CBSE Class 10 Science Chapter Notes: Periodic Classification of Elements

The Evolution of the Modern Periodic Table

In 1913, the English physicist Henry Moseley made a revolutionary breakthrough by demonstrating that the atomic number (represented by the symbol Z) of an element is a far more fundamental property for classification than its atomic mass. This discovery corrected historical attempts and led to the formulation of the Modern Periodic Law, which states: 'The physical and chemical properties of elements are periodic functions of their atomic numbers.'

The resulting Modern Periodic Table systematically arranges elements into 18 vertical columns known as Groups and 7 horizontal rows known as Periods. Elements within the same group share identical valence shell configurations, directly translating to similar chemical reactivities. Conversely, when moving left-to-right across a period, the atomic number sequentially increases by one unit. The number of valence electrons also increases incrementally, while the total number of electron shells remains constant across that specific period.

Comparison of Early Classification Theories

Must Remember: Key Periodic Trends

• Valency: Across a period (left to right), valency first increases from 1 to 4 and then gradually decreases to 0. Down a group, valency remains constant because the number of valence electrons does not change.
• Atomic Size (Atomic Radius): Decreases moving left to right across a period due to an increase in nuclear charge which pulls the valence shell closer. Increases down a group because new electron shells are added.
• Metallic Character (Electropositive Nature): Decreases across a period because the tendency to lose electrons decreases. Increases down a group because atomic size increases, making it easier to lose electrons.
• Non-Metallic Character (Electronegative Nature): Increases across a period as elements gain electrons more easily due to higher effective nuclear charge. Decreases down a group.
• Electronegativity: The capacity of an atom to attract shared electrons increases across a period and decreases down a group.
• Chemical Reactivity: For metals, reactivity decreases across a period and increases down a group. For non-metals, reactivity increases across a period and decreases down a group.
• Nature of Oxides: Oxides of elements on the extreme left of the periodic table are basic (metallic oxides), while those on the right are acidic (non-metallic oxides).

Key Definitions to Memorize

Modern Periodic Law

The principle stating that the chemical and physical properties of elements are periodic functions of their atomic numbers.

Valency

The combining capacity of an atom, determined strictly by the number of valence electrons present in its outermost shell.

Atomic Radius

The shortest distance between the center of the nucleus of an atom and its outermost electron shell.

Effective Nuclear Charge

The net positive charge experienced by valence electrons from the nucleus, accounting for the shielding effect of inner-shell electrons.

Metalloids

Borderline elements that exhibit physical and chemical properties of both metals and non-metals (e.g., Boron, Silicon, Germanium).

Isotopes

Atoms of the same element that share the same atomic number but possess different mass numbers due to differing neutron counts.

High-Yield Revision Examples

• {"title":"Example 1: Predicting Periodic Table Position","description":"An element 'X' has an atomic number of 16. Determine its electronic configuration, period number, and group number.","solution":"1. Electronic Configuration: Divide the 16 electrons into shells: K=2, L=8, M=6. So, configuration is 2, 8, 6.\n2. Period: Since there are 3 filled shells (K, L, M), element 'X' belongs to Period 3.\n3. Group: There are 6 valence electrons in the outermost shell. For elements with 3 or more valence electrons, Group = 10 + valence electrons = 10 + 6 = Group 16 (Chalcogens)."}
• {"title":"Example 2: Comparing Atomic Radii","description":"Arrange Fluorine (F, Z=9), Nitrogen (N, Z=7), and Oxygen (O, Z=8) in increasing order of their atomic radii.","solution":"All three elements belong to the same period (Period 2). Moving left to right across a period, atomic size decreases due to the increase in effective nuclear charge which pulls the outer electrons closer. The sequence from left to right is N (7) -> O (8) -> F (9). Therefore, the increasing order of atomic radii is: F < O < N."}

Board Exam Traps & Marking Cues

1. Mendeleev's Gaps & Predicted Elements: Board papers frequently ask for Mendeleev's predicted elements. Remember this mapping clearly:

• Eka-Boron is Scandium (Sc)
• Eka-Aluminium is Gallium (Ga)
• Eka-Silicon is Germanium (Ge)

1. Using the Right Keywords: In reasoning questions about trends, examiners look for specific terms. Never just write 'atomic size decreases.' Always state: 'due to an increase in effective nuclear charge, which pulls the outermost shell closer to the nucleus.'

1. Noble Gases Trap: Note that Noble Gases (Group 18) have completely filled valence shells. Their atomic size is determined by Van der Waals radii, making them unexpectedly larger than the preceding halogens (Group 17). Keep this exception in mind if asked about a specific period's absolute size trend.

Quick Revision Checklist

• Why did Mendeleev leave vacant gaps in his Periodic Table? Mendeleev predicted the existence of elements that were undiscovered at that time. Instead of forcing existing elements into wrong positions, he left gaps and accurately forecasted their chemical properties.
• What happens to the electronegativity of elements as you move down a group, and why? Electronegativity decreases down a group. This is because atomic size increases, placing the valence shell farther from the nucleus, which reduces the nuclear pull on shared pairs of electrons.
• Why is the position of Hydrogen still considered anomalous? Hydrogen shows dual properties: it resembles alkali metals (Group 1) by forming positive ions, and also resembles halogens (Group 17) because it is a diatomic non-metal that can gain one electron to achieve stability.
• How does the basic nature of oxides vary across a period? The basic nature of oxides decreases across a period, while the acidic nature increases. Elements on the left form strongly basic metallic oxides, whereas elements on the right form acidic non-metallic oxides.

Frequently Asked Questions