Theme 4: Periodic properties of the elements (Lecture Slides Summary)

Theme 4: Periodic Properties of the Elements

• Overview of periodic properties of elements.

Importance of Chemistry in Art (Page 2)

• Impressionist oil paintings, such as Monet's, rely on chemistry.

• Inorganic salts suspended in organic media (hydrocarbons).

• Painters utilize a wide range of compounds from the periodic table.

• Periodic table organization is based on electron configurations.

  • Elements in the same column have similar valence electron arrangements.

  • Examples: Oxygen (O: [He]2s2 2p4) and Sulfur (S: [Ne]3s2 3p4).

Historical Development of the Periodic Table (Page 3)

• Discovery of elements has evolved since ancient times.

  • Stable elements mostly found as compounds, not in elemental form.

  • 19th-century advancements allowed isolation of elements from their compounds.

• Classification of elements became essential as discoveries grew.

Mendeleev and Meyer (Page 4)

• Dmitri Mendeleev (Russia) and Lothar Meyer (Germany) developed similar classification systems.

• Recognized periodicity in chemical properties by arranging elements by increasing atomic weight.

• Mendeleev's predictions for missing elements like gallium (Ga) and germanium (Ge) were later confirmed.

Henry Moseley and Atomic Numbers (Page 5)

• Moseley established atomic numbers in 1913 based on unique X-ray frequencies produced by elements.

• Identified atomic number as the number of protons, clarifying periodic table organization.

• Addressed discrepancies in atomic weights for accurate arrangement in the table.

Periodicity: Valence Electron Configurations (Page 6-7)

• Periodic table organizes elements by increasing atomic number.

• Similarities in properties arise from valence electron configurations:

  • Alkali metals (valence configuration ns1).

  • Alkaline earth metals (ns2).

  • p-block elements (ns2npx).

Effective Nuclear Charge (Page 8-9)

• Electrons are attracted to the nucleus (positive charge) while being repelled by each other.

• Coulomb's law applies: attraction depends on charge magnitude and distance.

• Effective nuclear charge (Zeff) accounts for electron shielding by core electrons.

  • Formula: Zeff = Z - S.

  • S represents the number of core electrons affecting Zeff.

• Increases going down and to the right. The size of the orbitals increases so the electrons are further apart.

Atomic Size and Ionic Size Trends (Page 13-17)

• Atomic size increases down a column due to higher principal quantum numbers (n).

• Size decreases across a row due to increasing effective nuclear charge, drawing electrons closer to the nucleus.

• Cations are smaller than their parent atoms, while anions are larger due to increased repulsion among added electrons.

• Isoelectronic series show ions with the same electron count but differing nuclear charges.

Ionization Energy (Page 24-27)

• Ionization energy is the energy required to remove an electron from an atom.

  • Trends:

    • Increases across a period (left to right).

    • Decreases down a group.

• Successive ionization energies increase as electrons are removed, showcasing a steep increase when reaching core electrons.

Electron Affinity (Page 37-38)

• Electron affinity reflects the energy change when an electron is added to an atom.

• Negative values indicate energy release; positive values indicate instability of formed anions (e.g., noble gases).

• Halogens have the most negative affinities due to their tendency to form stable negative ions.

  • Noble gases have positive affinities due to unfavorable conditions for adding electrons.

• Trends are less consistent compared to ionization energy across p-block and s-block elements.