Atomic and Molecular Masses
Dalton’s Atomic Theory (1808)
John Dalton proposed that:
- ✨ Matter is made of indivisible atoms.
- ✨ Atoms of the same element have identical mass; atoms of different elements have different masses.
- ✨ Compounds form when atoms combine in fixed ratios.
- ✨ Atoms are rearranged in reactions but never created/destroyed.
Limitation: Couldn’t explain why gases combine in whole-number volume ratios.
Atomic Mass ⚖️
Atoms are super tiny! We measure their mass in atomic mass units (u):
- 1 u = mass of one carbon-12 atom ÷ 12
- 1 u = \(1.66056 \times 10^{-24}\) g
Example: Hydrogen atom mass = \(1.6736 \times 10^{-24}\) g ≈ 1.008 u.
Average Atomic Mass
Most elements have multiple isotopes (atoms with different masses). Average mass is calculated using their natural abundance. Example for carbon:
| Isotope | Abundance | Mass |
|---|---|---|
| \(^{12}\text{C}\) | 98.892% | 12 u |
| \(^{13}\text{C}\) | 1.108% | 13.00335 u |
Average mass = \((0.98892 \times 12) + (0.01108 \times 13.00335)\) = 12.011 u.
Molecular Mass 🧪
Sum of atomic masses of all atoms in a molecule:
- Methane (\(\text{CH}_4\)): \[ \text{C} + 4\text{H} = 12.011\, \text{u} + 4 \times 1.008\, \text{u} = \bf{16.043\, \text{u}} \]
- Water (\(\text{H}_2\text{O}\)): \[ 2 \times 1.008\, \text{u} + 16.00\, \text{u} = \bf{18.02\, \text{u}} \]
Formula Mass 🔷
For ionic compounds (like NaCl) that don’t form molecules, we use formula mass:
- Sodium chloride (\(\text{NaCl}\)): \[ \text{Na} + \text{Cl} = 23.0\, \text{u} + 35.5\, \text{u} = \bf{58.5\, \text{u}} \]
Mole Concept 🥑
1 mole = \(6.022 \times 10^{23}\) particles (atoms, molecules, etc.). This is called Avogadro’s number (\(N_A\)).
- 1 mole of H atoms = \(6.022 \times 10^{23}\) atoms
- 1 mole of \(\text{H}_2\text{O}\) = \(6.022 \times 10^{23}\) molecules
Molar mass = mass of 1 mole of a substance (in grams). It matches atomic/molecular mass in u:
- Molar mass of water = 18.02 g/mol
- Molar mass of NaCl = 58.5 g/mol
Important for NEET! 🚀
- Avogadro’s number (\(6.022 \times 10^{23}\)) links microscopic atoms to measurable quantities.
- Average atomic mass calculations (e.g., carbon isotopes).
- Molecular vs. formula mass (covalent vs. ionic compounds).
- Mole conversions: Use molar mass to switch between grams ↔ moles ↔ particles.
