Two experiments are carried out to investigate one of the factors that affects the reaction rate between magnesium and dilute hydrochloric acid - NSC Physical Sciences - Question 5 - 2018 - Paper 2

The independent variable in this investigation is the form of magnesium used: ribbon in experiment 1 and powder in experiment 2.

The control variable is the concentration of the hydrochloric acid, which must remain constant across both experiments.

To calculate the volume of hydrogen gas produced, we first determine the moles of hydrogen generated using the balanced equation. The loss of mass over the time period from the graph yields approximately 1 mol of magnesium reacted, producing 1 mol of hydrogen. Using the molar volume of gas at standard conditions (25 dm³·mol⁻¹), the volume of hydrogen gas produced is:

$V(H_2) = n(H_2) imes 25 ext{ dm}^3 \text{ mol}^{-1} = 1 ext{ mol} imes 25 ext{ dm}^3 = 25 ext{ dm}^3$

For experiment 2, we first find the number of moles of hydrogen produced using the average rate:

$n(H_2) = ext{rate} imes ext{time} = 2.08 imes 10^{-4} ext{ mol·s}^{-1} imes (10 ext{ min} imes 60 ext{ s/min}) = 0.125 ext{ mol}$

Using the stoichiometry from the balanced equation, we have 1 mol of magnesium producing 1 mol of hydrogen. Therefore, the initial mass of magnesium used is:

$m(Mg) = n(Mg) imes 24 ext{ g/mol} = 0.125 ext{ mol} imes 24 ext{ g/mol} = 3 ext{ g}$

The collision theory states that for a reaction to occur, reactant particles must collide with sufficient energy and correct orientation. In experiment 2, magnesium powder has a larger surface area compared to magnesium ribbon. This increased surface area allows for more frequent collisions between magnesium and hydrochloric acid molecules, resulting in a faster reaction rate, which is evident by the steeper curve in the graph.