Periodic Classification of Elements Class 10 Notes

Periodic Classification of Elements Revision Notes Class 10

Early Attempts at Classification:

Dobereiner's Triads (1829):

Johann Wolfgang Dobereiner observed similarities among certain elements. He grouped them into triads where the atomic mass of the middle element was approximately the average of the other two, and they exhibited similar chemical properties.

Example: Lithium (Li), Sodium (Na), Potassium (K)

Limitation: Could only identify a few triads.

Newlands' Law of Octaves (1864):

John Newlands arranged elements in increasing order of atomic masses and noticed that every eighth element had properties similar to the first, much like the notes in a musical octave.

Limitations:

• Applicable only up to Calcium.
• Did not account for undiscovered elements.
• Placed dissimilar elements together to fit the octave rule.

Mendeleev's Periodic Table (1869):

Dmitri Mendeleev formulated the Periodic Law, stating that the properties of elements are a periodic function of their atomic masses.

Key Features:

1. Elements were arranged in increasing order of their atomic masses.
2. Gaps for Undiscovered Elements: Mendeleev boldly left gaps in his table for elements that were yet to be discovered. He predicted the properties of these elements (e.g., Eka-Boron, Eka-Aluminium, Eka-Silicon), which were later confirmed with the discovery of Scandium, Gallium, and Germanium, respectively.
3. Prediction of Properties: His predictions were remarkably accurate, lending strong support to his classification.
4. Correction of Atomic Masses: Mendeleev also corrected the atomic masses of some elements, such as Beryllium and Indium, based on their predicted positions.

Limitations:

1. Position of Isotopes: Isotopes (atoms of the same element with different atomic masses) posed a challenge as they would have to be placed in different positions if atomic mass were the sole criterion.
2. Position of Hydrogen: Hydrogen's position was ambiguous as it exhibited properties similar to both alkali metals and halogens.
3. Anomalous Pairs: Some elements with higher atomic masses were placed before elements with lower atomic masses to maintain chemical similarity (e.g., Argon (atomic mass 39.9) was placed before Potassium (atomic mass 39.1)).

Modern Periodic Table (Moseley's Periodic Table):

Henry Moseley (1913) demonstrated that atomic number (the number of protons in an atom) is a more fundamental property of an element than its atomic mass.

Modern Periodic Law:

The properties of elements are a periodic function of their atomic numbers.

Key Features:

Elements are arranged in increasing order of their atomic numbers.

Periods:

Horizontal rows (1 to 7). The period number corresponds to the number of electron shells in the atoms of the elements belonging to that period.

Groups:

Vertical columns (1 to 18). Elements within the same group have the same number of valence electrons and thus exhibit similar chemical properties.

Solution to Mendeleev's Limitations:

Isotopes:

Since isotopes have the same atomic number, they are placed in the same position in the modern periodic table.

Anomalous Pairs:

The problem of anomalous pairs was resolved as elements are now arranged by atomic number, not atomic mass. For example, Argon (atomic number 18) correctly comes before Potassium (atomic number 19).

Position of Hydrogen:

While still a subject of discussion, hydrogen is generally placed in Group 1 due to its single valence electron.

Classification	Basis	Key Feature	Limitation
Dobereiner's Triads	Atomic mass	Groups of 3; mean mass rule	Few triads identified
Newlands' Octaves	Atomic mass	Every 8th element similar	Only up to Ca; anomalies
Mendeleev's Table	Atomic mass	Periodic law; gaps for new	Isotope position; anomalies
Modern Table	Atomic number	18 groups, 7 periods	Some anomalies remain

Trends in the Modern Periodic Table:

The periodic table allows us to predict the properties of elements based on their position.

Valency:

The combining capacity of an element.

Across a Period:

Valency first increases (from 1 to 4) and then decreases (from 4 to 0).

Down a Group:

Valency remains the same.

Atomic Size (Atomic Radius):

The distance from the center of the nucleus to the outermost shell of an isolated atom.

Across a Period (left to right):

Atomic size generally decreases. This is because the nuclear charge increases, pulling the electrons closer to the nucleus.

Down a Group (top to bottom):

Atomic size generally increases. This is due to the addition of new electron shells, increasing the distance between the nucleus and the outermost electrons.

Metallic Character:

The tendency of an atom to lose electrons.

Across a Period:

Metallic character decreases (elements become more non-metallic). This is because the effective nuclear charge increases, making it harder for atoms to lose electrons.

Down a Group:

Metallic character increases. This is due to the increase in atomic size and the decreasing effective nuclear charge, making it easier for atoms to lose electrons.

Non-Metallic Character:

The tendency of an atom to gain electrons.

Across a Period:

Non-metallic character increases.

Down a Group:

Non-metallic character decreases.

Electronegativity:

The tendency of an atom to attract electrons towards itself in a chemical bond.

Across a Period:

Increases.

Down a Group:

Decreases.

Nature of Oxides:

Metallic oxides are generally basic (e.g., Na₂O, MgO).

Non-metallic oxides are generally acidic (e.g., CO₂, SO₂).

Amphoteric oxides show both acidic and basic properties (e.g., Al₂O₃, ZnO).