Conservation of Mass¶
Syllabus mapping
CH-11-02 — Quantitative Chemistry. Covers: relating the law of conservation of mass to stoichiometry; a laboratory experiment measuring the mass and volume of gas produced in a reaction; the features of open and closed systems; Aboriginal cooking with coals and enclosed pit ovens as examples of open and closed systems (explicit, examinable content integration point); a practical investigation modelling conservation of mass; solving problems involving mass changes in reactions.
Wording above is paraphrased for teaching use — check the official syllabus before treating any line as verbatim NESA text.
You need to know: what the law of conservation of mass says and why it means equations must balance, the difference between open and closed systems, and how to explain (and calculate) a change in mass you measure during a reaction.
Assumed knowledge check¶
Before this page makes sense, you should be comfortable with:
- Reading and balancing a chemical equation — reactants on the left, products on the right
- Writing chemical formulas
Core¶
The law of conservation of mass¶
In a chemical reaction, mass is not created or destroyed. The total mass of the reactants equals the total mass of the products.
Why? Because a chemical reaction only rearranges atoms — bonds break and new bonds form, but every atom you started with is still there at the end, just attached to different partners. Same atoms → same mass.
This is the reason chemical equations must balance: there must be the same number of each type of atom on both sides.
| Reactants | Products | |
|---|---|---|
| Hydrogen atoms | \(2 \times 2 = 4\) | \(2 \times 2 = 4\) ✓ |
| Oxygen atoms | \(2\) | \(2 \times 1 = 2\) ✓ |
The word stoichiometry (stoy-key-OM-etry) just means the number relationships between the amounts of reactants and products in a balanced equation. The law of conservation of mass is what makes stoichiometry work — and it's the starting point for everything else in Quantitative Chemistry.
Worked example
When \(4.0\ \text{g}\) of hydrogen gas reacts completely with \(32.0\ \text{g}\) of oxygen gas, what mass of water forms?
Mass in = mass out, so:
Open and closed systems¶
Here's the catch: if you put a reaction on a balance, the reading often does change. That's not mass being destroyed — it's mass escaping (or arriving), and it depends on what kind of system you're using.
| Open system | Closed system | |
|---|---|---|
| Can matter get in or out? | Yes — gases can escape, or be taken in from the air | No — sealed |
| Can energy (heat) get in or out? | Yes | Yes |
| Example | An open beaker; a campfire | A sealed flask or bottle; a sealed bag |
| Does the measured mass stay the same? | Not if a gas is made or used up | Yes |
A reaction that makes a gas, in an open beaker: marble chips (calcium carbonate) react with hydrochloric acid to make carbon dioxide gas.
In an open beaker, the \(\ce{CO2}\) bubbles out into the room, so the balance reading goes down. Seal the same reaction in a flask with a balloon over the top and the reading stays the same — the gas is trapped, so its mass is still being weighed.
A reaction that uses a gas, in an open container: when magnesium burns in air, it combines with oxygen from the air.
The white powder left behind weighs more than the magnesium you started with. That extra mass is the oxygen that joined on.
So in an open system, ask yourself: is a gas leaving, or is a gas joining?
Aboriginal and Torres Strait Islander content — placeholder, not yet written
NESA lists Aboriginal cooking processes using coals and enclosed pit ovens as examples of open and closed systems as explicit, examinable syllabus content, not an optional cultural add-on. The specific practices, and whose Knowledge they are, need proper research — ideally with input from your school's Aboriginal Education team, local Community, or a community-verified source — rather than a generic paraphrase.
Deliberately left unwritten here rather than filled with unverified content.
Solving mass-change problems¶
The method is always the same: mass before − mass after = mass of gas that left (or the reverse, if a gas joined).
Worked example — a gas leaves
A flask of hydrochloric acid and marble chips has a mass of \(152.40\ \text{g}\) at the start. When the bubbling stops, the mass is \(151.96\ \text{g}\). What mass of carbon dioxide was produced?
The flask was open, so the \(\ce{CO2}\) escaped. That's the "missing" mass.
Worked example — a gas joins
\(0.24\ \text{g}\) of magnesium ribbon is burned in a crucible. The magnesium oxide produced has a mass of \(0.40\ \text{g}\). What mass of oxygen combined with the magnesium?
Practical¶
Two suggested pracs — both use the same idea, so they pair well:
1. Measuring the mass and volume of gas produced. React a weighed amount of a carbonate (e.g. marble chips or sodium hydrogen carbonate) with excess acid in a conical flask on a balance. Record the mass every 30 seconds until it stops changing. At the same time (or in a repeat run), collect the gas in a gas syringe, or an upturned measuring cylinder full of water, to measure its volume. You'll use the mass and volume together on the Properties of Gases page.
