Periodic Law, Atomic Structure, and Chemical Trends

Mendeleev's Periodic Law and Atomic Structure

Item 2.1. Mendeleev's periodic law stated that the properties of chemical elements do not vary arbitrarily but in a systematic manner with their atomic mass.

Principal Merits of Mendeleev's Table

• The periodic table demonstrated the periodic properties of chemical elements.
• It gave absolute priority to chemical properties.
• When the properties of elements did not fit, he did not hesitate to leave gaps in the table.
• He correctly forecast the atomic mass and properties that elements occupying the reserved spots in the table should have.

Moseley's Contribution

Moseley determined the wavelengths of X-rays produced by different elements.

3. Different Energy Levels

• Level: Where electrons have the same principal quantum number n. These states are part of the same shell or energy level.
• Sublevel: Each level has sublevels formed by electrons that, besides the principal quantum number, have in common the secondary quantum number (l).
• Orbital: A region where electrons share the same principal and secondary quantum numbers. Some differ by the magnetic quantum number (m). Electrons that share these same quantum numbers are said to belong to the same orbital.

3.2. Pauli Exclusion Principle

In a single atom, there cannot be two electrons with all four quantum numbers being equal. A maximum of two electrons can fit in an orbital.

3.3. Hund's Rule

Electrons are placed singly in orbitals of each sublevel, evenly avoiding occupying the same orbital until necessary.

3.4. Aufbau Principle

The electrons in an atom occupy orbitals of the lowest energy first.

3.5. Madelung's Rule

The energy of an orbital depends on the quantum numbers that define it: n, l, and m.

5. Organization of the Periodic Table

• Period: Consisting of all elements whose highest energy level is the same.
• Group: Integrated by elements that have a similar electronic configuration in their highest energy level.
• Representative Elements: Electronic configuration ending in s or p orbitals:
  • ns¹: Alkali metals
  • ns² np⁵: Halogens
  • ns² np⁴: Chalcogens
  • ns² np³: Nitrogen group (pnictogens)
  • ns² np²: Carbon group
  • ns² np¹: Boron group
• Transition Elements: ns² (n-1)d^x
• Inner Transition Elements: ns² (n-2)f^x

6. Atomic and Ionic Radii

Atomic Radius: Defined as half the distance between the nuclei of two identical atoms joined together.

Ionic Radius: The distance between a positive ion and a negative ion in an ionic solid, which is the sum of the ionic radii of both.

6.2. Ionization Energy

The energy required to remove the last electron from an atom in its gaseous state.

6.3. Electron Affinity

The energy involved when an atom in a gaseous state captures an electron to form a negative ion.

6.4. Electronegativity

A measure of the attraction of an atom for the shared electron pair through which it is bonded to another atom.