Practice Questions

List any three rules used for the calculation of the oxidation number of an element in a compound or ion.

Define the term 'reductant' or 'reducing agent'.

Identify the oxidizing agent in the following reaction: $\mathrm{Zn(s)} + \mathrm{Cu^{2+}(aq)} \rightarrow \mathrm{Zn^{2+}(aq)} + \mathrm{Cu(s)}$

Explain 'reduction' according to the classical idea. Provide one example for the addition of hydrogen and one for the removal of oxygen.

Analyze the reaction $2H_2(g) + O_2(g) \rightarrow 2H_2O(l)$ and identify the species undergoing oxidation.

What is the oxidation number of an element when it is in its free or uncombined state? Give two examples.

Justify why fluorine, the most powerful oxidizing agent, exhibits only a -1 oxidation state in its compounds and cannot undergo disproportionation.

Calculate the oxidation number of Phosphorus in Magnesium Phosphate, $Mg_3(PO_4)_2$.

Justify the classification of the thermal decomposition of potassium chlorate ($\text{KClO}_3$) as an intramolecular redox reaction.

Summarize the key differences between the classical concept and the electron transfer concept of redox reactions. Provide one example reaction that can be explained by both concepts.

Describe the difference between an oxidant and a reductant in a redox reaction.

List and describe the five main steps involved in balancing a redox reaction using the oxidation number method in an acidic solution.

Using Stock notation, represent the compound Manganese(VII) oxide.

Define oxidation and reduction in terms of electron transfer. Identify the species being oxidized and reduced in the reaction: $2\mathrm{Na(s)} + \mathrm{Cl_2(g)} \rightarrow 2\mathrm{NaCl(s)}$

What is a redox couple?

Identify the type of redox reaction for each of the following: (i) $\mathrm{C(s)} + \mathrm{O_2(g)} \rightarrow \mathrm{CO_2(g)}$ (ii) $2\mathrm{H_2O(l)} \rightarrow 2\mathrm{H_2(g)} + \mathrm{O_2(g)}$ (iii) $\mathrm{Zn(s)} + \mathrm{CuSO_4(aq)} \rightarrow \mathrm{ZnSO_4(aq)} + \mathrm{Cu(s)}$

Describe the primary function of a salt bridge in an electrochemical cell like the Daniell cell.

Identify the reducing agent in the reaction: $3CuO(s) + 2NH_3(g) \rightarrow 3Cu(s) + N_2(g) + 3H_2O(l)$

A strip of nickel metal is placed in a 1M solution of silver nitrate. Analyze the reaction that occurs and write the balanced net ionic equation. Identify which metal is more reactive based on this observation.

Analyze the following reaction and identify the substance oxidized, substance reduced, oxidizing agent, and reducing agent. $I_2(s) + 10HNO_3(aq) \rightarrow 2HIO_3(aq) + 10NO_2(g) + 4H_2O(l)$

Compare the reactions of zinc metal with dilute hydrochloric acid and with cold water. Explain the difference in reactivity and classify the type of redox reaction in both cases.

Critique the classical definition of oxidation (addition of oxygen/electronegative element) using the reaction $2\text{Na(s)} + \text{H}_2(\text{g}) \rightarrow 2\text{NaH(s)}$. Justify why the electron transfer concept is a superior model for this specific case.

Design a simple galvanic cell using cobalt and nickel electrodes. Given the standard electrode potentials: $\text{Co}^{2+}(\text{aq}) + 2\text{e}^- \rightarrow \text{Co}(\text{s})$, $E^\ominus = -0.28$ V $\text{Ni}^{2+}(\text{aq}) + 2\text{e}^- \rightarrow \text{Ni}(\text{s})$, $E^\ominus = -0.25$ V

Justify your choice of anode and cathode, write the half-reactions, and formulate the overall cell reaction.

Evaluate the statement: 'The concept of fractional oxidation number, as in tetrathionate ion ($\text{S}_4\text{O}_6^{2-}$), is physically meaningless because electrons cannot be transferred in fractions.' Propose the correct interpretation.

Evaluate the two methods for balancing redox reactions (Oxidation Number vs. Half-Reaction). Propose a specific type of reaction where the half-reaction method is clearly superior and justify your choice.

Propose an experimental procedure using chlorine water ($\text{Cl}_2(\text{aq})$) and an organic solvent like $\text{CCl}_4$ to distinguish between aqueous solutions of sodium bromide (NaBr) and sodium iodide (NaI). Justify the underlying redox principles.

Hydrogen peroxide ($\text{H}_2\text{O}_2$) can act as both an oxidizing and a reducing agent. Formulate two separate balanced net ionic equations to justify this dual behavior. In the first, it oxidizes $\text{Fe}^{2+}$ to $\text{Fe}^{3+}$ in an acidic solution. In the second, it reduces $\text{MnO}_4^-$ to $\text{Mn}^{2+}$ in an acidic solution. Critique why $\text{H}_2\text{O}_2$ acts as a reducing agent in one case and an oxidizing agent in the other.

