Define the term rate of reaction.

The rate of reaction is the change in concentration of reactants or products per unit time.

One controlled variable for this investigation is the mass/surface area of magnesium carbonate, temperature, or pressure.

The mass of the reaction mixture and flask decreases because carbon dioxide gas ($\text{CO}_2(\text{g})$) is produced during the reaction and escapes from the flask.

From the graph for curve B:
Initial mass of flask + contents at $t=0 \text{ s}$ is $144.50 \text{ g}$.
Final mass of flask + contents at $t=120 \text{ s}$ is $143.24 \text{ g}$.

Mass of $\text{CO}_2$ produced:

Molar mass of $\text{CO}_2$:

Moles of $\text{CO}_2$ produced:

Volume of $\text{CO}_2$ produced (molar gas volume is $24.5 \text{ dm}^3 \cdot \text{mol}^{-1}$):

Average rate of $\text{CO}_2$ production:

Rounding to three significant figures:
The average rate is $\boxed{\text{0.00585 dm}^3 \cdot \text{s}^{-1}}$.

Experiment 1 has the lowest $\text{HCl}$ concentration ($0.1 \text{ mol} \cdot \text{dm}^{-3}$). A lower concentration means a slower reaction rate, which corresponds to the curve with the least steep slope (slowest mass decrease).
Therefore, curve A represents Experiment 1.
According to collision theory, a lower concentration of $\text{HCl}$ means there are fewer $\text{HCl}$ particles per unit volume. This leads to a decrease in the frequency of effective collisions between $\text{HCl}$ and $\text{MgCO}_3$ particles, resulting in a slower reaction rate.

The final mass of $\text{CO}_2(\text{g})$ produced in Experiment 2 will be THE SAME as that in Experiment 3.
This is because the problem states that hydrochloric acid is in excess. This means that magnesium carbonate ($\text{MgCO}_3$) is the limiting reactant in both experiments. Since the initial amount of the limiting reactant ($\text{MgCO}_3$) is the same for all experiments, the total amount of product ($\text{CO}_2$) formed will be the same, regardless of the $\text{HCl}$ concentration.