CBSE Class 11 Chemistry Revision Notes Chapter 3

Class 11 Chemistry Revision Notes for Chapter 3: Classification of Elements and Periodicity in Properties

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Class 11 Chemistry Revision Notes for Chapter 3 – Classification of Elements and Periodicity in H3 – Properties – [Free Download]

Access Class 11 Chemistry Chapter 3 Classification of Elements and Periodicity in Properties

Introduction

The chapter introduces concepts of periodic classification of elements and will explore the properties of elements based on their size, valency, and electronic configurations. In this chapter, students will learn about Modern Periodic Law. An atom’s electronic configuration leads logically to a periodic classification. The periodic table groups all elements together and has them arranged in an order of matching properties or dissimilar ones.

Note: Similar elements, also known as like elements, are arranged in the same vertical column. These vertical columns are called groups or families due to the similar chemical properties of the elements. The chemical properties and valence electron number of all members of a family of elements are the same. The horizontal rows are called periods. 

Historical Development of Periodic Table

All previous classification attempts of the elements used their atomic weights.

Dobereiner’s Triads

Dobereiner proposed that each of the triads (groups of three elements) has a middle element with an atomic weight that is roughly halfway between the other two. The defining characteristic of the triads was that when the elements in a trio were ordered according to their increasing atomic weights, the atomic weight of the middle element was approximately equal to the arithmetic mean of the other two elements. The law of triads did not apply to all elements and thus was dismissed.

Newland’s Law of Octaves

The Law of Octaves, which was first proposed in 1865, was named after the musical octaves. It predicted that if elements were arranged in increasing order of their atomic weights, every eighth element would have characteristics eligible to be compared to the first one. However, it appeared that this was only accurate for elements up to calcium.

Lothar Meyer’s Curve

Lothar Meyer determined that after plotting various physical properties such as atomic volume, melting point and boiling point against the atomic weight, a periodically repeated pattern was obtained. Meyer noticed the change in the length of repetition, unlike Newland.

Mendeleev’s Periodic Law

Periodic law was first published by Dmitri Mendeleev. The law states that “The properties of the elements are a periodic function of their atomic weights.” 

Mendeleev’s system of classification involved arranging elements in horizontal rows and vertical columns of the periodic table in order of increasing atomic weights. Additionally, he made predictions about the existence of some elements that had not yet been discovered. 

Modern Periodic Law

In 1913, Henry Moseley observed the X-ray spectra of elements and discovered regularities in them. When √v ( v is the frequency of X-rays emitted) of emitted X-rays, was plotted against the atomic number (Z), the result was a straight line.

Due to this observation, there were corresponding changes made to  Mendeleev’s periodic law. The Modern Periodic law  states: “The physical and chemical properties of elements are periodic functions of their atomic numbers”

Periodic Table of Elements

Prediction of Block, Period, and Group :

1. Group – Similar electronic configuration 
2. Block – The orbital where the last electron enters
3. Period – Maximum value of the principal quantum number. Example : p – block – 10 + number of valence electron

Properties of an Element

The properties of an element have a periodic dependence on its atomic number.

Atomic radius

The atom’s lack of a distinct boundary makes it difficult to precisely measure the atomic size. There are various approaches to finding this.

Covalent Radius 

When two atoms of non-metallic elements are joined by a single bond in a covalent molecule, the distance between them that can be measured is known as the covalent radius. 

Eg. Bond distance in Cl = 198 pm, Covalent radius = 99pm

Covalent radius for A−A : For the same atom, rA = dA−A / 2

For two different atoms, dA−B= rA + rB− 0.09(χA−χB) 

Van der Waals Radius 

It refers to the radius of an imaginary hard sphere representing the distance of the closest approach to another atom.

