The Law of Conservation of Mass
The law of conservation of mass is one of the most fundamental principles in all of science. It states that in a chemical reaction, the total mass of the reactants (starting materials) equals the total mass of the products (substances formed). Matter is neither created nor destroyed — it is only rearranged.
Historical Development
The law is most closely associated with the French chemist Antoine-Laurent Lavoisier (1743–1794), often called the "father of modern chemistry." Through carefully controlled and precisely measured experiments, Lavoisier demonstrated that the total mass of substances in a sealed container does not change during a chemical reaction.
Before Lavoisier, chemists struggled to explain combustion. The dominant theory was phlogiston — a hypothetical substance thought to be released by materials when they burned. The problem was that some metals increase in mass when they burn (form rust or calx). Phlogiston theorists awkwardly explained this by claiming phlogiston had "negative weight." Lavoisier's quantitative experiments demolished phlogiston theory:
- He showed that when a metal burns in a closed container, the mass increase of the metal oxide exactly equals the mass decrease of the air inside (specifically, oxygen being consumed).
- His classic 1789 textbook, Traité Élémentaire de Chimie, formalized the principle: "Nothing is created in the operations of art or of nature, and it can be taken as an axiom that in every operation there is an equal quantity of matter before and after."
Atomic Explanation
Conservation of mass follows directly from atomic theory. In a chemical reaction: - Chemical bonds are broken and new bonds are formed. - Atoms are rearranged into new combinations (different molecules). - No atoms are created, destroyed, or converted into other types of atoms.
Because each atom has a definite mass, and the number of each type of atom is the same before and after the reaction, the total mass must be conserved.
Example — Combustion of hydrogen: 2H₂ + O₂ → 2H₂O
Count atoms: - Reactants: 4 H atoms + 2 O atoms - Products: 4 H atoms + 2 O atoms ✓
Mass check: - Reactants: 2(2.016) + 1(32.00) = 4.032 + 32.00 = 36.03 g - Products: 2(18.015) = 36.03 g ✓
Balanced Chemical Equations
The law of conservation of mass is the reason chemical equations must be balanced — the same number and type of each atom must appear on both sides of the arrow.
Unbalanced: H₂ + O₂ → H₂O - Left: 2 H, 2 O - Right: 2 H, 1 O ← imbalanced!
Balanced: 2H₂ + O₂ → 2H₂O - Left: 4 H, 2 O - Right: 4 H, 2 O ✓
Balancing equations requires adjusting coefficients (the numbers in front of formulas) — never changing the subscripts, which would change the identity of the substance.
Real-World Examples
Rusting of iron: 4Fe + 3O₂ → 2Fe₂O₃
When iron rusts, it gains oxygen from the air. A sample of iron left outdoors gets heavier as it rusts — not because mass was created, but because oxygen from the atmosphere is incorporated into the product (iron(III) oxide, Fe₂O₃). If you measured the iron and the oxygen consumed, their combined mass would equal the mass of the rust formed.
Burning wood: When wood burns, the solid material seems to almost disappear, leaving only a small pile of ash. Has mass been lost? No — the apparent disappearance is because most of the products are gases (CO₂, H₂O vapor, small amounts of CO and soot) that escape into the air. If the reaction were conducted in a sealed container that captured all gases, the total mass inside would be unchanged.
Photosynthesis: 6CO₂ + 6H₂O + light energy → C₆H₁₂O₆ + 6O₂
Plants build glucose from CO₂ and H₂O. The mass of glucose and oxygen produced equals the mass of CO₂ and water consumed — only the form has changed.
Conservation of Mass vs. Conservation of Energy
Conservation of mass and conservation of energy are distinct classical principles, but Einstein's famous equation E = mc² revealed they are actually unified: mass and energy can be interconverted.
In nuclear reactions (fission, fusion), a tiny amount of mass is converted into enormous amounts of energy. The Sun converts about 4 million tons of mass into energy every second through nuclear fusion.
However, for all ordinary chemical reactions, the mass change due to energy exchange (ΔE/c²) is immeasurably small — less than 1 part in 10¹⁰. For practical chemistry, the law of conservation of mass holds to extraordinary precision.
Applications
Stoichiometry — calculating the amounts of reactants and products in chemical reactions — is entirely based on conservation of mass. Every time a chemist asks "how much product will I get from this many grams of starting material?", they are applying Lavoisier's law.
Industrial process control: Chemical plants use mass balances — accounting for all input and output materials — to monitor reaction efficiency, detect leaks, and optimize yield. If mass is unaccounted for, either a reaction is incomplete, a by-product is being formed, or material is being lost.
Environmental monitoring: Conservation of mass underlies calculations of pollutant discharge, nutrient cycling, and carbon budgets. The global carbon cycle tracks carbon atoms as they move through atmosphere, oceans, living organisms, and rocks.
Forensic chemistry: When investigators analyze a crime scene, conservation of mass guides the identification of substances (e.g., calculating expected yields from precursor chemicals found at a scene).
Demonstrating Conservation of Mass
A classic laboratory demonstration: dissolve sodium sulfate (Na₂SO₄) and barium chloride (BaCl₂) in separate beakers, weigh the total system, then mix the solutions. A white precipitate of barium sulfate (BaSO₄) forms immediately. Reweighing the combined system shows no change in mass — confirming Lavoisier's law.
Na₂SO₄(aq) + BaCl₂(aq) → BaSO₄(s) + 2NaCl(aq)