3-Chemistry-Module-Chapter-3-Periodicity-of-Elements-1

Chapter III: PERIODICITY OF ELEMENTS

Introduction to Periodicity

  • Elements exhibit repetitive structures in their properties and reactivity, similar to crystals.

  • Periodicity refers to this unique repetitiveness of properties occurring at regular intervals.

  • The periodic table summarizing these elements includes details like atomic numbers, masses, and more.

Early Attempts to Classify Elements

  • Johann Wolfgang Dobereiner (1817)

    • Proposed the Law of Triads where elements grouped in threes have similar properties.

    • Example: Lithium (7), Sodium (23), Potassium (39) with sodium's atomic mass being the average of the other two.

  • John Newlands (1860)

    • Proposed the Law of Octaves stating properties repeat every eight elements based on atomic number.

    • Example: First octave elements (1H to 8O) display similar properties to the second octave (9F to 16S).

    • This law was later rejected due to inconsistencies, particularly past Calcium (20Ca).

Development of the Modern Periodic Table

  • Dmitri Mendeleev and Julius Lothar Meyer (1869)

    • Independently arranged elements by increasing atomic mass, forming groups with similar properties.

    • Mendeleev’s arrangement allowed predictions of unknown elements like Gallium (31Ga).

    • Modern periodic table organizes elements by increasing atomic number (proton number).

Structure of the Modern Periodic Table

  • Horizontal Arrays (Periods): Represent energy levels of elements.

  • Vertical Arrays (Groups/Families): Group elements with similar properties.

  • 118 Elements: Classified into groups based on different references:

    • Metals (left), Nonmetals (right), Metalloids (middle).

Classifications of Elements

Main Group Elements (s and p blocks)

  • Group 1: Alkali Metals

    • Highly reactive, +1 charge, example: Lithium, Sodium, Potassium.

  • Group 2: Alkaline Earth Metals

    • Slightly reactive, +2 charge, example: Beryllium, Magnesium.

  • Group 13: Boron Group (Triels)

    • +3 charge, weakly reactive, example: Aluminum, Gallium.

  • Group 14: Carbon Group (Tetrels)

    • Either +4 or -4 charge, example: Carbon, Silicon.

  • Group 15: Nitrogen Group (Pnictogens)

    • -3 charge, example: Nitrogen, Phosphorus.

  • Group 16: Oxygen Group (Chalcogens)

    • -2 charge, example: Oxygen, Sulfur.

  • Group 17: Halogens

    • -1 charge, highly reactive, example: Fluorine, Chlorine.

  • Group 18: Noble Gases

    • Stable, zero charge, example: Helium, Neon.

Transition Metals

  • Located in the d block with partially or fully filled d orbitals.

  • Multiple oxidation states, example: Vanadium (charges +2, +3, +4, +5).

Inner Transition Metals

  • Contain f block elements, includes Lanthanides and Actinides.

  • Lanthanides: Mostly non-radioactive with +4 max oxidation state.

  • Actinides: Mostly radioactive with +6 max oxidation state.

Electronic Configuration Notations

  • Electron Configuration: Represents electron distribution across orbitals (e.g., 4p5).

  • Valence Electrons: Key in bonding; can be determined by electron configuration.

Trends in Atomic Properties

  1. Atomic Radius

    • Decreases across periods, increases down groups.

  2. Metallicity

    • Decreases across periods, increases down groups.

  3. Ionization Energy

    • Increases across periods, decreases down groups.

  4. Electronegativity

    • Increases across periods, decreases down groups.

  5. Electron Affinity

    • Increases across periods, decreases down groups.

Conclusion

  • Understanding periodicity helps in recognizing trends and predicting properties of elements based on their positions in the periodic table.