Mass and Equations

This section explains mass and equations, covering conservation of mass, relative formula mass and mass changes when the reactant or product is a gas. 

Conservation of Mass

In chemistry, the law of conservation of mass states that mass is neither created nor destroyed in a chemical reaction. This means that the total mass of reactants before the reaction will be the same as the total mass of the products after the reaction, provided the system is closed.

When performing calculations in quantitative chemistry, it's important to know how to convert between masses and amounts of substances involved in reactions. This typically involves working with the relative atomic mass (Ar) and the relative formula mass (Mr).

Relative Formula Mass (Mr)

The relative formula mass (Mr) of a compound is the sum of the relative atomic masses (Ar) of all the atoms in its formula.

For example:

Water (H₂O): 

• Hydrogen (H) has an atomic mass of 1, and oxygen (O) has an atomic mass of 16.
• Mr of H₂O = (2 × 1) + (1 × 16) = 18 g/mol.

To calculate the mass of a substance from the number of moles, you can use the following formula:

$$\text{Mass} = \text{Number of moles} \times \text{Relative formula mass}$$ 

Mass Changes when a Reactant or Product is a Gas

When a gas is involved in a chemical reaction, the mass of the system may appear to change if the gas is not contained. If a gas is a product of a reaction, the mass of the product can be difficult to measure because it might escape into the atmosphere. On the other hand, if a gas is a reactant, the mass of the system may decrease as the gas is consumed in the reaction.

For example, if a solid reactant is heated and produces a gas, the mass of the system will decrease if the gas escapes. However, if the gas is collected and measured, the total mass remains constant.

Example: Copper Carbonate Heating

Consider the thermal decomposition of copper carbonate:

$$\text{Copper carbonate (CuCO₃)} \xrightarrow{\text{heat}} \text{Copper oxide (CuO)} + \text{Carbon dioxide (CO₂)}$$

• Before heating: Copper carbonate is a solid. The mass of the copper carbonate can be weighed.
• During heating: Copper carbonate decomposes to form copper oxide (solid) and carbon dioxide (gas). The carbon dioxide gas escapes, so if the system is not sealed, the mass of the system will decrease.
• After heating: The remaining solid is copper oxide. If you were to collect and measure the mass of the carbon dioxide gas, you would find that the total mass remains the same, even though the gas escaped during the reaction.

Example Calculation:

If you start with 10 g of copper carbonate and heat it to produce copper oxide and carbon dioxide, you might find that 8 g of copper oxide is formed and 2 g of carbon dioxide escapes.

Balanced Equation:

$$\text{CuCO₃ (s)} \rightarrow \text{CuO (s)} + \text{CO₂ (g)}$$

Another Example: Magnesium Reacting with Hydrochloric Acid

Let’s consider another example of a reaction where we can apply quantitative chemistry. When magnesium reacts with hydrochloric acid, magnesium chloride and hydrogen gas are produced:

$$\text{Mg (s)} + 2\text{HCl (aq)} \rightarrow \text{MgCl₂ (aq)} + \text{H₂ (g)}$$

Steps:

1. Write the balanced equation.
2. Calculate the molar masses:

• Magnesium (Mg): 24 g/mol
• Hydrochloric acid (HCl): 36.5 g/mol
• Magnesium chloride (MgCl₂): 95 g/mol
• Hydrogen gas (H₂): 2 g/mol

3. Use the mole ratio to determine the amount of products formed from a given mass of reactants.

For example, if 12 g of magnesium reacts with excess hydrochloric acid, the number of moles of magnesium can be calculated as:

$$\text{Number of moles of Mg} = \frac{12 \, \text{g}}{24 \, \text{g/mol}} = 0.5 \, \text{mol}$$ 

From the balanced equation, we know that 1 mole of magnesium reacts to form 1 mole of hydrogen gas. Therefore, 0.5 mol of magnesium will produce 0.5 mol of hydrogen gas.

$$\text{Mass of H₂} = 0.5 \, \text{mol} \times 2 \, \text{g/mol} = 1 \, \text{g}$$ 

Through these examples, you can see how the mass of reactants and products can be calculated using the mole concept and balanced chemical equations. This is crucial in understanding how mass is conserved and how to perform quantitative calculations in chemistry.