Atomic Structure
🔹 Atomic Structure and Empty Space
- Atoms consist of a central nucleus surrounded by electron shells.
- The nucleus contains protons and neutrons:
- It is extremely small compared to the whole atom (approximately 1/100,000th of atom’s diameter).
- It contains most of the atom’s mass.
- Electrons occupy the region outside the nucleus, arranged in shells (also called energy levels).
- The atom is mostly empty space:
- If the nucleus were the size of a marble, the nearest electron would be several meters away.
- This explains how alpha particles could pass through atoms in the Rutherford scattering experiment.
🔹 Subatomic Particles
| Particle | Relative Charge | Relative Mass |
|---|---|---|
| Proton | +1 | 1 |
| Neutron | 0 | 1 |
| Electron | -1 | 1/1836 |
- Protons:
- Positively charged.
- Found in the nucleus.
- Number of protons = atomic number = determines the element.
- Neutrons:
- No charge (neutral).
- Also found in the nucleus.
- Contribute to mass, but not to charge.
- Electrons:
- Negatively charged.
- Occupy orbitals in energy levels (shells).
- In a neutral atom, number of electrons = number of protons.
🔹 Key Atomic Terms and Definitions
- Atomic number (Z):
- Also called proton number.
- The number of protons in the nucleus.
- Determines the identity of the element.
- Also equals number of electrons in a neutral atom.
- Mass number (A):
- Also called nucleon number.
- Total number of protons and neutrons in the nucleus.
- A = Z + number of neutrons.
- Isotopes:
- Atoms of the same element (same proton number) with different numbers of neutrons.
- Same chemical properties, different physical properties (e.g. density, rate of diffusion).
- Example:
- Carbon-12: 6 protons, 6 neutrons.
- Carbon-13: 6 protons, 7 neutrons.
🔹 Distribution of Mass and Charge in the Atom
- Mass:
- Almost entirely concentrated in the nucleus (protons + neutrons).
- Electrons have negligible mass compared to nucleons.
- Charge:
- Positive charge is located in the nucleus (due to protons).
- Electrons provide negative charge in surrounding shells.
- In a neutral atom, the total charge is zero: number of protons = number of electrons.
🔹 Behaviour of Protons, Neutrons, and Electrons in Electric Fields
- Proton Beam:
- Positively charged → deflects towards the negative plate.
- Deflection is moderate, due to its relatively large mass and +1 charge.
- Electron Beam:
- Negatively charged → deflects towards the positive plate.
- Deflection is large, due to very small mass and -1 charge.
- Neutron Beam:
- No charge → no deflection in an electric field.
- Continues in a straight line.
- Factors affecting deflection:
- Charge-to-mass ratio: Greater the ratio, greater the deflection.
- Velocity of particles: Slower particles deflect more for a given field strength.
🔹 Calculating Subatomic Particles in Atoms and Ions
Given: Atomic number (Z), mass number (A), and charge:
- Number of protons = atomic number (Z).
- Number of neutrons = A – Z.
- Number of electrons:
- For neutral atom: same as number of protons.
- For cations (+ve ions): electrons = Z – charge.
- For anions (–ve ions): electrons = Z + charge.
Examples:
- Sodium atom (Na):
- Z = 11, A = 23
- Protons = 11
- Neutrons = 23 – 11 = 12
- Electrons = 11 (neutral atom)
- Sodium ion (Na⁺):
- Protons = 11
- Electrons = 11 – 1 = 10
- Chloride ion (Cl⁻):
- Z = 17, A = 35
- Protons = 17
- Neutrons = 18
- Electrons = 17 + 1 = 18
🔹 Atomic Radius and Ionic Radius: Trends and Explanations
🔸 Atomic Radius
- Defined as the distance from the nucleus to the outermost electron.
- Measured using X-ray diffraction or calculated from covalent bond lengths.
Trend Across a Period (Left to Right):
- Atomic radius decreases across a period.
- Reasons:
- Nuclear charge increases (more protons).
- Electrons added to same energy level.
- Increased electrostatic attraction pulls electrons closer.
- No increase in shielding → greater effective nuclear charge.
Example:
- Atomic radius from Na to Cl decreases:
- Na: 186 pm
- Mg: 160 pm
- Al: 143 pm
- Si: 118 pm
- P: 110 pm
- S: 104 pm
- Cl: 99 pm
Trend Down a Group:
- Atomic radius increases down a group.
- Reasons:
- New electron shells are added → increased distance from nucleus.
- Increased shielding from inner electrons.
- Effective nuclear attraction on outer electrons is reduced.
Example:
- Atomic radius from F to I increases:
- F: 71 pm
- Cl: 99 pm
- Br: 114 pm
- I: 133 pm
🔸 Ionic Radius
- Radius of an atom’s ion in a crystal lattice or in a gas phase.
Cations (Positive Ions):
- Smaller than parent atoms.
- Reasons:
- Loss of electrons → fewer electron-electron repulsions.
- Remaining electrons pulled closer due to unchanged nuclear charge.
- Often lose an entire shell.
Example:
- Na (neutral) = 186 pm
- Na⁺ = 102 pm
Anions (Negative Ions):
- Larger than parent atoms.
- Reasons:
- Gain of electrons → increased electron-electron repulsion.
- Expansion of electron cloud.
- Nuclear charge spread over more electrons → reduced pull per electron.
Example:
- Cl (neutral) = 99 pm
- Cl⁻ = 181 pm
🔹 Additional Notes and Advanced Understanding
- Mass spectrometry uses ion deflection in magnetic/electric fields to determine isotopic masses.
- The concept of effective nuclear charge (Z_eff) explains periodic trends:
- Z_eff = Z – shielding effect.
- Across a period: Z increases, shielding nearly constant → Z_eff increases → atomic size decreases.
- Down a group: shielding increases → Z_eff roughly constant or decreases → atomic size increases.
- Shielding effect:
- Inner electrons repel outer electrons.
- Reduces the net attractive force between nucleus and valence electrons.
🔹 Real-World Applications and Connections
- Rutherford’s gold foil experiment:
- Provided evidence for the nuclear model.
- Showed that most alpha particles passed through, but some were deflected sharply.
- Periodic Table organization is based on atomic number (Z), not mass number.
- Isotopes play important roles in:
- Medical imaging (e.g. Technetium-99m)
- Carbon dating (Carbon-14)
- Nuclear power and weapons (Uranium-235 vs. Uranium-238)
- Ion size impacts lattice energies, solubility, melting points of ionic compounds.