Mass Changes in Reactions

Conservation of Mass in Reactions

In any chemical reaction, the total mass of the reactants is always equal to the total mass of the products. This is known as the conservation of mass. It means that atoms are neither created nor destroyed during a chemical reaction; they are simply rearranged to form new substances.

This principle applies only in a closed system, where no substances can enter or leave during the reaction. If the system is open, gases or other substances might escape or enter, causing apparent changes in mass.

For example, when magnesium ribbon burns in air, it reacts with oxygen to form magnesium oxide. If this happens in a sealed container, the mass before and after the reaction stays the same, showing conservation of mass.

Mass Changes with Gases

When gases are involved in reactions, the mass of the substances can appear to change if the system is not closed. This happens because gases can enter or leave the reaction mixture.

• Mass increases if a gas from the surroundings reacts and becomes part of the product. For example, when iron rusts, it reacts with oxygen from the air, so the mass of the iron increases as it forms iron oxide.
• Mass decreases if a gas is produced and escapes from the reaction mixture. For example, when a metal carbonate reacts with acid, carbon dioxide gas is released, causing the mass of the reaction mixture to decrease if the gas escapes.

These changes occur because the system is not closed, and gas particles can move in or out.

Observing Mass Changes

To accurately observe mass changes during reactions involving gases, experiments must be carried out in sealed containers or using equipment that prevents gases from escaping or entering.

For example, using a sealed flask or a closed reaction vessel ensures that all reactants and products remain inside, so the total mass remains constant and can be measured accurately.

Accurate measurement of mass before and after the reaction is essential. This involves using a balance that measures to at least two decimal places and ensuring the apparatus is stable and free from vibrations. Precise balances help detect small mass changes accurately, improving experimental reliability.

Applications of Mass Changes

Understanding mass changes helps predict the mass of products formed in reactions. By knowing the mass of reactants and which gases are involved, chemists can calculate expected masses of products.

A common real-life example is the rusting of iron. Iron reacts with oxygen and water from the air to form iron oxide (rust), increasing the mass of the iron object. This explains why rusty iron weighs more than the original iron.

Chemical equations can be used to relate the masses of reactants and products, helping to predict outcomes in industrial processes and laboratory experiments.

Example: Calculating Mass Increase When Iron Rusts

Iron reacts with oxygen to form iron oxide:

$\text{4Fe} + 3\text{O}_2 \rightarrow 2\text{Fe}_2\text{O}_3$

If 56 g of iron reacts completely with oxygen, what is the mass of oxygen that reacted?

First, find the relative formula masses (Mr). Mr is the sum of the relative atomic masses of all atoms in a formula unit.

• Fe = 56
• O = 16
• $\text{Fe}_2\text{O}_3 = (2 \times 56) + (3 \times 16) = 112 + 48 = 160$

From the equation, 4 moles of Fe (4 × 56 = 224 g) react with 3 moles of $\text{O}_2$ (3 × 32 = 96 g) to form 2 moles of $\text{Fe}_2\text{O}_3$ (2 × 160 = 320 g).

If 56 g of Fe reacts (which is 1 mole), then the mass of $\text{O}_2$ reacting is:

$\frac{96}{224} \times 56 = 24 \text{ g}$

So, the mass of oxygen that reacts is 24 g, and the total mass of iron oxide formed is $56 + 24 = 80 \text{ g}$.