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In 1792, German chemist Jeremias Benjamin Richter noticed that chemicals combine in fixed, predictable weight proportions—a discovery for which he coined the term stoichiometry. If you burn an unmeasured ribbon of magnesium in air, how can a basic balance scale reveal the exact balanced equation without any guessing?
To find the balanced reaction, we convert laboratory masses into moles—the standard SI counting unit where 1 mole=6.022×10231\text{ mole} = 6.022 \times 10^{23} particles, calculated using n=mMn = \frac{m}{M}.
A stoichiometric coefficient is the whole-number multiplier placed in front of a formula in a balanced equation that represents the exact molar ratio of reacting particles.
How do we apply this mole ratio method to real laboratory combustion data?

Deducing Coefficients from Experimental Masses

Suppose 4.86 g4.86\text{ g} of magnesium metal reacts completely with 3.20 g3.20\text{ g} of oxygen gas (O2\text{O}_2) to produce 8.06 g8.06\text{ g} of white magnesium oxide powder.
First, we divide each substance's mass (mm) by its molar mass (MM), which is the mass in grams per one mole of that substance (MMg=24.31 g/molM_{\text{Mg}} = 24.31\text{ g/mol}, MO2=32.00 g/molM_{\text{O}_2} = 32.00\text{ g/mol}, MMgO=40.31 g/molM_{\text{MgO}} = 40.31\text{ g/mol}).
Dividing all molar amounts by the smallest value (0.100 mol0.100\text{ mol}) gives an integer ratio of 2:1:22 : 1 : 2, directly yielding the equation 2Mg+O2→2MgO2\text{Mg} + \text{O}_2 \rightarrow 2\text{MgO}.
What happens if one of the product masses escapes as a gas, or the resulting mole ratio contains a fraction?

Handling Missing Masses and Non-Integer Ratios

Consider heating 12.26 g12.26\text{ g} of solid potassium chlorate (KClO3\text{KClO}_3, M=122.55 g/molM = 122.55\text{ g/mol}) in a crucible, leaving behind 7.46 g7.46\text{ g} of solid potassium chloride (KCl\text{KCl}, M=74.55 g/molM = 74.55\text{ g/mol}) after gas escapes.