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CBSE Class 11 Chemistry Notes Chapter 3 PDF Download

Here, we have provided CBSE Class 11 Chemistry Notes Chapter 3 Classification of Elements and Periodicity in Properties. Students can view these ch 3 chemistry class 11 notes before exams for better understanding of the chapter.
authorImageAnanya Gupta21 Jul, 2025
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CBSE Class 11 Chemistry Notes Chapter 3

CBSE Class 11 Chemistry Notes Chapter 3:  In ch 3 chemistry class 11 notes, we learn about how elements are grouped and organized based on their properties.

This organization helps us understand the behavior of different elements better. We study the periodic table, which is like a big chart showing all the elements arranged in a specific order.

This order is based on things like the number of protons in an atom and how the electrons are arranged around the nucleus. By studying this, we can see patterns in how elements behave.

For example, some elements are more reactive than others, and some are better conductors of electricity. Understanding these patterns helps us understand chemistry better.

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

Dobereiner's Triads

Dobereiner's Triads were proposed by the German chemist Johann Wolfgang Dobereiner. He attempted to classify elements with similar properties into groups of three elements each, called "triads.

" The key idea behind these triads was that the atomic mass of the element in the middle of the triad would be approximately equal to the average of the atomic masses of the other two elements.

For instance, one of Dobereiner's triads consisted of lithium (Li), sodium (Na), and potassium (K). The atomic mass of lithium is 6.94, and that of potassium is 39.10. Interestingly, the atomic mass of sodium (22.99) falls nearly in the middle of the atomic masses of lithium and potassium.

Dobereiner's Triads had limitations:

Not all elements known at that time could be classified into triads. Only four triads were identified by Dobereiner, leaving many elements unaccounted for.

Newland’s Octaves

John Newlands, an English scientist, made significant contributions to the early development of the periodic table. In 1866, he arranged the 56 known elements in increasing order of atomic mass. Newlands observed a pattern where every eighth element exhibited properties similar to the first. This observation led to the formulation of Newland’s Law of Octaves. According to this law, when elements are arranged in increasing order of atomic mass, the properties of two elements with an interval of seven elements between them would be similar.

Newland’s octaves had limitations:

The classification of elements based on octaves was only successful up to calcium. The discovery of noble gases posed a challenge to Newland’s arrangement, as they did not fit into the periodic pattern without disrupting it completely.

Mendeleev’s Periodic Table

Dmitri Ivanovich Mendeleev, a Russian chemist, introduced his periodic table in 1869. He noticed that the properties of elements, both physical and chemical, repeated periodically with their atomic masses. Mendeleev's Periodic Law, also known as Mendeleev’s Law, states that the chemical properties of elements are a periodic function of their atomic weights.

Advantages of Mendeleev’s Periodic table:

It accommodated newly discovered elements without disturbing the table's structure, including germanium, gallium, and scandium. Mendeleev’s table helped correct inaccurate atomic weights that were prevalent at the time. It introduced variations from the strict order of atomic weights.

Limitations of Mendeleev’s Periodic table:

Hydrogen's position in the group of alkali metals contradicted its halogen-like qualities. Isotopes were not considered, leading to inconsistencies in the placement of elements like protium, deuterium, and tritium.

Some elements were anomalously positioned, such as cobalt (atomic mass 58.9) appearing before nickel (atomic mass 58.7). These methods laid the groundwork for the modern periodic table, with Mendeleev being its most significant contributor. He is often referred to as the Father of the Modern Periodic Table, and the modern periodic law is also called Mendeleev’s Law in his honor.

Modern periodic table

In 1913, English physicist Henry Moseley conducted groundbreaking research on the characteristic x-rays emitted by different metals. He discovered a direct correlation between the square root of the frequency of these x-rays and the atomic number of the elements. Based on his findings, Moseley formulated the modern periodic law, which states:

"The physical and chemical properties of elements are periodic functions of their atomic numbers."

Unlike atomic mass, which is influenced by the mass of protons and neutrons in the nucleus, the atomic number solely determines the number of electrons in an atom. Since chemical properties primarily depend on the arrangement of electrons in different energy levels, elements with different electronic configurations exhibit distinct chemical behaviors. Therefore, Moseley argued that atomic number, not atomic mass as proposed by Mendeleev, should be the basis for classifying elements in the periodic table. The periodic repetition of similar properties among elements grouped by their atomic numbers is known as periodicity. This concept highlights the predictable patterns in the chemical and physical properties of elements as their atomic numbers increase.

