Lecture on the Periodic Table by Professor David

Introduction

• The periodic table is a well-known symbol in chemistry.
• Elements appear randomly arranged, but the arrangement reveals nature's patterns.
• Mid-1800s, chemists sought ways to table elements; Dmitri Mendeleev's arrangement succeeded due to data correlation and predictive power.

Mendeleev's Periodic Table

• Arranged elements into rows (periods) and columns (groups).
• Grouped elements with similar behavior together.
• Predicted the existence and properties of not-yet-discovered elements.
• Currently organized into metals, metalloids, and nonmetals.
• Similar group behavior due to having the same number of valence electrons.

Example: Group 1 Elements

• All have one valence electron (one in outermost shell).
• As you move down the table:
  • Number of shells increases (n increases).
  • Outer shell's valence remains one electron.
• Group 2 Elements: Two electrons in the outermost shell.
• Number of valence electrons determines many characteristics.

Periodic Trends

1. Atomic Radius

• Size of an atom.
• Increases downward (adding shells).
• Decreases to the right (one more proton per element increasing electromagnetic attraction, thus shrinking radius).

2. Ionic Radius

• Different from atomic radius.
• Electrons repel each other:
  • Adding an electron increases the atom's size.
  • Removing an electron decreases size.
• Ions with the same electron configuration: Radii decrease as atomic number increases.

3. Ionization Energy

• Energy required to remove an electron (from outermost shell).
• Electromagnetic force drops quickly with distance.
• Farther electrons from the nucleus are easier to remove.
• Trend is opposite to atomic radius.
  • Large atoms (like Francium) ionize easily.
  • Small atoms (like Helium) require more energy to ionize.
• Atoms prefer full outer shells; it's easier for Group 1 to lose one electron.
• Helium has stable, full shells requiring more energy for ionization.

Successive Ionization Energies

• Removing more electrons requires more energy.
• Huge energy jump after removing an electron from a full shell.
• Orbital symmetry can explain deviations:
  • Example: Oxygen has lower ionization energy than Nitrogen due to stability reasons.

4. Electron Affinity

• Opposite of ionization energy.
• Energy released when an atom gains an electron.
• Increases to the right (disregarding noble gases).
  • Fluorine has high affinity (gain one electron to complete shell).
  • Elements in the lower left prefer losing electrons.
• Deviations explained by orbital symmetry.

5. Electronegativity

• Ability of an atom to hold onto electrons.
• Increases to the right (disregarding noble gases).
• Smaller atoms with more protons (like Fluorine) hold electrons tightly.

Recap

• Key Trends:
  • Atomic radius increases downward.
  • Ionization energy, electron affinity, and electronegativity increase to the right.

Conclusion

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