The rate of reaction between magnesium ribbon and dilute hydrochloric acid at room temperature is investigated - Edexcel - GCSE Chemistry - Question 9 - 2018 - Paper 1

To calculate the rate of reaction from the tangent drawn on the graph in Figure 12, we first need to determine the volume of hydrogen evolved at that point (let's assume it is 60 cm³ at 40 seconds).

Next, calculate the change in volume over the change in time:

$ext{Rate of reaction} = \frac{\text{change in volume}}{\text{change in time}} = \frac{60 - 45}{40 - 30} = \frac{15}{10} = 1.5 \, \text{cm}^3 \, \text{s}^{-1}$

The graph would depict a steeper curve to the left of the printed curve and would eventually merge back down. This indicates an increased reaction rate due to the larger surface area of powdered magnesium.

To calculate the number of moles, we use the formula:

$\text{Number of moles} = \frac{\text{mass}}{\text{molar mass}}$

Substituting the values:

$\text{Molar mass of Mg} = 24 \, \text{g/mol}$

$\text{Number of moles of Mg} = \frac{0.1 \, \text{g}}{24 \, \text{g/mol}} = \frac{0.1}{24} = 0.00417 \, \text{moles}$

The balanced equation shows that 2 moles of HCl reacts with 1 mole of Mg. If we have 0.5 mol of HCl, we will need:

$\frac{0.5 \, \text{moles of HCl}}{2} = 0.25 \, \text{moles of Mg}$

Since we only have 0.5 mol of Mg, the magnesium is in excess.

The experiments illustrate that increasing the concentration of A or the temperature will increase the rate of reaction. With higher temperatures, particles move faster, increasing collision frequency. Similarly, higher concentrations of A lead to more particles in a given volume, causing more frequent successful collisions, thereby speeding up the reaction.