Why Burning a Candle Does Not Actually Destroy Any Mass

When a candle burns, it seems to disappear — yet if you trapped every gas it gave off, the total mass would not change. Chemistry rests on a few such rules about how substances combine, and on the idea that matter is made of atoms with definite masses.

This lesson covers the laws of chemical combination, Dalton's atomic theory, and atomic, molecular and formula mass.

The laws of chemical combination state that mass is conserved in reactions, that a compound always contains its elements in a fixed proportion by mass, that the masses of one element combining with a fixed mass of another are in simple whole-number ratios, and that reacting gases combine in simple volume ratios.

Law of conservation of mass. Matter is neither created nor destroyed in a chemical reaction.

Worked example. When 2 g of hydrogen reacts completely with 16 g of oxygen, the water formed has a mass of $2+16=18$ g.

Law of definite proportions. A given compound always contains the same elements in the same proportion by mass, whatever its source — pure water from a river or a laboratory always has hydrogen and oxygen in the mass ratio 1 : 8.

Law of multiple proportions. When two elements form more than one compound, the masses of one element that combine with a fixed mass of the other are in a small whole-number ratio.

Worked example. Carbon forms carbon monoxide and carbon dioxide.

- In carbon monoxide, 12 g of carbon combines with 16 g of oxygen
- In carbon dioxide, 12 g of carbon combines with 32 g of oxygen

$$\frac{16\text{ g}}{32\text{ g}}=\frac{1}{2}$$

so the masses of oxygen are in the simple ratio 1 : 2.

Gay Lussac's law of gaseous volumes. At the same temperature and pressure, gases combine or form in simple whole-number volume ratios — 1 volume of hydrogen and 1 volume of chlorine give 2 volumes of hydrogen chloride.

Avogadro's law. Equal volumes of all gases at the same temperature and pressure contain equal numbers of molecules.

An everyday example. A sealed pressure cooker of dal on the stove has the same total mass before and after cooking, as long as no steam escapes.

The substance. Mass seems to vanish only when a gas escapes — burning wood leaves light ash because carbon dioxide and water vapour leave, not because mass is destroyed.

Dalton's atomic theory states that matter consists of indivisible atoms, that atoms of one element are identical in mass and properties, that compounds form when atoms of different elements combine in fixed ratios, and that reactions rearrange atoms without creating or destroying them.

Postulates:

- Matter consists of tiny, indivisible particles called atoms
- All atoms of a given element have identical properties, including identical mass; atoms of different elements differ in mass
- Compounds form when atoms of different elements combine in a fixed ratio
- Chemical reactions reorganise atoms — atoms are neither created nor destroyed

How it explains the laws:

- Atoms are not created or destroyed, so mass is conserved
- A compound always has its atoms in a fixed ratio, so its composition by mass is definite
- Atoms combine only as whole units, so different compounds of the same two elements show whole-number mass ratios

Limitations:

- Atoms are divisible — they contain electrons, protons and neutrons
- Isotopes of one element have different masses
- Isobars — atoms of different elements — can have the same mass
- It cannot explain the law of gaseous volumes or why atoms combine at all

An everyday example. A child building different shapes from the same set of toy blocks mirrors the theory — the blocks are rearranged, never created or destroyed.

The substance. Dalton's theory explains the laws of combination but not the structure of atoms — the existence of subatomic particles shows that atoms can in fact be divided.

Atomic mass is the mass of an atom relative to one-twelfth of a carbon-12 atom, molecular mass is the sum of the atomic masses of all atoms in a molecule, and formula mass is the sum of the atomic masses in one formula unit of an ionic compound, which has no separate molecules.

Atomic mass:

- Measured in atomic mass units (u), where 1 u is one-twelfth of the mass of a carbon-12 atom
- Carbon-12 has a mass of exactly 12 u; hydrogen has about 1.008 u and oxygen about 16.00 u

Average atomic mass. Most elements exist as a mixture of isotopes, so their atomic mass is an average weighted by natural abundance.

Worked example. Chlorine is about 75 per cent chlorine-35 and 25 per cent chlorine-37.

$$\text{Average atomic mass}=35\times 0.75+37\times 0.25=26.25+9.25=35.5\text{ u}$$

Molecular mass. Add the atomic masses of all atoms in the molecule.

Worked example. For glucose, ${\mathrm{C}}_6{\mathrm{H}}_{12}{\mathrm{O}}_6$:

$$6(12.01)+12(1.008)+6(16.00)=72.06+12.10+96.00=180.16\text{ u}$$

Formula mass. Ionic compounds such as sodium chloride exist as a three-dimensional lattice of ions, so one formula unit is used.

Worked example. For sodium chloride, NaCl: $23.0+35.5=58.5$ u.

An everyday example. The table salt on a kitchen shelf contains no NaCl molecules at all — only a lattice of sodium and chloride ions, which is why chemists speak of its formula mass.

The substance. An average atomic mass need not match any single atom — no chlorine atom has a mass of 35.5 u; the value is a weighted average.