Classification of elements in modern periodic table

The modern periodic table comprises 18 vertical columns known as groups (1-18) and 7 horizontal rows known as periods. Each period and group represent distinct characteristics of the elements.
  • The first period consists of only two elements: Hydrogen and Helium.
  • The second period includes eight elements, starting from Lithium and ending with Neon.
  • The third period also contains eight elements, from Sodium to Argon.
  • The fourth period encompasses eighteen elements, extending from Potassium to Krypton.
  • The fifth period also holds eighteen elements, ranging from Rubidium to Xenon.
  • The sixth period consists of thirty-two elements.
  • The seventh period is incomplete.
Elements in the periodic table are classified into four blocks based on their electronic configuration: s-block, p-block, d-block, and f-block.
  • The elements in the 1st and 2nd groups are termed s-block elements, with a general electronic configuration of ns1-2.
  • Elements in groups 13th to 18th are known as p-block elements, having a general electronic configuration of ns2 np1-6.
  • The elements from the 3rd to 12th groups belong to the d-block, with a general electronic configuration of (n-1)d1-10 ns1-2.
  • Lanthanides and actinides, placed separately at the bottom of the periodic table, are known as f-block elements, with a general electronic configuration of (n-2)f1-14 (n-1)d0-1 ns2.
This systematic classification helps in understanding the properties and behavior of different elements within the periodic table.

Periodic properties and their trends

The periodic properties of elements are closely related to their electronic configuration and exhibit a gradual change as we move down a group or across a period. Key physical properties such as melting points, boiling points, density, and enthalpy of fusion and vaporization are influenced by electronic configuration. However, our focus lies primarily on properties like atomic and ionic radii, ionization enthalpy, electron gain enthalpy, and electronegativity.

Atomic and Ionic Radii:

Atomic Radii: The distance from the nucleus to the outermost electron shell defines atomic radii, which can be covalent, van der Waals, or metallic.

Covalent radii: Half the distance between nuclei of adjacent atoms in a single covalent bond.

Van der Waals radii: Half the internuclear distance between atoms of neighboring molecules in a solid.

Metallic radii: Half the distance between nuclei of adjacent atoms in a metallic crystal.

Atomic radii decrease across periods due to increased effective nuclear charge and increase down groups owing to additional electron shells and shielding effect.

Ionic Radii: The effective distance from the nucleus to the electron cloud for ions formed from neutral atoms.

Ionic radii follow the same trend as atomic radii, decreasing across periods and increasing down groups.

Properties of an Element

Atomic radius is a fundamental concept in chemistry that refers to the size of an atom. However, due to the electron's probability distribution, it's challenging to measure the exact size of an isolated atom. Therefore, different forms of atomic radius are utilized depending on the atom's surroundings.

These include covalent radius, van der Waals' radius, and metallic radius. Covalent radius is defined as half the distance between the nuclei of two identical atoms bonded by a single covalent bond. If the bond is formed between different elements, the covalent radius is calculated based on their electronegativities.

Van der Waals' radius is half the internuclear distance between adjacent atoms of neighboring molecules in the solid state. Metallic radius, also known as crystal radius, is half the distance between the nuclei of neighboring metal atoms in a metallic crystal lattice. In general, the order of atomic radii is: covalent radius < metallic radius < van der Waals' radius. The variation of atomic radii in the periodic table follows certain trends.

Along a period, as we move from left to right, the covalent and van der Waals' radii generally decrease due to increasing effective nuclear charge. Along a group, the atomic radius increases with the increasing number of atoms in the group, as the valence electrons move farther away from the nucleus with increasing principal quantum number (n). Ionic radius, which refers to the size of ions, follows similar trends as atomic radii.

Cations (positive ions) are smaller than their parent atoms due to the removal of electrons, while anions (negative ions) are larger due to the addition of electrons. Isoelectronic species, which have the same number of electrons, exhibit variations in radius depending on their nuclear charges.

Ionization Enthalpy

The energy required to remove an electron from the outermost orbit of an isolated gaseous atom. Increases across periods due to increased nuclear charge and decreases down groups.

Electron Gain Enthalpy

The change in enthalpy when a gaseous atom gains an electron to form a monovalent anion. Increases across periods and decreases down groups, with exceptions for half-filled or fully filled orbitals.

Electronegativity

  • The ability of an atom to attract shared electrons in a covalent bond.
  • Fluorine is the most electronegative, while cesium is the least.
  • Electronegativity increases across periods and decreases down groups.

Periodic Trends in Chemical Properties

Valence electrons are crucial in determining the chemical behavior of atoms, as they reside in the outermost shell and are involved in bonding. The valence, or valency, of an atom corresponds to the number of these electrons. For s- and p-block elements, the valence is typically equal to the number of valence electrons.

Transition and inner transition elements, however, exhibit variable valences due to the presence of d- or f-electrons, although common valences are often 2 and 3. Understanding valence periodicity helps comprehend how valence states change across periods and within groups in the periodic table.  Along a period, the number of valence electrons increases from 1 to 8 as we move from left to right. However, when considering the valence in compounds relative to hydrogen or oxygen, it peaks at 4 before declining to 0. For instance, in Na2O, oxygen, being more electronegative, gains two electrons from each sodium atom, leading to an oxidation state of -2, while sodium loses one electron, resulting in an oxidation state of +1. Within a group, the number of valence electrons remains constant, resulting in elements within the group sharing the same valence. For instance, all alkali metals in group 1 have a valence of one, while alkaline earth metals in group 2 have a valence of two. Noble gases in group 18 have zero valence, as they are chemically inert and do not form bonds.

