Evolution of Atomic Theory

Revision of Atomic Theory

• Dalton's Atomic Theory: Proposed that all matter is composed of indivisible and indestructible atoms.
• Thomson's Plum Pudding Model: Suggested atoms are spheres of positive charge with negative electrons embedded throughout.
• Rutherford's Gold Foil Experiment: Led to the nuclear model with a small, dense, positively charged nucleus surrounded by mostly empty space containing electrons.
• Bohr's Model: Introduced quantized energy levels for electrons orbiting the nucleus, explaining atomic spectra.
• Quantum Mechanical Model: Describes electrons in terms of probability distributions (orbitals) and quantum numbers, incorporating wave-particle duality.
• Atomic Number (Z): Number of protons in the nucleus, defining the element.
• Mass Number (A): Total number of protons and neutrons in the nucleus.
• Isotopes: Atoms of the same element with different numbers of neutrons, hence different mass numbers.
• Atomic Mass Unit (amu): Defined as 1/12 the mass of a carbon-12 atom.
• Relative Atomic Mass (Ar): Weighted average of the masses of isotopes of an element, relative to the atomic mass unit.

Periodic Table and Trends (Atomic Radius)

• Atomic Radius: Typically refers to the distance from the nucleus to the outermost electron orbital.
• Measured in picometers (pm) or angstroms (Å).
• Trends in Atomic Radius:
  • Down a Group: Atomic radius generally increases due to the addition of electron shells.
  • Across a Period: Atomic radius generally decreases due to increasing nuclear charge (effective nuclear charge) attracting electrons more strongly.
• Effective Nuclear Charge: The net positive charge experienced by an electron in a multi-electron atom.
• Shielding Effect: Inner electrons shield outer electrons from the full nuclear charge.
• Factors Affecting Atomic Radius:
  • Number of electron shells
  • Nuclear charge
  • Shielding effect
• Metallic Radius: Half the distance between the nuclei of adjacent atoms in a solid metal.
• Covalent Radius: Half the distance between the nuclei of two atoms joined by a single covalent bond.
• Van der Waals Radius: Half the distance between the nuclei of two non-bonded atoms in a solid.

Periodic Table and Trends (Ionisation Energy & Electron Affinity)

• Ionization Energy (IE): The energy required to remove one mole of electrons from one mole of gaseous atoms in their ground state.
• Measured in kJ/mol.
• Trends in Ionization Energy:
  • Down a Group: IE generally decreases because the outermost electrons are farther from the nucleus and more shielded.
  • Across a Period: IE generally increases because of increasing nuclear charge and decreasing atomic radius.
• Successive Ionization Energies: The energy required to remove subsequent electrons.
• Core Electrons: Electrons in inner shells require significantly more energy to remove due to higher nuclear attraction.
• Electron Affinity (EA): The change in energy when one mole of electrons is added to one mole of gaseous atoms to form one mole of gaseous ions.
• Can be either exothermic (negative EA) or endothermic (positive EA).
• Trends in Electron Affinity:
  • Down a Group: EA generally decreases (becomes less negative) due to increasing atomic size and shielding.
  • Across a Period: EA generally increases (becomes more negative) due to increasing nuclear charge.
• Exceptions to EA trends:
  • Group 2 elements have EA close to zero due to stable filled s-subshells.
  • Group 15 elements have lower EA than expected due to half-filled p-subshells.
• Factors Affecting IE and EA:
  • Nuclear charge
  • Atomic radius
  • Shielding effect
  • Electron configuration
• Relationship to Reactivity:
  • Low IE indicates a tendency to lose electrons (metals).
  • High EA indicates a tendency to gain electrons (nonmetals).

Periodic Table/Trends

• Periodic Law: The properties of elements are periodic functions of their atomic numbers.
• Groups (Vertical Columns): Elements in the same group have similar chemical properties due to the same number of valence electrons.
• Periods (Horizontal Rows): Elements in the same period have different properties that change gradually across the row.
• Metals: Typically lustrous, conductive, malleable, and ductile; tend to lose electrons.
• Nonmetals: Typically dull, non-conductive, and brittle; tend to gain electrons.
• Metalloids (Semimetals): Have properties intermediate between metals and nonmetals.
• Alkali Metals (Group 1): Highly reactive, readily lose one electron to form +1 ions.
• Alkaline Earth Metals (Group 2): Reactive, readily lose two electrons to form +2 ions.
• Halogens (Group 17): Highly reactive, readily gain one electron to form -1 ions.
• Noble Gases (Group 18): Generally inert due to stable, filled valence shells.
• Transition Metals (Groups 3-12): Exhibit variable oxidation states and form colored compounds.
• Lanthanides and Actinides: Inner transition metals with f-electrons.
• Trends in Metallic Character:
  • Down a Group: Metallic character increases.
  • Across a Period: Metallic character decreases.
• Trends in Electronegativity:
  • Electronegativity: A measure of the ability of an atom in a chemical compound to attract electrons.
  • Down a Group: Electronegativity generally decreases.
  • Across a Period: Electronegativity generally increases.
• Trends in Acidity/Basicity of Oxides:
  • Across a Period: Oxides trend from basic (left) to acidic (right).
  • Down a Group: For nonmetals, acidity of oxides increases.
• Diagonal Relationships: Similarities between elements diagonally adjacent to each other in the periodic table (e.g., Li and Mg, Be and Al, B and Si).
• Electron Configuration and the Periodic Table: The periodic table is organized based on electron configurations, with blocks corresponding to the filling of s, p, d, and f orbitals.