Electron Transfer Reactions (HL IB Chemistry)

Common household bleach is a cleaning product which smells like chlorine gas and is therefore, also called chlorine bleach.
It contains a mixture of sodium chlorate (NaOCl), sodium chloride and water and can be made by dissolving chlorine gas in a solution of sodium hydroxide.

The mixing of household bleach with ammonia during cleaning should be avoided, as a
redox reaction between the ammonia and the chlorate(I) ions in bleach will generate toxic chlorine gas and hydrazine (N₂H₄). 

The overall redox reaction for this reaction is shown below. 

            2NH₃ (aq) + 2ClO^- (aq) → N₂H₄ (aq) + Cl₂ (g) + 2OH^- (aq)

Due to the risks associated with chlorine-based bleach, alternative bleaches are often used instead. These bleaches are based on peroxides such as hydrogen peroxide.  

Manganate(VII) ions oxidize hydrogen peroxide to oxygen gas. The reaction is carried out with both species under acidic conditions. 

[1]

Explain how the oxidation number of the oxygen atom in H₂O₂ is different from its oxidation state in other compounds.

Metals can often be seen written as a list, from the most reactive metal to the least reactive metal. This list is known as the reactivity series of metals and can be used to predict the feasibility of a reaction. 

Below is a section of the reactivity series of metals, ordered from most to least reactive: 

            Calcium

            Magnesium

            Aluminium

            Zinc

            Iron

            Tin

            Lead 

A piece of zinc was placed into a solution of iron(II) sulfate and a solution of magnesium sulfate.

Predict, giving a reason, whether a reaction would occur in each solution. 

Copper is below lead on the reactivity series shown in part (a). A piece of zinc was placed into a solution of copper(II) sulfate. Write the half equation for the zinc and identify the type of reaction taking place.

Many chemical reactions are redox reactions as they involve the transfer of electrons.

The following reaction is an example of a common redox reaction:

      5Fe²⁺ (aq) + MnO₄^- (aq) +8H⁺ (aq) → 5Fe³⁺ (aq) + Mn²⁺ (aq) + 4H₂O (l) 

Deduce the oxidation numbers of iron and manganese in the above reaction, both as reactants and as products.           

State which substance is reduced.

The amount of iron in some dietary iron supplements was analyzed by redox titration. Four tablets were crushed and dissolved in 50.0 cm³ of 2.00 mol dm^-3 sulfuric acid. The solution was then transferred to a 250 cm³ volumetric flask and made up to 250 cm³ with distilled water. 

A 25.0 cm³ sample of the iron tablets solution was titrated against 0.00500 mol dm^-3 potassium manganate(VII) and 25.8 cm³ was needed for complete reaction. 

Determine the amount of iron, in mol, in one tablet.

Halide ions, such as chloride, Cl^-, can be identified using chemical tests. If an unknown compound is dissolved in dilute nitric acid, and then silver nitrate solution is added, a precipitate will form if the unknown solution contains halide ions. The precipitate formed will be a silver halide. 

The general equation for the precipitation reaction of halide ions with silver nitrate solution is:

            AgNO₃ (aq) + X^- (aq) → AgX (s) + NO₃^- (aq)

Halide ions can also react with other halogens in aqueous solutions. Chlorine reacts in a redox reaction with an aqueous solution of sodium bromide, to form sodium chloride and bromine.

         Cl₂ (aq) + NaBr (aq) → NaCl (aq) + Br₂ (aq)

Chlorine also oxidizes sulfur dioxide (SO₂) in aqueous solutions to sulfate ions (SO₄^2-) under acidic conditions.

Use the two half-equations from part (c) to construct the overall redox equation for this reaction.  

The iron of railway lines rusts when it comes into contact with water and oxygen. The overall redox equation for the rusting of iron is as follows: 

         4Fe (s) + 3O₂ (g) + 6H₂O (g) → 4Fe(OH)₃ (s) 

Define the term reduction.

State, with a reason, the oxidizing agent in this reaction in part (a).

A student investigates the rate of rusting of a piece of iron under different conditions.           

Figure 1 shows the set-up of the students’ experiment.

Figure 1

Predict in which test tube(s) the iron metal will not rust. Explain your answer.

Deduce the oxidation number of each of the stated elements in the ions and compounds to complete Table 1 below.

Table 1 

Aluminium is present in the Earth’s crust in aluminium ore, called bauxite. A number of processes are done to this ore, to extract the aluminium from it. The bauxite is initially purified to produce aluminium oxide, Al₂O₃. Electrolysis is then carried out on the molten Al₂O₃, to extract the aluminium.

Another ionic compound which can undergo electrolysis is molten lead bromide.

State the products formed at each electrode during the electrolysis of molten lead bromide, giving the equations at each electrode with state symbols.  

Draw a labelled diagram of the apparatus suitable to carry out the electrolysis of molten lead bromide. Include the direction of electron flow, the negative electrode (cathode), the positive electrode (anode) and the electrolyte.

The list below shows three metals from the activity series in order of reactivity.

