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

Classification of Elements and Periodicity in Properties explains how elements are arranged in the periodic table according to atomic number and electronic configuration. In CBSE Class 11 Chemistry Chapter 3, students learn periodic law, element blocks, atomic size, ionisation enthalpy, electron gain enthalpy, electronegativity and reactivity trends.

The periodic table turns a large number of elements into a system students can study through patterns. As atomic number increases, electronic configuration changes in a regular order, and similar properties appear again in groups and periods.

These CBSE class 11 chemistry revision notes chapter 3 focus on how an element’s position helps predict its size, valency, metallic character and chemical reactivity. The chapter also connects historical classification, Modern Periodic Law and periodic trends with the present form of the periodic table.

Key Takeaways

• Modern Periodic Law: Element properties are periodic functions of their atomic numbers.
• Groups and periods: Elements in the same group have similar valence shell electronic configurations.
• Blocks: s, p, d and f blocks depend on the orbital that receives the last electron.
• Periodic trends: Atomic radius, ionisation enthalpy, electron gain enthalpy and electronegativity change in predictable ways.

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Classification of Elements and Periodicity in Properties Class 11 Chemistry Notes: Chapter Overview

The periodic table gives a systematic way to study elements. As atomic number increases, elements show patterns in electronic configuration, size, valency, metallic character and chemical reactivity.

This chapter begins with early attempts to classify elements and then moves to Modern Periodic Law. It also explains how groups, periods and blocks help students understand periodic trends across the periodic table.

Why Do We Need to Classify Elements?

Elements are the basic units of matter. As more elements were discovered, studying each element separately became difficult. Classification placed elements with similar properties together and made comparison easier.

Early Attempts to Classify Elements

Before the modern periodic table, scientists arranged elements using atomic weights and similarities in properties.

Dobereiner’s Triads

Dobereiner’s Triads were groups of three elements with similar properties. In each triad, the atomic weight of the middle element was nearly the average of the atomic weights of the other two elements.

This idea showed that elements could follow a pattern. It was later rejected because many elements could not be arranged into triads.

Newlands’ Law of Octaves

Newlands’ Law of Octaves arranged elements in increasing order of atomic weights. Newlands observed that every eighth element had properties similar to the first, similar to musical octaves.

Newlands’ idea was important because it introduced repetition of properties. It could not explain all known elements, so it was replaced by better classifications.

Lothar Meyer’s Curve

Lothar Meyer plotted physical properties such as atomic volume, melting point and boiling point against atomic weight. He observed a repeated pattern in these properties.

This supported the idea of periodicity. Mendeleev’s work became more popular because it arranged elements more completely and predicted undiscovered elements.

Mendeleev Periodic Law

Mendeleev Periodic Law states that the properties of elements are periodic functions of their atomic weights.

Mendeleev arranged elements in horizontal rows and vertical columns so that elements with similar properties appeared in the same group. He also left gaps for elements that had yet to be discovered.

Limitations of Mendeleev’s Periodic Table

Mendeleev’s periodic table was a major development, but some questions remained.

These problems were solved when atomic number became the basis of classification.

Modern Periodic Law

Modern Periodic Law states that the physical and chemical properties of elements are periodic functions of their atomic numbers.

Henry Moseley showed that atomic number is more fundamental than atomic mass. This changed the basis of the periodic table.

Long Form of the Periodic Table

The long form of the periodic table arranges elements in seven periods and eighteen groups. Periods are horizontal rows. Groups are vertical columns.

Elements in the same group have similar valence shell electronic configurations. This is why they show similar chemical properties.

Groups in the Periodic Table

Groups are vertical columns numbered from 1 to 18. Elements in the same group have similar outer electronic configurations.

Group position helps predict valency, ion formation and chemical reactivity.

Periods in the Periodic Table

Periods are horizontal rows. The period number shows the highest principal quantum number of the valence shell.

IUPAC Nomenclature of Elements Above Atomic Number 100

New elements are given temporary systematic names before official names are accepted. These names are based on numerical roots.

For example, atomic number 120 is named unbinilium with the temporary symbol Ubn.

Electronic Configuration and Periodic Table

Electronic configuration explains the position of an element in the periodic table. The orbital that receives the last electron decides the block.

The period number depends on the highest principal quantum number. The group depends on valence electrons and outer electronic configuration.

s-block Elements

s-block elements include Group 1 and Group 2 elements. Their last electron enters the s-orbital.

Group 1 elements are alkali metals. Group 2 elements are alkaline earth metals. They are reactive metals and usually form positive ions.

p-block Elements

p-block elements include Group 13 to Group 18 elements. Their last electron enters a p-orbital.

This block contains metals, non-metals, metalloids, halogens and noble gases. Group 18 elements have complete valence shells and show very low chemical reactivity.

d-block Elements

d-block elements are placed from Group 3 to Group 12. Their last electron enters the d-orbital of the penultimate shell.

These elements are also called transition elements. They commonly show variable oxidation states and form coloured compounds.

f-block Elements

f-block elements are placed separately at the bottom of the periodic table. Their last electron enters an f-orbital.

