Lecture on Chemistry: Periodic Table and Atomic Structure

Introduction to Chemistry and Elements

• Robert Boyle:

  • Known as the father of chemistry.
  • Defined an element in 1661.
  • An element is a substance that cannot be split into simpler substances by chemical means.

• Humphry Davy:

  • Discovered many elements by passing electricity through compounds.
  • Early 19th century marked efforts to discover more elements.

Development of the Periodic Table

• Early Attempts:

  • Johann de Bruyne observed that some elements had similar properties.
  • Introduced the concept of triads: groups of three elements with similar properties.

• John Newlands (1860s):

  • Proposed Newlands Octaves: properties repeat every eighth element.
  • Faced issues due to undiscovered elements and incorrect assumptions.

• Dmitri Mendeleev:

  • Created the periodic table by arranging elements by increasing atomic weight, leaving gaps for undiscovered elements.
  • Made predictions about elements and properties, some elements were reversed for properties to align.

• Henry Moseley (1913):

  • Discovered the atomic number using X-rays, leading to the modern periodic table arranged by atomic number.

Modern Periodic Table

• Elements are arranged by increasing atomic number.
• Properties of elements reoccur periodically.

Atomic Structure

• Atomic Number:

  • Defined as the number of protons in the nucleus of an atom.

• Ions:

  • Atoms or groups of atoms with a charge; ions have unequal numbers of protons and electrons.

• Subatomic Particles in Ions:

  • To determine the number of subatomic particles, calculate based on electron loss or gain.

Isotopes and Mass Spectrometry

• Isotopes:

  • Atoms of the same element with different mass numbers due to varying neutrons.

• Mass Spectrometer:

  • Instrument for discovering isotopes and determining atomic masses.
  • Five processes: Vaporisation, Ionisation, Acceleration, Separation, Detection.

Electron Configuration

• Sub-levels:

  • Includes s, p, d (for leaving certificate, only these are needed).
  • S: 1 orbital, max 2 electrons.
  • P: 3 orbitals, max 6 electrons.
  • D: 5 orbitals, max 10 electrons.

• Examples:

  • Fluorine: 9 electrons, configuration gets to 2p5.
  • Phosphorus: 15 electrons, configuration reaches 3p3.
  • Magnesium Ion: Has lost 2 electrons, configuration gets to 2p6.

Exceptions in Electron Configuration

• Chromium and Copper:
  • Electrons rearrange for stability to achieve half-filled or completely filled sublevels.

Principles and Rules

• Hund’s Rule:

  • Electrons occupy orbitals singly before filling them in pairs.

• Pauli Exclusion Principle:

  • No more than two electrons per orbital, with opposite spins.

Conclusion

• Understanding of periodic table arrangement and atomic structure is vital.
• Emphasis on modern periodic law, electron configurations, and foundational chemistry principles.