4: Stoichiometry

Stoichiometry: Quantifying Chemical Change

Stoichiometry, derived from Greek words meaning "element measuring," is the mathematical backbone of chemistry. It quantifies relationships between reactants and products in chemical reactions based on the law of conservation of mass and the mole concept.

The Mole and Avogadro’s Number

The mole (mol) is the SI unit for amount of substance. One mole contains Avogadro’s number ($6.022 × 10^{23}$) of particles (atoms, molecules, ions). This links microscopic entities to measurable masses. The molar mass (g/mol) of any element or compound equals its atomic/molecular mass in grams. For example:

• Carbon (C): 12.01 g/mol
• Water (H₂O): 18.02 g/mol (2×1.008 + 16.00).

Balancing Chemical Equations

Chemical equations must be balanced to obey mass conservation. Adjust coefficients (whole numbers preceding formulas) to ensure atom counts match on both sides. For example:
$\text{Unbalanced: } \ce{CH4 + O2 -> CO2 + H2O}$
$\text{Balanced: } \ce{CH4 + 2O2 -> CO2 + 2H2O}$

Stoichiometric Calculations

Balanced equations provide mole ratios (conversion factors between substances). For:
$\ce{N2 + 3H2 -> 2NH3}$
The ratio $\ce{H2 : N2}$ is 3:1, meaning 3 moles of H₂ react with 1 mole of N₂.

Mass-to-mass conversions involve three steps:

1. Convert given mass → moles (using molar mass).
2. Convert moles of Substance A → moles of Substance B (using mole ratio).
3. Convert moles of B → mass (using molar mass).

Limiting Reactant and Yield

In reactions with multiple reactants, the limiting reactant is consumed first and dictates the maximum product (theoretical yield). Identify it by:

1. Calculating moles of each reactant.
2. Using mole ratios to determine which reactant produces less product.

Percent yield compares actual experimental yield to theoretical yield:
$\text{Percent Yield} = \left( \frac{\text{Actual Yield}}{\text{Theoretical Yield}} \right) \times 100\%$
Low percent yield indicates side reactions, incomplete purification, or measurement errors.