Formulate a balanced net ionic equation for the redox titration of ferrous ions ($\text{Fe}^{2+}$) with dichromate ions ($\text{Cr}_2\text{O}_7^{2-}$) in an acidic medium, resulting in ferric ions ($\text{Fe}^{3+}$) and chromic ions ($\text{Cr}^{3+}$). A student suggests that since the dichromate solution is orange and the chromic ion solution is green, the endpoint can be detected without an indicator. Critique this suggestion and propose a more reliable method for endpoint detection in this specific titration.

Name the specific type of redox reaction in which an element in one oxidation state is simultaneously oxidized and reduced.

Calculate the oxidation number of the central carbon atom (marked with *) in the structure of carbon suboxide ($O=\stackrel{+2}{C}=\stackrel{*}{C}=\stackrel{+2}{C}=O$).

Demonstrate why the decomposition of silver carbonate, $Ag_2CO_3(s) \rightarrow Ag_2O(s) + CO_2(g)$, is not a redox reaction, whereas the decomposition of silver oxide, $2Ag_2O(s) \rightarrow 4Ag(s) + O_2(g)$, is a redox reaction.

Balance the following redox reaction in an acidic medium using the half-reaction method: $MnO_4^-(aq) + C_2O_4^{2-}(aq) \rightarrow Mn^{2+}(aq) + CO_2(g)$

Balance the following redox reaction using the oxidation number method in a basic medium. Also, identify the oxidant and reductant. $Cr(OH)_3(s) + IO_3^-(aq) \rightarrow I^-(aq) + CrO_4^{2-}(aq)$

The oxoanions of chlorine are $\text{ClO}^-$, $\text{ClO}_2^-$, $\text{ClO}_3^-$, and $\text{ClO}_4^-$. Evaluate their ability to undergo disproportionation. For each species that can disproportionate, formulate a balanced chemical equation for the reaction. Justify why one of them cannot.

Explain what is meant by a disproportionation reaction. Use the decomposition of hydrogen peroxide as an example to illustrate your explanation.

The reaction of white phosphorus ($P_4$) with aqueous NaOH is a disproportionation reaction. Formulate the balanced ionic equation for this reaction, where the products are phosphine gas ($\text{PH}_3$) and the hypophosphite ion ($\text{H}_2\text{PO}_2^-$). Justify its classification as disproportionation by assigning oxidation numbers.

Explain the paradox of fractional oxidation numbers. Describe the actual oxidation states of the sulfur atoms in the tetrathionate ion, $\mathrm{S_4O_6^{2-}}$, and show how the fractional value is obtained.

Calculate the oxidation number of iron in $Fe_3O_4$. Explain why this compound gives a fractional average oxidation number and what the actual oxidation states of the iron atoms are, given that it is a mixed oxide of $FeO$ and $Fe_2O_3$.

Chlorine gas is passed through a hot concentrated solution of sodium hydroxide. The reaction results in the formation of sodium chloride and sodium chlorate. Write the balanced chemical equation for this disproportionation reaction.

Balance the following equation using the ion-electron (half-reaction) method in an acidic medium: $Zn(s) + NO_3^-(aq) \rightarrow Zn^{2+}(aq) + NH_4^+(aq)$

Design a series of experiments to identify the contents of three unlabelled beakers containing aqueous solutions of silver nitrate ($\text{AgNO}_3$), zinc sulfate ($\text{ZnSO}_4$), and magnesium sulfate ($\text{MgSO}_4$). You are provided with strips of copper (Cu) and zinc (Zn) metal. Justify your entire procedure based on the principles of competitive electron transfer and the relative positions of these metals in the electrochemical series ($E^\ominus_{\text{Mg}^{2+}/\text{Mg}} < E^\ominus_{\text{Zn}^{2+}/\text{Zn}} < E^\ominus_{\text{Cu}^{2+}/\text{Cu}} < E^\ominus_{\text{Ag}^{+}/\text{Ag}}$).

A reaction involves permanganate ion ($MnO_4^−$) reacting with hydrogen peroxide ($H_2O_2$) in an acidic solution to produce manganese(II) ion ($Mn^{2+}$) and oxygen gas ($O_2$). (a) Write the balanced half-reactions for oxidation and reduction. (b) Write the overall balanced net ionic equation. (c) Analyze the role of hydrogen peroxide in this reaction and compare it with its role in the reaction $2H_2O_2(aq) \rightarrow 2H_2O(l) + O_2(g)$.

A student attempts to balance the following reaction in a basic medium using the half-reaction method and presents the final incorrect answer. Critique the student's final equation, identify the likely errors in their process, and formulate the correctly balanced equation.

Student's Incorrect Equation: $2\text{MnO}_4^-(\text{aq}) + 3\text{I}^-(\text{aq}) + \text{H}_2\text{O}(\text{l}) \rightarrow 2\text{MnO}_2(\text{s}) + \text{I}_3^-(\text{aq}) + 2\text{OH}^-(\text{aq})$

Propose a reason why in the reaction of $\text{Pb}_3\text{O}_4$ with HCl, $\text{Pb(IV)}$ is reduced to $\text{Pb(II)}$, but with $\text{HNO}_3$, it remains as $\text{PbO}_2$.