Periodic Trends in Chemistry

1. Definition and Importance of Periodicity

• Periodicity: recurrent patterns observed across periods and groups in the periodic table.
• Physical Properties:
  • Influence boiling points, melting points, and atomic/ionic sizes.
  • Facilitate predictions about density and thermal conductivity.
• Chemical Properties:
  • Affect reactivity, electron affinity, and ionisation energy.
  • Impact elements' stability and bonding capabilities.

---

2. Arrangement of the Periodic Table

• Overview:
  • The table is organised into periods (rows) and groups (columns) to reflect elemental properties.
  • This layout aids in predicting trends in properties across various elements.
• Historical Context:
  • Dmitri Mendeleev: Created the initial structured arrangement by atomic weight, forming the foundation for the modern periodic framework.
  • Did you know? Mendeleev's insight has profoundly shaped our understanding of element prediction.

---

3. Electronic Configurations and State of Matter

• Electronic Configurations:
  • Denotes the distribution of electrons around an atom's nucleus.
  • Determines an element's state at room temperature.
  • Example: Sodium (Na): $1s^2 2s^2 2p^6 3s^1$
• States of Matter:
  • Solid: Atoms are closely packed, e.g., Iron (Fe).
  • Liquid: Atoms allow fluid movement, e.g., Mercury (Hg).
  • Gas: Atoms are widely separated with high energy, e.g., Chlorine (Cl).

Visual Illustrations and Examples

• Diagram: Depicting states of matter for metals, non-metals, and metalloids.

---

4. Key Trends in Groups and Periods

• Groups:
  • Similar chemical properties arise from having the same number of valence electrons.
• Periods:
  • Properties consistently evolve with increasing atomic numbers.
• Example:
  • Group 2: Notable for acid-base interactions.
  • Group 17: Reactivity varies due to changes in electronegativity.

---

5. Predicting Chemical Behaviour

• Predictive Power:
  • Trends facilitate predictions concerning element reactivity and bonding.

Worked Example

• Comparing Lithium and Potassium Reactivity:
  1. Both elements are in Group 1 (alkali metals)
  2. Potassium is below lithium in the group
  3. As we move down Group 1, reactivity increases because:
     • Atomic radius increases (more electron shells)
     • Valence electrons are further from the nucleus
     • Less energy is required to remove the outer electron
  4. Therefore, potassium reacts more vigorously with water than lithium

---

6. Atomic Radii and Ionisation Energy Trends

Atomic Radii

• Trend Across Periods:
  • Decreases: Greater nuclear charge draws electrons closer.
• Trend Down Groups:
  • Increases: Adding electron shells results in larger atomic size.

Ionisation Energy

• Across a Period:
  • Generally increases.
• Down a Group:
  • Generally decreases.

---

7. Electronegativity

• Definition: The strength of an atom's pull for electrons within a bond.
• Trend Across a Period: Rises from left to right.
• Trend Down a Group: Diminishes.

---

8. Reactivity with Water

Alkali Metals (Group 1)

• Trend: Reactivity increases down the group.
• Reactions:
  • Lithium: $2\mathrm{Li(s)} + 2\mathrm{H}_2\mathrm{O(l)} \rightarrow 2\mathrm{LiOH(aq)} + \mathrm{H}_2\mathrm{(g)}$
  • Sodium: $2\mathrm{Na(s)} + 2\mathrm{H}_2\mathrm{O(l)} \rightarrow 2\mathrm{NaOH(aq)} + \mathrm{H}_2\mathrm{(g)}$

Alkaline Earth Metals (Group 2)

• Trend: Less vigorous reactions.
• Reaction Example:
  • Calcium: $\mathrm{Ca(s)} + 2\mathrm{H}_2\mathrm{O(l)} \rightarrow \mathrm{Ca(OH)}_2\mathrm{(aq)} + \mathrm{H}_2\mathrm{(g)}$

Safety Considerations

---

9. Key Takeaways

• Key Concepts:
  • Understanding Periodicity: Essential for forecasting element behaviour and chemical interactions.
  • Applications: Crucial in industrial and pharmaceutical innovations.