Atomic Structure

🔹 Particles in the Atom and Atomic Radius

⚛️ Subatomic Particles:

• Atoms are made of protons, neutrons, and electrons.
• Protons: Positive charge (+1), relative mass = 1
• Neutrons: Neutral charge (0), relative mass = 1
• Electrons: Negative charge (–1), relative mass ≈ 1/1836

⚗️ Experimental Discovery:

• J.J. Thomson discovered electrons using the cathode ray tube experiment.
• Rutherford’s Gold Foil Experiment discovered the nucleus:
  → Most alpha particles passed through, but a few deflected.
  → Conclusion: Atoms have a small, dense, positively charged nucleus.
• Millikan’s Oil Drop Experiment: Determined the charge of the electron.

📏 Atomic Radius:

• Atomic radius decreases across a period (↑ nuclear charge, electrons pulled closer).
• Increases down a group (↑ electron shells, more shielding).

🔹 Isotopes, Relative Atomic Mass (RAM), and Relative Molecular Mass (RMM)

🧪 Isotopes:

• Definition: Atoms of the same element with the same number of protons but different numbers of neutrons.
• Example:
  • Hydrogen: ¹H, ²H (deuterium), ³H (tritium)
  • Carbon: ¹²C, ¹³C, ¹⁴C

📊 Relative Atomic Mass (RAM):

Weighted average mass of the isotopes of an element, relative to 1/12 of the mass of a ¹²C atom.

Formula: RAM=∑(isotopic mass×abundance)/100

🧮 Relative Molecular Mass (RMM):

• Sum of the relative atomic masses of all atoms in a molecule.
• E.g., for H₂O:
  RMM = (2 × 1.0) + 16.0 = 18.0

🔹 Electrons, Energy Levels and Atomic Orbitals

🌀 Electron Configuration:

• Electrons occupy orbitals (s, p, d, f) in order of increasing energy.
• Aufbau Principle: Fill lower energy orbitals first.
• Pauli Exclusion Principle: Max 2 electrons per orbital, opposite spins.
• Hund’s Rule: Electrons occupy all orbitals singly before pairing.

📋 Order of Filling:

1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p

🔢 Electron Configurations (Z = 1 to 36):

Examples:

• H (Z = 1): 1s¹
• He (Z = 2): 1s²
• O (Z = 8): 1s² 2s² 2p⁴
• Na (Z = 11): 1s² 2s² 2p⁶ 3s¹
• Cl (Z = 17): 1s² 2s² 2p⁶ 3s² 3p⁵
• Ca (Z = 20): 1s² 2s² 2p⁶ 3s² 3p⁶ 4s²
• Fe (Z = 26): [Ar] 4s² 3d⁶
• Kr (Z = 36): [Ar] 4s² 3d¹⁰ 4p⁶

🌀 Atomic Orbitals:

• s-orbital: spherical (1 per level)
• p-orbitals: dumbbell-shaped (3 per level from n=2)
• d-orbitals: clover-shaped (5 per level from n=3)

🔹 Ionisation Energy

⚡ Definition:

• The first ionisation energy is the energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous ions: X(g)→X+(g)+e−X(g)

📈 Factors Affecting Ionisation Energy:

• Nuclear charge (↑ nuclear charge = ↑ ionisation energy)
• Atomic radius (↑ radius = ↓ ionisation energy)
• Shielding (↑ inner electron shielding = ↓ ionisation energy)

🔍 First Ionisation Energies Across Period 3 (Na to Ar):

• General trend: Increases across the period due to increasing nuclear charge and constant shielding.
• Two exceptions:
  1. Mg → Al: Drop due to Al’s electron entering a higher-energy 3p orbital.
  2. P → S: Drop due to electron pairing in 3p orbital causing repulsion.

The electronic configurations of elements using successive ionisation energy data

1. What is successive ionisation energy?

• The first ionisation energy (IE₁) is the energy required to remove the first electron from a gaseous atom.
• The second ionisation energy (IE₂) is the energy required to remove the second electron, and so on.
• Successive ionisation energies generally increase gradually, but there is a large jump when electrons from a core (inner) shell are removed after all outer (valence) electrons have been lost.

2. How to use the data to deduce electronic configuration:

Step 1: Examine the ionisation energy values.

• Identify where there is a large jump in the successive IE values.
• The electrons before the jump are the valence electrons, and the jump occurs when core electrons start being removed.

Step 2: Count the number of electrons removed before the large jump.

• This number gives the number of valence electrons.

Step 3: Deduce the shell structure.

• Use the number of electrons and the periodic table group to assign the full electron configuration.

Example:

Successive ionisation energies (in kJ/mol) for an element X:

IE₁ = 590
IE₂ = 1145
IE₃ = 4912
IE₄ = 6492
IE₅ = 8780

Step 1: Identify the large jump.

• Jump occurs between IE₂ (1145) and IE₃ (4912) → After removing 2 electrons, the next is much harder → 2 valence electrons.

Step 2: Determine total electrons.

• Element is in Group 2 → 2 valence electrons → atomic number = 4 (Li would be 3, Be is 4)

Step 3: Write electronic configuration:

• X = Be → 1s² 2s²

Tips:

• The first large jump always indicates removal of core electrons.
• The number of electrons removed before the jump = number of valence electrons.
• Compare with the periodic table to deduce the element and write the full configuration.

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