Deduce which of the three metals is the strongest reducing agent.

A voltaic cell can be made by joining two half-cells together, such as  Zn/Zn²⁺ and Ni/Ni²⁺.

Write a balanced equation for the overall reaction taking place when the two half-cells are connected together, and state which species is undergoing oxidation.

Cell diagrams are a way to represent the redox reactions taking place in voltaic cells.

Write a cell diagram for the overall cell reaction taking place in part (b).

Complete the partially labelled diagram in Figure 1, of the apparatus used in the voltaic cell in part (b). Show the direction of the movement of the electrons and ions in the cell.                                                      

Figure 1

Ethene, C₂H₄, can be made into a number of useful compounds. A reaction sequence for this is shown below:

The product of step 2 can undergo combustion.

Give the reagents and conditions needed to carry out step 3.

The product of step 2 has a higher boiling point than the product of step 3.

State the names of the products of step 2 and 3, and explain the difference in their boiling points.

Two isomeric compounds are shown below in Figure 1.

 Figure 1

Compound B, CH₃CH(CH₃)CH(OH)COOH, can be oxidized into compound C.

The same reaction in part (b) can be used to oxidize ethanol into ethanal or ethanoic acid, depending on the reaction conditions.

Outline how the reaction conditions can be changed to produce ethanal or ethanoic acid.

Some standard electrode potential data are shown in Table 1 which you will need to answer the following questions.

Table 1 

Deduce the species from Table 1 that is the weakest oxidising agent. Explain your choice.

Give the conventional representation of the cell that is used to measure the standard electrode potential of copper/copper(II) ions as shown in Table 1 in part (a).

A voltaic cell is made from nickel in a solution of nickel(II) chloride and copper in a solution of copper(II) sulfate.

Calculate the EMF of this cell using the values given in Table 1 in part (a).

Two half-cells, involving species in Table 1, are connected together to give a cell with an EMF = +0.30 V

Aqueous copper(II) sulfate can be electrolysed using passive or active electrodes. Passive electrodes can be made of platinum and active electrodes from copper. 

Draw a labelled diagram of an electrolytic cell for this process using platinum electrodes and identify in which direction electrons flow.

Write the half equations taking place at each electrode in part a), including state symbols, and state what is seen at each electrode.

Write the half equations taking place at each electrode when using copper electrodes, including state symbols, and state what is seen at each electrode.

State what happens to the colour and acidity of the electrolyte when using platinum and copper electrodes in the electrolysis of aqueous copper(II) sulfate.

State the conditions under which the EMF of a redox reaction will be spontaneous.

Using Sections 1 & 19 of the Data Booklet, calculate ΔGᶿ for the following reaction and state whether the reaction is spontaneous under standard conditions.

 Fe²⁺ (aq) + Ni(s) → Fe(s) + Ni²⁺ (aq)

Suggest, with a reason, how a non-spontaneous reaction could be made spontaneous.

Using Table 2, predict and write overall equations for all the spontaneous reactions. 

Table 2

Metals coatings on other metals can be achieved using electroplating. Three beakers containing solutions of Sn(NO₃)₄, Co₂(SO₄)₃, Pb(NO₃)₂, were set up as electrolytic cells and used to electroplate the metals. The same amount of current was passed through the cells for the same length of time.           

State and explain in which cell would the greatest amount of metal be produced and identify the electrode where the metals are deposited. 

Apart from current and time, identify two factors that influence the amount of cobalt deposited in the Co₂(SO₄)₃ cell.

State two reasons why electroplating of metals is carried out.

A nickel teaspoon is electroplated with silver using sodium argentocyanide. Predict the mass changes at each electrode.

An organic reaction sequence is shown below.

State the IUPAC names of the four substances in the sequence.

Classify the reactions in (a) and give the names of the reagents in each step.

Give the reaction conditions for step 3 in (a)

Draw a displayed formula of an isomer of C₄H₁₀O that gives two signals in an ¹H NMR spectrum.

The following reaction pathway is used to produce Compounds A and B, which when reacted together, form a branched ester molecule, Compound C. 

Suggest suitable reagents and conditions for the synthesis of Compound A via Step 1 and give the name for this type of reaction.

In order for the ester to be produced, the ketone in part (a) must be converted to another compound, B.

Outline how ethanol can be synthesised from ethane in two steps. State the reaction conditions and reagents and name the type of reaction taking place.

The four step synthesis to form propan-1-ol from a ketone is outlined below.

Propanal is a versatile organic building block used in the synthesis of plastics and rubber chemicals.

Propanal can be produced from propanoic acid in the following two-step reaction. 

	Step 1		Step 2
Propanoic acid		Propan-1-ol		Propanal

State the reaction type, including suitable reagents, for Steps 1 and 2.

Suggest why it is not possible to convert propanoic acid directly to propanal using the reagent you identified for Step 1 in (a).

Explain why Step 2, in (a), is completed by distillation.

Identify, explain your reasoning, which of the three organic compounds, from the reaction scheme in (a), would be distilled first.