Rvan der waal = dA−A / 2

Metallic Radius (Crystal Radius) 

It refers to half of the internuclear distance between two metal cores of a crystal. 

d= rA+rA / 2

rA = d / 2

Therefore, Van der Waals radius > Metallic radius > Covalent radius

Variation of Atomic Radii in the Periodic Table

• At the end of a period, the atomic radius decreases as the atomic nuclear charge increases. Outer valence electrons are in the same shell as inner valence electrons.
• As we move down a group, atomic radius increases because newer valence electrons are added. The primary quantum number (n) increases inside the group, valence electrons move further away.
• Ionic radius shows similarities to an atomic radius. The removal of an electron produces a cation, which shrinks in size while maintaining the same nuclear charge. A cation like sodium (Na+) will have a 95 pm radius, compared to Na having 186 pm. Anions are made when electrons are added to the atom, causing the electron repulsion higher, and the net nuclear charge to decrease. Eg. The atomic radius of the fluoride ion, Fl– (anion), is 136 pm, while the fluorine atom has 64 pm.
• Isoelectronic species are atoms and ions with the same number of electrons but with different radii as a result of different nuclear charges. An anion with a greater negative charge will have a larger radius than a cation with a greater positive charge.

Note: Moving from left to right in a period, the covalent radius and van der Waals radius tend to usually decrease as the atomic number increases.

General Trend

Ionisation Energy

The least amount of energy necessary to separate an electron from an isolated atom’s outermost orbit while it is in a gaseous state is known as ionisation energy.

Units of I.E/I.P

An electron volt is the energy gained by an electron while travelling under a potential difference of one volt. Ionisation enthalpy is expressed in units of an electron volt (eV) per atom, kilocalories per mole (kcal mol-1) or kilojoules per mole (kJ mol-1). 

1 electron volt per atom = 23.06 kcal mol−1 = 96.49 kJ mol−1

Important Points

• As atomic size decreases, the ionisation enthalpy of an atom increases.
• An increase is observed in ionisation enthalpy as the screening defect decreases.
• Rise is observed in ionisation energy in case of orbitals being entirely or half-filled
• The increasing order of penetration of different orbitals is f < d < p < s

Electron Gain Enthalpy

As an electron gets added to a neutral gaseous atom to turn it into a negative ion, an enthalpy change accompanying it is called the electron gain enthalpy. This can be used to assess how easy it is to add an electron to produce anions.

X (g) + e– → X– (g)

The element determines whether this addition is exothermic or endothermic. Halogens (Group 17 ) take up electrons to achieve stable noble gas electrical configurations and thus have high negative electron gain enthalpies. Noble gases do not release electrons easily and thus have high positive  ΔegH values.

Variation of Electron Gain Enthalpy

Along a period, from left to right, ΔegH increasingly turns negative. As the atom gets smaller, the nucleus exerts more of a pull on the electron to be added.

Along a group, from top to bottom, ΔegH becomes less negative owing to the growing size of atoms and an enhancement in nuclear charge.

Electronegativity

The propensity of an atom to attract a shared pair of electrons to itself is known as electronegativity. It can be determined using Pauling’s scale and is derived as follows: 

χA−χB)=0.208 [EA−B− (EA−A× EB−B)1/2]½

Mulliken suggested,

Electronegativity = ( Ionisation potential + Electron affinity ) / 2

• 16(χA−χB)+3.5(χA−χB)2 gives the percentage ionic character, where χA and χB are electronegativities of A and B.
• If subtracting the electronegativities of combining atoms achieves a value of 1.7, the bond is 50% covalent and 50% ionic.
• The oxide takes on a basic character when the difference is quite big between the electronegativities of oxygen and the element.

The Periodic Trends of Elements in the Periodic Table

Periodic Trends in Chemical Properties 

Periodicity of Valence or Oxidation States

An atom’s valency is determined by the electrons in its outermost shell.  The shell is called the valence shell, and the orbitals contained in this shell are called valence orbitals. Generally, valency can be determined as 8 minus the number of valence electrons. The most frequent values for transition and inner transition elements are 2 and 3, though they can have other values.

• Along a period, valency is increased as we move from left to right. However, with hydrogen and oxygen, it first increases from 1 to 4 and then reduces to zero. In the production of the Na2O  molecule, the more electronegative oxygen receives two electrons from each of the two sodium atoms, giving it an oxidation state of –2. In contrast, sodium’s valence shell electronic configuration of 3s1 causes it to lose one electron to oxygen and have an oxidation state of +1.
• Across a group, the valency of all atoms remains the same. All elements in group 2 have valency 2. For inert gases, the valency is zero since they are inert gases. 

