Conservation of mass

Light a candle and watch it burn down through the evening — it looks like matter is vanishing into thin air. It isn't. The wax is combining with oxygen from the air and leaving as invisible gases and heat, and if you could trap every atom in a sealed jar and weigh it, the reading would not budge. This is the Law of Conservation of Mass: in any physical or chemical change, matter is neither created nor destroyed, only rearranged into new forms. It is one of the oldest and most reliable laws in chemistry, first demonstrated carefully by Antoine Lavoisier using sealed vessels and a balance. This article states the law precisely, shows the closed-system experiment that proves it, and works through the two situations — combustion and gas release — where students most often "lose" mass on paper that was never really lost in the flask.

The core insight

In a closed system, the total mass of the reactants always equals the total mass of the products, because matter is only rearranged in a physical or chemical change, never created or destroyed.

The core idea

The law, precisely. In any physical or chemical change occurring in a closed system — one where nothing enters or escapes — the total mass before the change equals the total mass after it.

\text{Total mass of reactants (including any gases involved)} = \text{Total mass of products}

where: "reactants" means every substance that goes into the change, including any gas drawn in from the surrounding air, and "products" means every substance that comes out, including any gas released.

Why "closed system" is the whole game. The law is exact, but a careless measurement can make it look false. If a container is open and a gas escapes during the reaction, the flask will weigh less afterward — not because mass was destroyed, but because some of it left unmeasured. If a gas is drawn in from the air and combines with the reactant, the solid left behind can weigh more than what you started with, because an invisible ingredient joined in. Both cases obey the law perfectly; only the bookkeeping was incomplete. To test the law properly, you must either seal the system completely or carefully account for every gas that enters or leaves.

Two classic traps to recognise on sight. In combustion, oxygen gas from the air is a real reactant even though nobody weighs it beforehand — a burnt sample can look heavier than the original solid. In gas-releasing reactions, such as a carbonate reacting with an acid to release carbon dioxide, an open flask feels lighter afterward only because the escaped gas was never weighed. Neither case breaks the law; both simply remind you that "mass of the solid on the balance" is not the same as "mass of the whole system."

Figure

A stoppered flask before and after a reaction shows the same balance reading on both sides, while an open flask releasing gas shows a lighter reading afterward

Derivation AApplying the closed-system mass balance to combustion

The goal is to show, using nothing more than subtraction, that a burnt solid appearing heavier than it started is not a violation of the Law of Conservation of Mass but a sign that a gas reactant from the air must be included in the total, and to find exactly how much of that gas took part.

Step 1. Write the mass balance as an equation. For any closed change, mass of reactants equals mass of products: $m_{\text{solid}} + m_{\text{gas in}} = m_{\text{product}}$ , where $m_{\text{solid}}$ is the starting weighable solid, $m_{\text{gas in}}$ is the mass of any gas from the air that combined with it, and $m_{\text{product}}$ is the final substance weighed on the balance.

Step 2. Identify what was actually measured. In a combustion experiment you can weigh $m_{\text{solid}}$ before burning and $m_{\text{product}}$ after, but you cannot weigh $m_{\text{gas in}}$ directly, since it came from the surrounding atmosphere and was never placed on the balance.

Step 3. Rearrange to solve for the missing gas mass: $m_{\text{gas in}} = m_{\text{product}} - m_{\text{solid}}$ . This single subtraction recovers exactly how much oxygen must have joined the reaction, even though it was invisible to the balance beforehand.

Step 4. Check the balance holds. Substituting the recovered $m_{\text{gas in}}$ back into Step 1 always gives an exact equality, because it was derived from that very equation — the "extra" mass the student sees is not created mass, it is unweighed reactant mass finally accounted for.

Worked example

Question: A student burns a 5 g strip of magnesium ribbon completely in air. The magnesium oxide formed weighs 8.3 g. The student claims this proves conservation of mass is wrong, since only 5 g was weighed at the start. Explain whether the student is correct, and find the mass of oxygen that reacted.

