Unit 4: Classification of Elements and Periodic Table

Modern Periodic Law: Properties of elements are periodic functions of their atomic numbers. When elements are arranged by increasing atomic number, their chemical and physical properties show periodic trends.
Modern Periodic Table: The table arranges elements based on atomic number into 7 periods (horizontal rows) and 18 groups (vertical columns). Elements in the same group have similar chemical properties due to the same number of valence electrons.

Groups:
- There are 18 groups in the periodic table.
- Elements in the same group have the same number of valence electrons and exhibit similar chemical properties.
- Example: Group 1 (Alkali metals), Group 17 (Halogens), Group 18 (Noble gases).
Periods:
- There are 7 periods.
- Each period corresponds to the number of electron shells an element has.
- Across a period, elements transition from metals to non-metals.
Blocks:
- Elements are categorized into blocks based on their electron configuration.
  - s-block: Groups 1 and 2 (Alkali and Alkaline earth metals).
  - p-block: Groups 13 to 18 (includes metalloids, non-metals, and halogens).
  - d-block: Transition metals (Groups 3 to 12).
  - f-block: Lanthanides and actinides, often shown below the main table.

The International Union of Pure and Applied Chemistry (IUPAC) follows the modern periodic law to classify elements.
Elements are assigned a unique atomic number and symbol, following standard nomenclature rules. The periodic table is divided into main groups (s and p blocks), transition metals (d block), and inner transition metals (f block).

Nuclear Charge: The total charge of all the protons in the nucleus. As the atomic number increases, the nuclear charge increases.
Effective Nuclear Charge (Z_eff): The net positive charge experienced by an electron in an atom. It is lower than the actual nuclear charge due to the shielding effect of inner electrons.
$Z_{\text{eff}} = Z - \sigma$
Where $Z$ is the atomic number and $\sigma$ is the shielding constant.

Definition: The distance from the nucleus to the outermost electron.
Trend:
- Across a period: Atomic radius decreases due to increasing nuclear charge, pulling electrons closer to the nucleus.
- Down a group: Atomic radius increases because additional electron shells are added.

Definition: The radius of an atom's ion.
Trend:
- Cations (positive ions): Smaller than their neutral atoms due to the loss of electrons and decreased electron-electron repulsion.
- Anions (negative ions): Larger than their neutral atoms due to the gain of electrons and increased electron-electron repulsion.

Definition: The energy required to remove an electron from a gaseous atom.
Trend:
- Across a period: Increases due to increased nuclear charge, making it harder to remove an electron.
- Down a group: Decreases because the outer electrons are farther from the nucleus and are easier to remove.

Definition: The energy change that occurs when an atom gains an electron.
Trend:
- Across a period: Generally increases as atoms more readily accept electrons to complete their valence shells.
- Down a group: Generally decreases because the added electron is further from the nucleus.

Definition: The ability of an atom to attract electrons in a chemical bond.
Trend:
- Across a period: Increases because of increasing nuclear charge.
- Down a group: Decreases because the atomic radius increases, making it harder for the nucleus to attract bonding electrons.

Definition: Refers to the tendency of an element to lose electrons and form positive ions (cations).
Trend:
- Across a period: Decreases as elements become less likely to lose electrons (non-metals).
- Down a group: Increases because atoms are more willing to lose electrons due to the larger atomic radius and lower ionization energy.

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