Anomalous Properties of Second Period Elements

Diagonal relationships exist between the elements of the second period and the elements of the third period. When these elements are placed diagonally next to each other in the periodic table, they have properties that are similar. For example, Lithium from group 1 resembles Magnesium from group 2. These anomalies are attributed to their short sizes, huge charge/radius ratio and high electronegativities. The group’s first member has only two valence orbitals (2s and 2p) that can be used for bonding, while the group’s second member has nine valence orbitals (3s, 3p, 3d). Since the other members of the groups can expand their valence shell to accommodate more than four pairs of electrons, the first member of each group has a maximum covalency of four (for example, boron can only form [BF4]). such as aluminium [AlF6]3-. Furthermore, the first member of the p-block elements exhibits a greater capacity for forming multiple p-p bonds with both other second-period elements and with itself (e.g., C=C, C≡C, N=N, N≡N).

Periodic Trends and Chemical Reactivity

Reactivity of Metals

The tendency to lose electrons from the outer shell of metals determines their reactivity.

• Across a period, from left to right, metal reactivity will decrease.
• Across a group, from top to bottom, metal reactivity increases due to increasing size and thus an increased tendency to lose electrons. 

Reactivity of Non–Metals

The formation of anions by non-metals upon gaining electrons determines their reactivity.

• Across a period, from left to right, the reactivity of non-metals increases as anions are formed during reactions. 
• Moving down in a group, the reactivity of non-metals decreases as the tendency to accept electrons decreases.  

Inert Pair Effect

Further along group 13 to group 16, the tendency of bond formation by elements decreases as the atomic size increases. The s-electrons are more tightly gripped and deny participating in bond formation, making the lower oxidation stable. 

Summary and Important Points to Remember

• The s -, p -, d -, and f – blocks of elements are named for the orbital that receives the last electron. These blocks’ electronic configuration is as follows:

1. s – block: Noble Gas ns1−2, but hydrogen has a 1s1 configuration.
2. p – block: Noble Gas ns2np1−6 
3. d – block: Noble Gas (n−1)d1−10ns1−2
4. f – block: Noble Gas (n−2)f1−14(n−1)d0−1ns2

• Elements can be classified as metals, nonmetals and metalloids 

1. Metals constitute 78% of all elements.
2. Nonmetals are less than twenty in number
3. Metalloids exist on the boundary between metals and nonmetals. 

Some Important Facts About Elements

• Fluorine – The most electronegative element in the periodic table
• Tungsten (W) – The metal with the highest melting point 3380°C 
• Bromine – A liquid non-metal (at room temperature)
• Chlorine – The highest negative electron gain enthalpy of any element.

Class 11 Chemistry Chapter 3 Notes Download

A detailed understanding and insight into various concepts regarding trends in the periodic table are crucial to cracking all questions related to Class 11 Chemistry Chapter 3. Download Class 11 Chemistry Chapter 3 PDF for studying these points at your own convenience. 

Important Terms Determining Periodic Tendency in Element Properties

• Atomic sizes 
• Ionisation enthalpies
• Electron gain enthalpies
• Electronegativity
• Chemical reactivity

These phenomena, which result from an element-based chemical reaction, are explained in detail in Chemistry Class 11 Chapter 3 Notes. Additionally, there are structural characteristics in the Chapter 3 Chemistry Class 11 Notes that are important to understand.

Terms Related to Atomic Structure in Class 11 Chemistry Chapter 3 Notes

• Groups – They are the vertical rows of the periodic table that are all unrelated to one another.
• Periods – The periodic table’s horizontal rows are referred to as rows.
• Metalloids – Elements that exhibit both metallic and nonmetallic characteristics are called metalloids.

From the perspective of the exam, Chapter 3 of Chemistry for Grade 11 contains many important points. Students can refer to the revision notes carefully crafted by subject matter experts to maximise their exam preparation and easily grasp concepts.