A student sets up a titration to determine the amount of iron(II) sulfate in an iron tablet. They titrate the iron(II) sulfate solution with potassium manganate(VII) solution.

The iron(II) sulfate solution is acidified before titration to stop the manganate ion forming unwanted manganese dioxide. 

Explain the effect that not acidifying the iron(II) sulfate would have on the final calculation of the estimated mass of iron.

The student dissolved the iron tablet in excess sulfuric acid and made the solution up to 250 cm³ in a volumetric flask. 25.0 cm³ of this solution was titrated with 0.0100 mol dm^-3 potassium manganate(VII) solution. The average titre was found to be 26.65 cm³ of potassium manganate(VII) solution. 

Iron sulfate reacts with chromium to produce chromium(III) sulfate, Cr₂(SO₄)₃ , and iron

Deduce the overall ionic equation for the reaction occurring 

Molten potassium bromide can be electrolysed using graphite electrodes.

State the half equations for the oxidation and reduction processes and deduce the overall cell reaction, including state symbols. 

Explain why solid potassium bromide does not conduct electricity. 

A voltaic cell is made from a half-cell containing a magnesium electrode in a solution of magnesium nitrate and a half-cell containing a silver electrode in a solution of silver(I) nitrate.

State the oxidation state of phosphorus in the following compounds.

The tetrathionate ion is shown below:

Sodium tetrathionate can be formed by reacting sodium thiosulfate, Na₂S₂O₃, with iodine.

Describe the expected observation to show that this reaction had gone to completion.

15.00 cm³ of ethanedioic acid, H₂C₂O₄ (aq), requires 10.30 cm³ of a 0.250 mol dm^-3 solution of sodium hydroxide, NaOH (aq), for complete neutralisation using a phenolphthalein indicator for the first permanent colour change.

15.00 cm³ of the same H₂C₂O₄ solution required 12.35 cm³ of potassium permanganate solution, KMnO₄ (aq), solution for complete oxidation to carbon dioxide and water in the presence of dilute sulfuric acid to further acidify the H₂C₂O₄ solution for the first permanent colour change. 

H₂C₂O₄ (aq) + 2NaOH (aq) → Na₂C₂O₄ (aq)+ 2H₂O (l) 

[2]

Deduce the following half equations and overall redox equation for the reaction outlined in part a).

MnO₄^- (aq) to Mn²⁺ (aq) .........................

H₂C₂O₄ (aq) to CO₂ (g) ...........................

Overall equation ......................................

Calculate the concentration, in mol dm^-3, of the potassium manganate(VII), KMnO₄, solution.

Use section 19 of the data booklet to draw the electrochemical cell for the feasible reaction of Ag / Ag⁺ and Al / Al³⁺. Write the conventional representation, including state symbols, for this cell.

Write the conventional representation, including state symbols, for this cell.

Explain why the salt bridge connecting the silver and aluminum electrodes cannot be made with potassium chloride solution.

The silver half cell is replaced with a magnesium half cell. Deduce the reading on the voltmeter.

Use section 19 of the data booklet and the information below to determine if the following reaction is feasible at 298 K.

2KMnO₄ (aq) + 5H₂O₂ (aq) + 6HCl (aq) → 2MnCl₂ (aq)+ 8H₂O (l) + 5O₂ (g) + 2KCl (aq)

   O₂ (g) + 2H⁺ + 2 e^-  Η₂O₂    E^θ = 0.68 V

The reaction of copper oxide and sulfuric acid is shown below. Use section 19 of the data booklet to explain why the reaction is thermodynamically feasible.

Suggest a reason why the reaction does not occur despite being thermodynamically feasible.

Explain why the following does not represent the standard hydrogen electrode.

The standard electrode potential for Zn²⁺ (aq) + 2e^- → Zn (s) is –0.76 V. State the meaning of the minus sign in the value of –0.76 V.

Zinc coating on metals serves as physical protection which prevents rust from affecting the underlying metal surface. This is achieved by electroplating.
 

Using section 19 of the data booklet deduce the full equation for the Cr₂O₇^2- (aq) / Cr³⁺ (aq) and Br₂ (l) / Br^- (aq) cell.

Determine the value for E^Θcell value for the cell outlined in part a). 

Use your answer to part b) and sections 1 and 2 of the data booklet to determine whether the reaction in part a) reaction is spontaneous.

An electrochemical cell has a free energy change of -14.475 kJ mol^-1. Use the information in the table to determine the cell representation of the electrochemical cell. 

Aqueous sodium tetrahydridoborate, NaBH₄, is a common reducing agent.

State the IUPAC name of the two isomers with the formula C₃H₆O that can be reduced by aqueous NaBH₄.

State the IUPAC name of two non-cyclic isomers with the formula C₃H₆O that cannot be reduced by aqueous NaBH₄.

When NaBH₄ is used as a reducing agent followed by the addition of acid, the reduction products of ketones can exhibit optical isomerism, while the reduction products of aldehydes cannot.

Deduce the structure when the following compound is reduced using NaBH₄.