QCAA Chemistry Oxidation and reduction

This is a thermal decomposition reaction where the oxidation states of Calcium (+2), Carbon (+4), and Oxygen (-2) remain unchanged.

This is a combination reaction where no elements change their oxidation states (Ca remains +2, O remains -2, H remains +1).

This is an acid-base neutralization reaction where ions exchange partners without any change in oxidation numbers.

During recharge, they operate as electrolytic cells because an external power source drives the non-spontaneous reverse reaction; they only act as galvanic cells during discharge.

This describes fuel cells. Rechargeable batteries are self-contained and have a fixed amount of reactants enclosed within the cell.

Higher temperatures typically increase the rate of unwanted side reactions, which degrades the battery faster and reduces its overall lifespan.

In galvanic cells, reduction always occurs at the cathode, which is the positive electrode, while oxidation occurs at the negative anode.

Only secondary cells are rechargeable and act as electrolytic cells during the charging process; primary cells cannot be recharged.

Galvanic cells convert chemical energy into electrical energy; the conversion of electrical energy into chemical energy occurs during recharging (electrolysis), which primary cells cannot undergo.

This option incorrectly ranks S as stronger than Q. In Cell 2, Q acts as the anode (oxidized) and S as the cathode (reduced), which demonstrates that Q is a stronger reducing agent than S.

This option incorrectly identifies R as the strongest reducing agent. R produces the largest voltage when reduced by Q, indicating it has the highest reduction potential and is therefore the weakest reducing agent.

This option incorrectly ranks Q as stronger than T. In Cell 3, T acts as the anode (oxidized) and Q as the cathode (reduced), which demonstrates that T is a stronger reducing agent than Q.

This incorrect answer assumes that only 1 mole of electrons is transferred per mole of $O_2$, rather than the actual 4 moles required by the half-equation.

This incorrect answer uses the molar mass of atomic oxygen ($16.0$ g/mol) instead of molecular oxygen ($O_2$, $32.0$ g/mol) to calculate the final mass.

This value represents the number of moles of electrons produced ($n = \frac{36000}{96500} = 0.373$ mol), not the mass of $O_2$ consumed in grams.

Hydrogen fuel cells do not produce carbon dioxide ($CO_2$) during operation; they generate electricity through an electrochemical reaction rather than combustion.

While this describes the efficient energy conversion mechanism of a fuel cell, it does not explain the composition of the emissions or why they are less polluting.

This statement describes the operational requirement of a fuel cell (continuous fuel supply) but is unrelated to the environmental impact or chemical composition of its emissions.

While researching literature can provide background information, it is not an experimental method to directly test the student's specific hypothesis about the cell's current.

Measuring the conductivity of the zinc electrode alone does not measure the overall current produced by the electrochemical reactions in the galvanic cell.

Although it directly measures current, using two different cells introduces confounding variables, such as slight differences in electrolyte concentration or internal resistance, which could affect the results.

This is incorrect. If manganese had an oxidation state of +1, the total charge of the ion would be $+1 + 4(-2) = -7$, which does not match the actual charge of -1.

This is incorrect. While +2 is a very common oxidation state for manganese (e.g., in $\text{Mn}^{2+}$ salts), it is not the oxidation state found in the permanganate ion.

This is incorrect. Manganese is in Group 7 and has only 7 valence electrons, so its maximum possible oxidation state is +7. It cannot reach +8.

In all electrochemical cells, electrons flow through the external circuit from the anode (where they are released via oxidation) to the cathode.

Only electrolytic cells convert electrical energy into chemical energy; fuel cells function as galvanic cells, converting chemical energy into electrical energy.

The polarity depends on the cell type: in fuel cells (galvanic), the anode is negative, whereas in electrolytic cells, the anode is positive.

$\text{SgO}_3$ is on the reactant side of a reduction half-reaction, meaning it accepts electrons and acts as an oxidizing agent, not a reducing agent.

While $\text{Sg}_2\text{O}_5$ can act as a reducing agent by reversing the first equation (oxidation potential of $+0.046 \text{ V}$), it does not have the most positive oxidation potential.

$\text{SgO}_2$ can act as a reducing agent by reversing the second equation, but its oxidation potential would be $-0.11 \text{ V}$, making it a relatively weak reducing agent.

While both fuels produce the same amount of greenhouse gases (1 mole of $\text{CO}_2$ per mole of fuel), they produce a different number of electrons during oxidation.

This is incorrect because they produce the same amount of greenhouse gases (1 mole of $\text{CO}_2$ per mole of fuel) and a different number of electrons.

This is incorrect because they produce the same amount of greenhouse gases, as both fuels contain exactly one carbon atom per molecule and thus yield the same amount of $\text{CO}_2$.

While Fe is oxidised, the oxidising agent is the specific species that accepts electrons. In this reaction, the copper(II) ion ($\text{Cu}^{2+}$) accepts electrons, not neutral copper (Cu).

Oxidation involves the reactant losing electrons. Here, solid Fe is oxidised, not the product $\text{Fe}^{2+}$. Additionally, the oxidising agent is the ion $\text{Cu}^{2+}$, not neutral Cu.

The species being oxidised is the reactant Fe, which loses electrons to become $\text{Fe}^{2+}$. The ion $\text{Fe}^{2+}$ is the product of oxidation, not the substance being oxidised.

$\text{H}^+$ is a product of the recharging reaction (the reverse of the discharge reaction), not a reactant, so it cannot act as the reducing agent.

The oxidation states of hydrogen and oxygen in $\text{H}_2\text{O}$ do not change during the reaction, meaning it does not act as an oxidising or reducing agent.

$\text{VO}_2^+$ is a product of the recharging reaction, not a reactant. While it acts as the oxidising agent during discharge, it does not play this role during recharge.