; it may be exothermic or endothermic.
(C) Be B

$1 \mathrm{~s}^2 2 \mathrm{~s}^2 \quad 1 \mathrm{~s}^2 2 \mathrm{~s}^2 2 \mathrm{p}^1$

The Be 2s subshell is fully filled
So, high energy is needed to remove $\mathrm{e}^{-}$as compared to B.

(D) $\mathrm {In} \mathrm {CCl}_4 \rightarrow$ due to partially positive charge $z_{\text {eff }} \uparrow, \mathrm{EN} \uparrow$
So, EN of $\mathrm{C} \Rightarrow \mathrm{CCl}_4>\mathrm{CH}_4$
(E) Cs is the most electropositive.

Hence, the correct answer is option (1).

Question 3. The most electronegative element in the periodic table is-

(1) Nitrogen
(2) Oxygen
(3) Chlorine
(4) Fluorine

Answer:

Fluorine has the highest electronegativity because of its small size, high nuclear charge, low shielding, and strong electron affinity.

Hence, the correct answer is option (4).

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