Chapter 5: Periodic Classification of Elements

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Introduction to Periodic Table

History:
- Döbereiner's Triads: Elements with similar properties grouped in threes, middle element had atomic weight approximately average of the others.
- Newlands' Law of Octaves: Every eighth element had properties similar to the first, but was limited to lighter elements.
- Mendeleev's Periodic Table: Organized elements by increasing atomic mass, predicting properties of undiscovered elements; had some anomalies.
Modern Periodic Table:
- Based on atomic number (number of protons), developed by Henry Moseley.

Modern Periodic Law

Statement: "The properties of elements are periodic functions of their atomic numbers."

Structure of the Modern Periodic Table

Periods: Horizontal rows; indicate the energy level of the outermost electron.
- There are 7 periods in the current table.
Groups: Vertical columns; elements in the same group have similar chemical properties due to the same number of valence electrons.
- Numbered 1 to 18, with the older system using Roman numerals with A and B (e.g., IA, IIA, etc.).

Types of Elements

Metals: Left side of the periodic table; lose electrons to form positive ions.
Non-metals: Right side; gain electrons to form negative ions.
Metalloids: Along the stair-step line, have properties intermediate between metals and non-metals.

Trends in the Periodic Table

Atomic Radius:

Decreases across a period: Due to the increase in nuclear charge pulling electrons closer.
Increases down a group: Additional electron shells increase size.

Ionization Enthalpy (Energy):

Increases across a period: More energy needed to remove an electron as nuclear charge increases.
Decreases down a group: Electrons are farther from the nucleus, easier to remove.

Electron Affinity:

Generally increases across a period: More energy released when an electron is added due to increased nuclear attraction.
Decreases down a group: For the same reason as ionization energy, though there are exceptions.

Electronegativity:

Increases across a period: Elements tend to attract electrons more as we move from left to right.
Decreases down a group: Due to increasing atomic size and distance from the nucleus.

Metallic Character:

Non-Metallic Character:

Valency:

Changes across a period: Varies from 1 to 4 then back to 1, then 0 for noble gases.
Constant in a group: For main group elements, valency is generally the same for elements in the same group.

Periodic Trends and Chemical Reactivity

Reactivity of Metals: Increases down a group (easier to lose electrons); decreases across a period.
Reactivity of Non-metals: Increases across a period (more electron affinity); decreases down a group.

Groups in Detail

Group 1 (Alkali Metals): Highly reactive, form +1 ions, stored in oil.
Group 2 (Alkaline Earth Metals): Less reactive than Group 1, form +2 ions.
Group 17 (Halogens): Very reactive non-metals, form -1 ions, reactivity decreases down the group.
Group 18 (Noble Gases): Inert, full valence shells, do not readily form compounds.

d-Block Elements (Transition Elements)

Placed in the middle of the periodic table between Groups 2 and 3.
Show variable oxidation states, colored compounds, catalytic properties, and form complexes.

f-Block Elements (Lanthanides and Actinides)

Placed at the bottom of the periodic table; not part of the main body due to space constraints.
Lanthanides follow Lanthanum; Actinides follow Actinium. Show similar chemical properties within their series.

Applications of the Periodic Table

Prediction of Properties: Helps in predicting chemical behavior of unknown elements.
Chemical Bonding: Understanding how elements will bond based on their position.
Industrial and Technological Uses: Selection of materials based on properties, like in electronics, batteries, and alloys.

Conclusion

The periodic table is a fundamental tool in chemistry, providing insights into the nature of elements, how they react, and their physical properties. The periodic trends are key to understanding chemical reactions and the design of new materials.