Anomalous Properties of Second Period Elements

Diagonal relationships in the periodic table occur when elements from the second period share similarities with elements from the third period when placed diagonally across each other, despite being in different groups. 

For instance, lithium in group 1 resembles magnesium in group 2, while beryllium in group 2 resembles aluminium in group 13, and so forth. This phenomenon arises due to their comparable attributes, such as size, charge-to-radius ratio, and electronegativity. The unique behavior of these elements is attributed to several factors. Firstly, the first member of each group has only four valence orbitals available for bonding (2s and 2p), whereas the second member of the group has nine valence orbitals (3s, 3p, 3d).  Consequently, the maximum covalency of the first member is limited to 4, while the subsequent members can expand their valence shell to accommodate more than four pairs of electrons.

For example, boron can only form compounds like [BF4]−, whereas aluminium can form complexes like [AlF6]3−. Additionally, the first member of p-block elements exhibits a greater tendency to form p-p multiple bonds within itself (e.g., C=C, C≡C, N=N, N≡N) and with other second-period elements (e.g., C=O, C=N, C≡N, N=O) compared to later members of the same groups. This propensity for multiple bonding contributes to their unique chemical behavior and diagonal relationship with elements in different periods and groups.

Periodic Trends and Chemical Reactivity

Reactivity in metals is determined by their ability to lose electrons from their outermost shell. When examining reactivity trends: Within a period: Moving from left to right across a period, the tendency to lose electrons decreases, leading to a decrease in metal reactivity.

For example, the reactivity order for third-period elements is Na > Mg > Al. Within a group: Reactivity increases as we move down a group because the tendency to lose electrons rises. In group 1, for instance, the reactivity order is Li < Na < K < Rb < Cs. On the other hand, non-metal reactivity is determined by their tendency to gain electrons and form anions: Within a period: Non-metal reactivity increases from left to right across a period, with non-metals tending to form anions during reactions.

For example, in the second period, reactivity follows the order C < N < O < F. Within a group: Non-metal reactivity decreases as we move down a group because the inclination to accept electrons decreases. The reactivity order for halogens is F > Cl > Br > I. Regarding oxides, the basic normal oxide formed by the leftmost element (e.g., Na2O) is the most basic, while the acidic normal oxide formed by the rightmost element is the most acidic (e.g., Cl2O7). Oxides of elements in the center may be amphoteric (e.g., Al2O3, As2O3) or neutral (e.g., CO, NO, N2O). The inert pair effect refers to the declining tendency of valence shell electrons to participate in bond formation as we move through groups 13-16.

This leads to lower oxidation states becoming more stable. This effect occurs because the increasing nuclear charge overcomes the rise in atomic size as we move down these groups, resulting in s-electrons becoming less willing to participate in bond formation, thus stabilizing lower oxidation states.

CBSE Class 11 Chemistry Notes Chapter 3 PDF

You can read the notes for Chapter 3 of CBSE Class 11 Chemistry, titled "Classification of Elements and Periodicity in Properties," by clicking on the PDF link provided below.

This chapter talks about how scientists organize elements based on their properties and why they arrange them this way. It helps you understand why some elements behave similarly and others don't.

The notes in the PDF cover all the important points in an easy-to-understand manner, making it easier for you to learn and remember

 

CBSE Class 11 Chemistry Notes Chapter 3

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Benefits of CBSE Class 11 Chemistry Notes Chapter 3 

Concept Clarity: These notes provide a structured overview of the classification of elements and periodic trends, helping students understand fundamental concepts clearly.

Comprehensive Coverage: The notes cover all essential topics and subtopics of the chapter, ensuring that students have a thorough understanding of the subject matter.

Simplified Explanation: Complex concepts are explained in a simplified manner, making it easier for students to grasp and retain information.

Quick Revision: Students can use these notes for quick revision before exams or assessments, helping them consolidate their knowledge effectively.

Practice Questions: The notes include practice questions and examples that enable students to test their understanding and reinforce their learning.

Exam Preparation:  By studying these notes, students can also  practice Important Questions for Class 11 Chemistry for their exams, as they cover the entire syllabus prescribed by the CBSE board.

CBSE Class 11 Chemistry Notes Chapter 3 FAQs

What is the Periodic Law?

The Periodic Law states that the properties of elements are a periodic function of their atomic numbers. This means that elements exhibit similar properties after certain regular intervals when arranged in order of increasing atomic number.

What are the different blocks in the periodic table based on electronic configuration?

Elements are classified into four blocks based on their electronic configuration: s-block, p-block, d-block, and f-block. Each block corresponds to a specific range of electron configurations.

How does atomic radius vary across a period and down a group?

Across a period, atomic radius generally decreases due to increasing effective nuclear charge. Down a group, atomic radius increases due to the addition of new energy levels and increased shielding effect.

Explain electron gain enthalpy and its periodic trends.

Electron gain enthalpy is the energy released when an atom gains an electron. It generally becomes more negative across a period (except for noble gases) and decreases down a group.

What is electronegativity? How does it vary in the periodic table?

Electronegativity is the tendency of an atom to attract shared electrons in a chemical bond. It generally increases across a period and decreases down a group.
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