The f-block includes lanthanoids and actinoids. They are also called inner transition elements.

Metals, Non-metals and Metalloids

Elements can also be classified as metals, non-metals and metalloids.

Metals are found mainly on the left and centre of the periodic table. Non-metals are found mainly on the right side.

Atomic Radius

Atomic radius means the size of an atom. Since an atom has no sharp boundary, atomic size is measured using internuclear distance.

General order:

Van der Waals radius > Metallic radius > Covalent radius

Trend of Atomic Radius

Atomic radius decreases from left to right across a period. This happens because nuclear charge increases while electrons enter the same shell.

Atomic radius increases down a group because new shells are added, and valence electrons move farther from the nucleus.

Ionic Radius

Ionic radius is the size of an ion. Cations are smaller than their parent atoms because they lose electrons. Anions are larger than their parent atoms because they gain electrons.

For example, Na⁺ is smaller than Na, while F⁻ is larger than F.

Isoelectronic Species

Isoelectronic species are atoms or ions with the same number of electrons.

Their sizes differ because their nuclear charges are different. In an isoelectronic series, the species with higher nuclear charge has smaller radius.

Ionisation Enthalpy

Ionisation enthalpy is the energy required to remove the most loosely held electron from an isolated gaseous atom in its ground state.

It is usually expressed in kJ mol⁻¹.

Trend of Ionisation Enthalpy

Ionisation enthalpy generally increases across a period because nuclear charge increases and atomic size decreases.

It generally decreases down a group because atomic size and shielding increase.

Exception: Nitrogen and Oxygen

Nitrogen has a stable half-filled 2p³ configuration. Oxygen has a 2p⁴ configuration, where one orbital contains paired electrons.

Removing an electron from oxygen is easier because electron-electron repulsion exists in the paired orbital. So, nitrogen has higher first ionisation enthalpy than oxygen.

Electron Gain Enthalpy

Electron gain enthalpy is the enthalpy change when an electron is added to a neutral gaseous atom to form an anion.

X(g) + e⁻ → X⁻(g)

Electron gain enthalpy may be negative or positive. A negative value means energy is released. A positive value means energy is absorbed.

Trend of Electron Gain Enthalpy

Electron gain enthalpy generally becomes more negative across a period. This happens because atomic size decreases and the nucleus attracts the added electron more strongly.

It generally becomes less negative down a group because atomic size increases.

Chlorine has more negative electron gain enthalpy than fluorine because fluorine is very small and has high electron-electron repulsion in its compact shell.

Electronegativity

Electronegativity is the tendency of an atom to attract the shared pair of electrons in a chemical bond.

It is a relative property and depends on the bonding environment.

Fluorine is the most electronegative element in the periodic table.

Atomic Radius, Ionisation Enthalpy and Electronegativity Together

Periodic trends are connected. When atomic size decreases, the nucleus attracts electrons more strongly.

Valency and Oxidation State

Valency is the combining capacity of an element. It depends on the number of valence electrons.

For representative elements, valency may be equal to the number of valence electrons or eight minus the number of valence electrons.

Oxidation state represents the charge an atom appears to have in a compound.

Periodicity of Valence

Across a period, valence first increases and then decreases for representative elements.

Down a group, valence generally remains the same because elements in the same group have the same number of valence electrons.

Anomalous Properties of Second Period Elements

The first element of many groups shows behaviour different from the rest of the group. This is especially seen in second-period elements.

The main reasons are small atomic size, high electronegativity, high charge-to-radius ratio and absence of d-orbitals.

Diagonal Relationship

Diagonal relationship means similarity between diagonally placed elements of the second and third periods.

Lithium and magnesium show similar behaviour. Beryllium and aluminium also show similar behaviour.

This happens because their ionic sizes and charge-to-radius ratios are similar.

Chemical Reactivity of Metals

Chemical reactivity of metals depends on the tendency to lose electrons.

Across a period, metallic reactivity decreases because ionisation enthalpy increases. Down a group, metallic reactivity increases because the outer electron is lost more easily.

Group 1 metals become more reactive from lithium to caesium.

Chemical Reactivity of Non-metals

Chemical reactivity of non-metals depends on the tendency to gain electrons.

Across a period, non-metallic reactivity increases. Down a group, non-metallic reactivity decreases because the attraction for an added electron becomes weaker.

Halogen reactivity decreases from fluorine to iodine.

Nature of Oxides

The nature of oxides changes across a period.

Oxides of metals are generally basic. Oxides of non-metals are generally acidic. Oxides of some elements in the middle may be amphoteric.

Inert Pair Effect

Inert pair effect is the tendency of the outer s-electrons to remain unshared in bonding, especially in heavier p-block elements.

This effect makes lower oxidation states more stable in heavier elements of groups 13 to 16.

Important Facts About Elements

Classification of Elements and Periodicity in Properties Class 11 Chemistry Chapter 3 CBSE Notes: Quick Revision Tables

Blocks of Elements

Periodic Trends Summary

Historical Development Summary

Key Terms from CBSE Class 11 Chemistry Revision Notes Chapter 3

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