Chapter 3 Classification of Elements and Periodicity in Properties class 11

 Chapter 3: Classification of Elements and Periodicity in Properties

By Aarish Sir 

Classification of Elements and Periodicity in Properties  Introduction

• Elements are the fundamental substances that make up all matter.

• Till now, 118 elements are known.

• To study such a large number of elements easily, scientists classified them into groups based on similar properties.

• This classification gave rise to the Periodic Table, which is one of the most important tools in chemistry.

2. Early Attempts of Classification

(a) Dobereiner’s Triads (1817)

• J.W. Dobereiner arranged elements in groups of three (triads) having similar chemical properties.

• The atomic mass of the middle element was roughly the average of the atomic masses of the other two.

Example:

• Li (7), Na (23), K (39) → Na ≈ (7+39)/2 = 23

• Limitation: Worked only for a few elements.

(b) Newlands’ Law of Octaves (1866)

• John Newlands arranged elements in increasing order of atomic mass.

• Every 8th element had properties similar to the first (like musical notes: sa, re, ga, ma, pa, dha, ni).

Example: Li, Be, B, C, N, O, F, Na → Li and Na similar.

• Limitation: Worked only up to calcium; failed for heavier elements.

3. Mendeleev’s Periodic Table (1869)

• Dmitri Mendeleev arranged elements in increasing order of atomic mass.

• He left gaps for undiscovered elements and predicted their properties (e.g., Eka-silicon = Germanium).

• Achievements:

  • First successful classification.

  • Predicted new elements correctly.

• Limitations:

  • Position of isotopes not explained.

  • Increasing order of mass not always correct (e.g., Ar (40) comes before K (39)).

4. Modern Periodic Law and Table

• Proposed by Moseley (1913).

• Modern Periodic Law: Physical and chemical properties of elements are a periodic function of their atomic number.

• This solved problems of Mendeleev’s table.

Features of Modern Periodic Table (Long Form)

• Horizontal rows → Periods (7 periods).

• Vertical columns → Groups (18 groups).

• Elements arranged according to atomic number.

• Group number indicates valence electrons, period number indicates valence shell.

5. Classification of Elements in the Periodic Table

1. s-block elements

   • Groups 1 and 2 (alkali metals and alkaline earth metals).

   • Valence electron in s-orbital.

   • Highly reactive metals.

2. p-block elements

   • Groups 13 to 18.

   • Valence electron in p-orbital.

   • Include metals, non-metals, metalloids, noble gases.

3. d-block elements (Transition elements)

   • Groups 3 to 12.

   • Valence electron enters d-orbital.

   • Show variable oxidation states, form colored ions.

4. f-block elements (Inner transition elements)

   • Lanthanides and actinides.

   • Valence electron enters f-orbital.

   • Mostly radioactive (actinides).

6. Periodic Properties and Their Trends

(a) Atomic Radius

• Distance from nucleus to outermost shell.

• Across a period: decreases (more nuclear charge pulls electrons closer).

• Down a group: increases (more shells are added).

(b) Ionic Radius

• Cation (+ ion): smaller than atom (due to loss of electrons).

• Anion (– ion): larger than atom (due to gain of electrons).

(c) Ionization Enthalpy

• Energy required to remove an electron from gaseous atom.

• Across a period: increases (nuclear charge increases).

• Down a group: decreases (outer electrons farther from nucleus).

(d) Electron Gain Enthalpy (Electron Affinity)

• Energy change when an electron is added to a neutral atom in gaseous state.

• Across a period: becomes more negative (tendency to gain e⁻ increases).

• Down a group: becomes less negative.

• Exception: Noble gases have positive values (do not accept electrons).

(e) Electronegativity

• Tendency of atom to attract shared pair of electrons in a bond.

• Across a period: increases.

• Down a group: decreases.

• Most electronegative element: Fluorine (F).

(f) Metallic and Non-Metallic Character

• Metallic character → tendency to lose electrons.

• Across a period: decreases.

• Down a group: increases.

• Non-metallic character shows opposite trend.

(g) Valency

• Number of valence electrons or combining capacity of an element.

• Across a period: first increases (1 to 4), then decreases (4 to 0).

• Down a group: remains same.

7. Anomalies in Periodic Properties

• Inert gases (Group 18) have very high ionization energy and zero electron affinity.

• Transition metals show irregular trends due to d-orbital electrons.

• Lanthanides show “lanthanide contraction” (gradual decrease in atomic radius).

8. Importance of Periodic Table

• Helps to study properties of elements in a systematic way.

• Predicts chemical behavior of unknown elements.

• Basis of modern inorganic chemistry.

Mendeleev’s Periodic Law. Physical and chemical properties of elements are periodic function of their atomic masses.
• Modem Periodic Law. Physical and chemical properties of the elements are periodic function of their atomic numbers.
• Groups. There are 18 groups. These are vertical rows.
• Periods. There are 7 periods. These are horizontal rows.
• Representative Elements. The S and P block of elements are known as representative elements.
• Transition Elements. They are also called d-block elements. They have general electronic configuration (n – 1) d^1-10 ns^0-2.
• Inner Transition Elements. Lanthanoids (the fourteen elements after Lanthanum) and actinides (the fourteen elements after actinium) are called inner transition elements. General electronic configuration is (n – 2) f^1-14(n – 1) d^0-1 ns².
They are also called f-block elements.
• Metals. Present on the left side of the periodic table. Comprise more than 78% of the known elements.
• Non-metals. Mostly located on the right hand side of the periodic table.
• Metalloids. Elements which line as the border line between metals and non-metals (e.g., Si, Ge, As) are called metalloids or semimetals.
• Atomic Radii and Ionic Radii, increase down the group decrease along the period.
• Ionization Enthalpy. Increases along the period and decreases down the group.
• Noble Gas Elements. Elements with symmetrical configuration are chemically inert in nature.
• Electric Nuclear Charge. Z = Nuclear charge – Screening constant.
• Electronegativity. Increases along a period decreases down the group,
• Chemical Reactivity. Chemical reactivity is highest at the two extremess of a period and lowest in the centre.
• Oxides of Elements. Oxides formed of the Elements on the left are basic and of elements
on the right are acidic in nature.
Oxides of elements in the centre are amphoteric or neutral