A student performs an experiment to investigate the reaction between magnesium carbonate and hydrochloric acid. The reaction is carried out in an open beaker. They record the loss in mass over time. S

2. (b)

(i) The reaction occurred for approximately 12.5 minutes. This is the point where the curve flattens, indicating that the reaction has stopped and no further mass is lost.

(ii) There was a loss in mass because carbon dioxide gas ($\text{CO}_2(g)$), a product of the reaction, escaped from the open vessel.

(iii) The reaction rate was fastest at the beginning of the reaction.
Reason: The rate of reaction is proportional to the concentration of reactants. At the beginning, the concentrations of $\text{MgCO}_3$ and $\text{HCl}$ are highest, leading to the highest frequency of effective collisions and thus the fastest reaction rate. The slope of the curve is steepest at the beginning.

(iv) Three reaction conditions that can affect the slope of the curve (i.e., the reaction rate) are:

• Concentration of hydrochloric acid
• Temperature of the reaction
• Surface area of magnesium carbonate (e.g., using powdered $\text{MgCO}_3$)
• Presence of a catalyst (though not typically used for this specific reaction, it's a general factor)

2. (c)
The reaction is: $\text{PCl}_5(g) \rightleftharpoons \text{PCl}_3(g) + \text{Cl}_2(g)$; $\Delta H = +95 \text{ kJmol}^{-1}$

(i) Expression for the equilibrium constant K:

(ii) Prediction of the effect on the equilibrium position:
I. Increased pressure: The equilibrium will shift to the left (towards reactants). This is because the reactant side has fewer moles of gas (1 mole) compared to the product side (2 moles), and increasing pressure favors the side with fewer gas moles to reduce the pressure.

II. Increased temperature: The equilibrium will shift to the right (towards products). This reaction is endothermic ($\Delta H = +95 \text{ kJmol}^{-1}$), so increasing the temperature favors the endothermic (forward) reaction to absorb the added heat.

III. Removal of chlorine: The equilibrium will shift to the right (towards products). Removing a product causes the system to produce more of that product to counteract the change, according to Le Chatelier's principle.

(iii) Sketch an energy profile diagram for the forward reaction:
The forward reaction is endothermic ($\Delta H = +95 \text{ kJmol}^{-1}$), meaning the products have higher energy than the reactants.

      Energy
         ^
         |      / \
         |     /   \  Activation Energy
         |    /     \
  Products |   -------
           |  /
           | /
           |/
  Reactants -------
           |
           +----------------------> Reaction Progress
           <---- ΔH = +95 kJmol⁻¹ ---->

3. (a)

(i) Two properties of iron that are characteristic of transition metals are:

• It forms colored compounds (e.g., $\text{Fe}^{2+}$ is green, $\text{Fe}^{3+}$ is brown/yellow).
• It exhibits variable oxidation states (e.g., +2 and +3).
• It acts as a catalyst (e.g., in the Haber process).
• It forms complex ions.

(ii) Equations for the extraction of iron and removal of impurities in the blast furnace:

• Production of reducing agent (carbon monoxide): $$\text{C}(s) + \text{O}_2(g) \longrightarrow \text{CO}_2(g)$$ $$\text{CO}_2(g) + \text{C}(s) \longrightarrow 2\text{CO}(g)$$
• Reduction of iron ore (hematite): $$\text{Fe}_2\text{O}_3(s) + 3\text{CO}(g) \longrightarrow 2\text{Fe}(l) + 3\text{CO}_2(g)$$
• Removal of acidic impurities (e.g., silica, $\text{SiO}_2$): $$\text{CaCO}_3(s) \longrightarrow \text{CaO}(s) + \text{CO}_2(g)$$ $$\text{CaO}(s) + \text{SiO}_2(s) \longrightarrow \text{CaSiO}_3(l)$$