Combustion analysis is a laboratory method used to determine the empirical formula of an organic compound that contains carbon, hydrogen, and sometimes oxygen. A known mass of the unknown sample is burned completely in excess oxygen, converting its carbon into CO2 and its hydrogen into H2O. By measuring the masses of CO2 and H2O produced, chemists can work backward to find how much carbon and hydrogen were in the original sample.
This matters because empirical formulas reveal the simplest whole-number ratio of atoms in a compound.
Understanding Chemistry: Combustion Analysis
A combustion analysis setup has to do more than make a sample burn. The gases formed must travel through a system that captures them separately and completely. A drying agent can absorb water vapor.
Another absorber can trap carbon dioxide. Each absorber is weighed before and after the experiment. Its increase in mass shows how much product it collected.
This is why leaks are a serious problem. Escaping gas makes the measured product mass too small.
Air entering the apparatus can cause a different error because air already contains water vapor and carbon dioxide. Good experiments use dry equipment, a steady oxygen flow, and careful weighing.
The calculation works because formulas describe fixed atom counts. Carbon dioxide has one carbon atom per molecule, so its amount tells chemists the amount of carbon that came from the sample. Water gives the hydrogen amount after accounting for its two hydrogen atoms.
Students often need to make several unit changes in the right order. Product mass becomes moles of product using molar mass. Those moles give moles of each element.
Element moles can then be changed into element masses when needed. Keeping units beside every number helps catch mistakes. A value in grams cannot be added directly to a value in moles.
Oxygen needs extra care when it is part of the unknown compound. It is usually not measured in a separate trap. Instead, its mass is found from what remains after the carbon and hydrogen masses are removed from the original sample mass.
This method makes oxygen more sensitive to error. If either collected product mass is slightly wrong, the calculated oxygen amount changes too.
The result can even become negative, which is a warning that a measurement, assumption, or calculation is wrong. Nitrogen, sulfur, and halogens require different methods because the simple mass difference approach no longer identifies oxygen alone.
The final whole-number ratio is not always obvious. After dividing all element mole amounts by the smallest amount, the results may be close to values such as one point five or one point three three. These values suggest multiplying every ratio by two or three before rounding.
Rounding too early is one of the most common errors. An empirical formula gives a ratio, not necessarily the actual number of atoms in a molecule.
For example, a compound with a ratio of carbon to hydrogen of one to two could have a molecular formula with one carbon and two hydrogen atoms, or a larger formula with the same ratio. A separate molar mass measurement is needed to distinguish those possibilities.
Key Facts
- Complete combustion converts carbon to CO2 and hydrogen to H2O.
- mol C = mol CO2 because each CO2 molecule contains 1 carbon atom.
- mol H = 2 × mol H2O because each H2O molecule contains 2 hydrogen atoms.
- mass C = mol C × 12.01 g/mol and mass H = mol H × 1.008 g/mol.
- If the compound contains oxygen, mass O = sample mass - mass C - mass H.
- Empirical formula steps: grams to moles, divide by the smallest mole value, multiply to get whole numbers.
Vocabulary
- Combustion analysis
- A method for determining the elemental composition of a compound by burning it and measuring the CO2 and H2O produced.
- Empirical formula
- The simplest whole-number ratio of atoms of each element in a compound.
- Complete combustion
- A reaction in which a substance burns in enough oxygen to convert carbon fully to CO2 and hydrogen fully to H2O.
- Mole ratio
- A ratio comparing the amounts of substances in moles, often taken from a chemical formula or balanced equation.
- Mass percent
- The percentage by mass of a particular element in a compound.
Common Mistakes to Avoid
- Using the mass of CO2 as the mass of carbon is wrong because CO2 also contains oxygen. Convert CO2 to moles, then use the 1 to 1 mole ratio to find moles of carbon.
- Using the mass of H2O as the mass of hydrogen is wrong because water also contains oxygen. Convert H2O to moles, then multiply by 2 to get moles of hydrogen atoms.
- Assuming oxygen comes directly from the measured gases is wrong when the compound itself may contain oxygen. Find oxygen by subtracting the calculated masses of carbon and hydrogen from the original sample mass.
- Rounding mole ratios too early can give an incorrect empirical formula. Keep several significant figures until the final whole-number ratio step.
Practice Questions
- 1 A 0.500 g sample of a compound containing C, H, and O is burned completely to produce 0.733 g CO2 and 0.300 g H2O. Find the empirical formula.
- 2 A hydrocarbon sample with mass 1.20 g produces 3.76 g CO2 and 1.54 g H2O during complete combustion. Determine its empirical formula.
- 3 In a combustion analysis experiment, why is it important that the sample burns completely in excess oxygen, and how would incomplete combustion affect the calculated empirical formula?