A mole is a unit that tells us how many particles (atoms, molecules, ions, etc.) are in a sample.
One mole contains exactly the same number of particles as there are atoms in 12 grams of pure carbon‑12.
This number is called Avogadro's constant:
\$N_A = 6.022\times10^{23}\,\text{mol}^{-1}\$
Think of it like a shoebox full of marbles. If you fill the box with exactly 6.022×10²³ marbles, that box holds one mole of marbles.
The same idea applies to any substance, no matter what it is.
The mass of one mole of a substance is called its molar mass.
It is expressed in grams per mole (g mol⁻¹) and is numerically equal to the average atomic or molecular weight.
| Compound | Molar Mass (g mol⁻¹) |
|---|---|
| H₂O | 18.02 |
| CO₂ | 44.01 |
| NaCl | 58.44 |
| C₆H₁₂O₆ | 180.16 |
At standard temperature and pressure (STP: 0 °C, 1 atm), one mole of any ideal gas occupies
\$22.4\;\text{L}\$.
This is called the molar volume.
Example: If you have 2 mol of nitrogen gas, the volume at STP is
\$2 \times 22.4 = 44.8\;\text{L}\$.
In a balanced chemical equation, the coefficients give the mole ratio between reactants and products.
For example, in the combustion of methane:
\$CH4 + 2\,O2 \rightarrow CO2 + 2\,H2O\$
The ratio of methane to oxygen is 1 : 2.
This means that for every 1 mol of methane, 2 mol of oxygen are required.
📌 Remember the value of Avogadro's constant: \$6.022\times10^{23}\,\text{mol}^{-1}\$.
📌 When converting between mass and moles, use the molar mass:
\$n = \frac{m}{M}\$
📌 For gases at STP, use the molar volume (22.4 L mol⁻¹).
If the temperature or pressure differs, apply the ideal gas law: \$PV = nRT\$.
📌 Practice converting between grams, moles, and particles.
The key is to keep the units consistent and use the appropriate conversion factor.