What the law says
Law of conservation of mass: in a chemical reaction, the total mass of the reactants equals the total mass of the products. Mass is not created and not destroyed.
- Reactants are the substances we start with.
- Products are the new substances that form.
So: mass of reactants = mass of products.
The French chemist Antoine Lavoisier showed this in the 1770s and 1780s. He did reactions in sealed glass vessels and weighed them very carefully before and after. The mass did not change. The Russian scientist Mikhail Lomonosov had made a similar finding a few years before.
Closed and open systems: why the balance sometimes changes
A closed system lets nothing in or out (a sealed flask). An open system lets gases move in or out (an open beaker).
Mass seems to go down
Marble chips in acid give off carbon dioxide gas. In an open beaker the gas escapes, so the balance shows less. The 'missing' mass is the gas in the air.
Mass seems to go up
When magnesium ribbon or steel wool burns, it joins with oxygen from the air. The product (an oxide) is heavier than the metal. The 'extra' mass is the oxygen that joined.
Rule: if you count every substance, including gases, the total never changes. To test the law, use a closed system.
Why mass is conserved: atoms are only rearranged
In a reaction, bonds between atoms break and new bonds form. But no atom is made, destroyed or changed into another kind. Each atom has a fixed mass. So if the same atoms are there before and after, the total mass must be the same.
Example: 2 H₂ + O₂ → 2 H₂O. Left: 4 hydrogen atoms + 2 oxygen atoms. Right: 4 hydrogen atoms + 2 oxygen atoms.
This is why we balance chemical equations: the number of each kind of atom must be the same on both sides. John Dalton used the conservation of mass as evidence for his idea that matter is made of tiny atoms that cannot be destroyed in chemical reactions.
Law of definite proportions
Law of definite (constant) proportions (Joseph Proust, about 1799): a pure compound always contains the same elements in the same ratio by mass, wherever it comes from.
- Water from a river, rain or a lab: always 1 g hydrogen to 8 g oxygen (1 : 8).
- Carbon dioxide: always 3 g carbon to 8 g oxygen (3 : 8).
If you mix elements in a different ratio, the extra amount of one element is simply left over; it does not join. Example: 3 g hydrogen + 16 g oxygen → 18 g water + 1 g hydrogen left over. Total before 19 g, total after 19 g.
Mass conservation in cycles of nature and industry
Because atoms are never lost, elements move round in cycles.
- Carbon cycle: plants take carbon dioxide from the air to make sugars; animals eat plants and breathe carbon dioxide out; burning fuel also returns carbon to the air.
- Water cycle: water evaporates, forms clouds, falls as rain and flows back to the sea.
- Industry and recycling: a factory must account for every kilogram. If 100 kg of reactants go in, 100 kg of products and waste come out. Recycling metals like aluminium keeps the same atoms in use instead of mining new ore.
Chemists use a mass balance (mass in = mass out) to design processes and to find where waste goes.
Limits of the law
In chemistry the law works perfectly for all practical purposes. In nuclear reactions (like in the Sun or a nuclear power station), a tiny amount of mass changes into a large amount of energy (E = mc²). There, mass alone is not conserved, but mass and energy together are. For school chemistry, use: mass of reactants = mass of products.
Key formulas and definitions
- Total mass of reactants = Total mass of products
- Mass of gas given off = mass before − mass after (open container)
- Mass of gas taken in = mass after − mass before (e.g. burning a metal)
- Water: hydrogen : oxygen = 1 : 8 by mass
- Carbon dioxide: carbon : oxygen = 3 : 8 by mass
Worked examples
1. 12 g of magnesium burns in oxygen and makes 20 g of magnesium oxide. How much oxygen joined?
Mass of reactants = mass of products. 12 g + oxygen = 20 g, so oxygen = 20 − 12 = 8 g.
2. In an open beaker, 50.0 g of acid and 10.0 g of marble chips react. After the fizzing stops the beaker contents weigh 55.6 g. What mass of carbon dioxide escaped?
Before: 50.0 + 10.0 = 60.0 g. After: 55.6 g. Gas escaped = 60.0 − 55.6 = 4.4 g.
3. Heating 100 g of calcium carbonate gives 56 g of calcium oxide and some carbon dioxide. Find the mass of carbon dioxide.
100 g = 56 g + CO₂, so CO₂ = 100 − 56 = 44 g.
4. How much water forms from 4 g of hydrogen and 40 g of oxygen? What is left over?
Ratio H : O = 1 : 8. 4 g hydrogen needs 4 × 8 = 32 g oxygen. Water = 4 + 32 = 36 g. Oxygen left = 40 − 32 = 8 g. Check: 44 g before, 36 + 8 = 44 g after.
5. Show that 2 H₂ + O₂ → 2 H₂O obeys the law by counting atoms.
Left: 2 × 2 = 4 H, 1 × 2 = 2 O. Right: 2 × 2 = 4 H, 2 × 1 = 2 O. Same atoms, same numbers, so the same mass.
6. A student finds 6 g of carbon gives 22 g of carbon dioxide. Another sample gives 33 g of carbon dioxide. How much carbon was in it?
Carbon : CO₂ = 6 : 22 = 3 : 11. For 33 g CO₂: carbon = 33 × 3 ÷ 11 = 9 g. Same ratio, as the law of definite proportions says.
Common mistakes
- Thinking mass is 'lost' when a gas escapes. It is still there, in the air.
- Thinking burning metal gains mass from nowhere. The extra mass is oxygen from the air.
- Testing the law in an open container and then saying the law is wrong. Use a closed container.
- Adding all the masses given, even the part that is left over unreacted, into the product. Only the amounts in the fixed ratio react; the rest stays as it was.