Step 1. List what was actually measured: the magnesium ribbon before burning, $m_{\text{solid}} = 5\ \text{g}$ , and the magnesium oxide product after burning, $m_{\text{product}} = 8.3\ \text{g}$ .

Step 2. Recognise the hidden reactant. When magnesium burns, it does not turn into magnesium oxide on its own — it combines chemically with oxygen gas drawn from the surrounding air. That oxygen is a genuine reactant, simply one that was never placed on the balance.

Step 3. Apply the mass balance from the derivation: $m_{\text{gas in}} = m_{\text{product}} - m_{\text{solid}} = 8.3\ \text{g} - 5\ \text{g} = 3.3\ \text{g}$ of oxygen combined with the magnesium.

Step 4. Check the full balance: $5\ \text{g (magnesium)} + 3.3\ \text{g (oxygen from air)} = 8.3\ \text{g (magnesium oxide)}$ — reactant mass exactly equals product mass. The student is wrong: the law still holds; 3.3 g of atmospheric oxygen was the missing reactant, not created mass.

Competitive stretch

If the same student instead burned a carbonate mineral, driving off carbon dioxide gas in an open crucible, the product left behind would weigh less than the starting mineral — by exactly the mass of the escaped $\text{CO}_2$ . The mass balance is identical in structure: $m_{\text{mineral}} = m_{\text{solid product}} + m_{\text{gas out}}$ , so $m_{\text{gas out}} = m_{\text{mineral}} - m_{\text{solid product}}$ . The direction of the "surprise" flips — losing mass instead of gaining it — but the reasoning is the same subtraction, just with the unweighed substance leaving instead of entering.

Common traps

Weighing only the visible solid and ignoring a gas reactant from the air. In combustion, atmospheric oxygen is a real reactant and must be added to the total, even though it was never on the balance beforehand.
Weighing an open container and expecting exact agreement. If a gas escapes, the measured mass drops even though nothing was truly destroyed — the law holds exactly only when the system is closed or the escaping gas is accounted for.
Mistaking a dramatic appearance for a change in mass. A precipitate forming or a solid turning to ash looks like a big transformation, but total mass is conserved whenever nothing physically enters or leaves.
Confusing conservation of mass with conservation of volume. The law concerns mass only; the volume of a gas or liquid product can differ hugely from the volume of the starting reactants.

Ask the tutor: Want to try more combustion or gas-release mass problems, or see how this law sets up balancing chemical equations later on?

If mass is always conserved, why does a burning candle seem to disappear?

Because the candle is not vanishing, it is combining with oxygen from the air and leaving as gases such as carbon dioxide and water vapour, plus heat and light. In an open room those gases simply diffuse away unweighed, so the visible wax shrinks while the total mass — wax plus oxygen consumed, equal to gases produced — stays exactly the same. If you burned the candle inside a large sealed container and weighed everything, including the trapped gases, the reading before and after would match.

Why does the barium chloride and sodium hydrogen sulphate experiment need a sealed flask?

Because the whole point is to prove the law holds exactly, with no gas sneaking in or escaping to blur the result. The two solutions sit in separate tubes inside one stoppered flask; when tilted, they mix and react to form a solid precipitate, but nothing crosses the stopper in either direction. Weighing before and after then gives a true test: if the readings match, as they do, it is direct evidence that mass is conserved, not just an assumption.

Does conservation of mass mean the number of substances stays the same too?

No — the law is only about total mass, not about how many substances or how much volume you end up with. A single reactant can break into several products, or several reactants can combine into one product; solids can turn into gases and gases into solids. What never changes is the sum of all the masses involved, provided nothing enters or leaves the system. Counting substances or measuring volume tells you nothing about whether mass was conserved.

Conservation of mass

The core idea

Worked example

Competitive stretch

Common traps

Before this

Where this leads