2. Modelling conservation of mass in a closed system. Do a reaction in a sealed container and show the mass doesn't change. Two options:
- Precipitation in a closed flask: put a solution of sodium carbonate in a conical flask and a small tube of calcium chloride solution inside it, stopper the flask and weigh it. Tip the flask so they mix (a white solid forms), then weigh again.
- Gas in a sealed bottle: put an effervescent tablet and some water in a plastic bottle with a balloon stretched over the top. Weigh before and after. Then repeat with the bottle open and compare.
Things worth discussing in your evaluation: was the seal really airtight? Did any liquid splash out? Some \(\ce{CO2}\) dissolves in water, so a gas volume measured over water is usually a little lower than expected.
Advanced
Linking mass changes to moles. The mass change on its own only tells you how much gas left. Once you've done the Mole Concept, you can go further. In the first worked example, \(0.44\ \text{g}\) of \(\ce{CO2}\) is \(\dfrac{0.44}{44.01} = 0.010\ \text{mol}\). From the equation, 1 mol of \(\ce{CaCO3}\) gives 1 mol of \(\ce{CO2}\) — so \(0.010\ \text{mol}\) of marble (\(1.0\ \text{g}\)) must have reacted. That's how real labs use mass loss to measure how much of a substance was in a sample.
Why an "open" result can still prove conservation of mass. In the magnesium example, the mass went up. That seems to break the law — until you account for the oxygen. A good exam answer always names what entered or left the system and why the total mass of everything involved still stayed the same.
Isolated systems. Some textbooks add a third type: an isolated system, where neither matter nor energy can get in or out (a perfect thermos flask is the closest everyday example). Truly isolated systems don't really exist, but the idea matters later for energy changes.
Extension
Beyond the syllabus for this course — won't appear in the HSC, included for interest / depth study inspiration.
Is mass exactly conserved? Einstein's \(E = mc^2\) says that whenever energy is released, a tiny amount of mass disappears with it. In a chemical reaction, this is so small it's impossible to measure — burning a whole kilogram of fuel "loses" well under a millionth of a gram. In nuclear reactions it's much bigger and very measurable; that's where the energy from nuclear power and the Sun comes from. For all chemistry in this course, mass is conserved.
A bit of history. In the 1770s and 1780s, Antoine Lavoisier did careful experiments in sealed containers and showed the total mass never changed — overturning the popular idea that burning released a substance called "phlogiston". His sealed-container approach is basically the closed-system prac above.
📎 Depth study idea
Compare the mass change for the same gas-producing reaction in an open container, a container with a loose lid, and a sealed one, and use the results to measure how much gas escapes — or investigate how temperature affects the volume of gas collected.
Video/visual resources¶
- 🖼️ Diagram needed — side-by-side: (1) an open beaker on a balance with CO₂ bubbles leaving and the balance reading going down; (2) the same reaction in a sealed flask with a balloon on top and the balance reading unchanged. TODO — placeholder until sourced/drawn.
- 🖼️ Diagram needed — gas collection set-ups: a gas syringe attached to a conical flask, and an upturned measuring cylinder in a water trough. TODO — placeholder until sourced/drawn.
- 🎥 TODO: source a short video of the open-vs-closed system demonstration (e.g. effervescent tablet with and without a balloon on a balance).
- 🖥️ PhET Simulation — Balancing Chemical Equations — balance equations and see the atoms counted on each side
Check yourself¶
-
State the law of conservation of mass, and explain why it's true using the idea of atoms.
Answer
In a chemical reaction, the total mass of the reactants equals the total mass of the products. It's true because a reaction only rearranges atoms — no atoms are created or destroyed, so the total mass stays the same.
-
What is the difference between an open system and a closed system?
Answer
In an open system, matter (and energy) can enter or leave — for example, gases can escape. In a closed system, matter can't enter or leave, but energy can.
-
Steel wool is weighed, heated strongly in air until it glows, then weighed again. The mass has increased. Explain.
Answer
The iron reacts with oxygen from the air to form iron oxide. The oxygen atoms that joined the iron add to the mass. Mass isn't created — it came from the air, because the steel wool was in an open system.
-
Sodium hydrogen carbonate is added to vinegar in an open beaker. The mass drops from \(85.62\ \text{g}\) to \(84.96\ \text{g}\). What mass of gas was produced, and what is the gas?
Answer
\(85.62 - 84.96 = 0.66\ \text{g}\). The gas is carbon dioxide — a carbonate (or hydrogen carbonate) reacting with an acid always produces \(\ce{CO2}\).
-
How could you change the set-up in question 4 so that the mass doesn't change?
Answer
Make it a closed system — for example, do the reaction in a sealed bottle or flask with a balloon over the top, so the \(\ce{CO2}\) can't escape.