Electron transfer reactions: notes and practice questions
- This topic covers electron transfer reactions, including oxidation, reduction, and electrochemical cells.
- Oxidation is electron loss, an increase in oxidation state, or oxygen gain/hydrogen loss.
- Reduction is electron gain, a decrease in oxidation state, or oxygen loss/hydrogen gain.
- Deduce oxidation states of atoms in compounds or ions.
- Identify oxidized and reduced species, and oxidizing and reducing agents.
- Half-equations show electron loss (oxidation) or gain (reduction).
- Electrochemical cells involve oxidation at the anode and reduction at the cathode.
- Voltaic cells convert spontaneous redox energy to electrical energy.
- Electrolytic cells use electrical energy for non-spontaneous reactions, such as molten salt electrolysis.
- Organic functional groups like alcohols, carboxylic acids, ketones, alkenes, and alkynes can undergo oxidation or reduction.
How it is examined
Oxidation state deduction is a 1-mark part in both papers. Balancing half-equations is 1 to 2 marks and the balance must be right in both atoms and charge. Voltaic cell diagrams attract "identify the anode", "state the direction of electron flow" and "explain the role of the salt bridge" at 1 mark each, and the salt bridge answer must mention ion movement to maintain electrical neutrality, not just "completes the circuit". At HL, E⦵cell calculations are [2] and the sign error from reversing the wrong half-cell is the classic loss. The organic oxidation sequence (primary alcohol to aldehyde with distillation, to carboxylic acid with reflux) is examined by asking for conditions as well as products.
HL: standard reduction potentials, the Faraday constant F, and ΔG⦵ = −nFE⦵cell. At SL nothing is given, but the guidance says displacement data will be supplied in the question. What is recall: oxidation state rules, anode is oxidation and cathode is reduction, the sign convention (anode negative in a voltaic cell, positive in an electrolytic cell), and the oxidation products of primary versus secondary versus tertiary alcohols.
- 3.2.1 Oxidation and reduction can be described in terms of electron transfer, change in oxidation state, oxygen gain or loss, or hydrogen loss or gain. Students deduce oxidation states of an atom in a compound or an ion, and identify the oxidized and reduced species and the oxidizing and reducing agents in a chemical reaction.
- 3.2.2 Half-equations separate the processes of oxidation and reduction, showing the loss or gain of electrons. Students deduce redox half-equations and equations in acidic or neutral solutions.
- 3.2.3 The relative ease of oxidation and reduction of an element in a group can be predicted from its position in the periodic table. The reactions between metals and aqueous metal ions demonstrate the relative ease of oxidation of different metals. Students predict the relative ease of oxidation of metals and of reduction of halogens, and interpret data regarding metal and metal ion reactions.
- 3.2.4 Acids react with reactive metals to release hydrogen. Students deduce equations for reactions of reactive metals with dilute HCl and H₂SO₄.
- Half-equations and equations are required in acidic or neutral solutions only. Balancing in alkaline solution is not on the syllabus.
- Names and formulas of specific oxidizing agents, and the mechanisms of oxidation, will not be assessed. The same for reducing agents and the mechanisms of reduction. Include the role of hydride ions in the reduction reaction.
- 3.2.12 The hydrogen half-cell `H⁺(aq) + e⁻ ⇌ ½H₂(g)` is assigned a standard electrode potential of zero by convention. It is used in the measurement of standard electrode potential, E⦵. Students interpret standard electrode potential data in terms of ease of oxidation and reduction.
- 3.2.13 Standard cell potential, E⦵cell, can be calculated from standard electrode potentials. E⦵cell has a positive value for a spontaneous reaction. Students predict whether a reaction is spontaneous in the forward or reverse direction from E⦵ data.
- 3.2.14 The equation `ΔG⦵ = −nFE⦵cell` shows the relationship between standard change in Gibbs energy and standard cell potential for a reaction. Students determine ΔG⦵ from E⦵ data.
- 3.2.15 During electrolysis of aqueous solutions, competing reactions can occur at the anode and cathode, including the oxidation and reduction of water. Students deduce from standard electrode potentials the products of the electrolysis of aqueous solutions.
Guiding questions
- What happens when electrons are transferred?
Linking questions
- Structure 3.1 What are the advantages and limitations of using oxidation states to track redox changes? Why does metal reactivity increase, and non-metal reactivity decrease, down the main groups? How can oxidation states be used to show that CH₄, CH₃OH, HCHO, HCOOH, CO₂ are given in increasing order of oxidation?
- Structure 2.3 The surface oxidation of metals is often known as corrosion. What are some of the consequences of this process?
- Structure 2.1 Under what conditions can ionic compounds act as electrolytes?
- Structure 3.2 How does the nature of the functional group in a molecule affect its physical properties, such as boiling point?
- Reactivity 1.3 Electrical energy can be derived from the combustion of fossil fuels or from electrochemical reactions. What are the similarities and differences? What is the difference between combustion and oxidation of an alcohol?
- Reactivity 2.3 Secondary cells rely on electrode reactions that are reversible. What are the common features of these reactions?
- Reactivity 3.4 Why are some reactions of alkenes classified as reduction reactions while others are classified as electrophilic addition reactions?
- Reactivity 1.4 (HL) How can thermodynamic data also be used to predict the spontaneity of a reaction?
- Structure 3.1 (HL) Why is there a colour change when an alcohol is oxidized by a transition element compound?
- Tool 1, Inquiry 2 Why are some redox titrations described as "self-indicating"? How is an electrolytic cell used for electroplating?
Practice questions
51 questions · 23 easy · 24 medium · 4 hardQuestion 1
EasyPaper 2 · calculator1 markSulfur is a versatile element, forming compounds with a wide range of oxidation states. In which of the following substances does sulfur exhibit the lowest oxidation number?
A
B
C
D
Recall the rules for assigning oxidation numbers. Hydrogen is typically +1, and oxygen is typically -2. For an element in its elemental form, the oxidation number is 0.
Question 2
MediumPaper 2 · calculator1 markA brass key is to be electroplated with copper. Which of the following statements are correct?
I The brass key should be the cathode.
II The anode is made of an inert material like graphite.
III The electrolyte could be aqueous copper(II) sulfate.
A. I only
B. II only
C. I and III only
D. I, II and III
Consider the purpose of each component in an electrolytic cell for electroplating. Where does reduction (the deposition of metal) occur? What should the anode be made of to maintain the concentration of metal ions in the solution? What must the electrolyte contain?
Question 3
HardPaper 1B · calculator16 marksNitrogen dioxide, a pollutant from car exhausts, reacts with water in the atmosphere to form nitric acid and nitrous acid. This is a disproportionation reaction.
(a) Deduce the oxidation states of nitrogen in the reactant and products.
Reactant:
Products: ,
(b) Explain, with reference to the equilibrium, why more nitrogen dioxide gas dissolves when the reaction occurs in alkaline rainwater.
(c) The solubility of nitrogen dioxide gas in water was measured by different scientific groups. A summary of their results is shown.
| Source | Temperature / °C | Solubility of gas in 0.100 dm³ of water |
|---|---|---|
| A | 0 | 0.380 dm³ |
| B | 10 | 0.28 dm³ |
| C | 20 | 200 cm³ |
| D | 25 | 0.15 L |
| E | 30 | 0.120 dm³ |
(i) Identify a problem in comparing the data from the different sources as presented in the table.
(ii) The units of solubility are converted to mol dm⁻³. Complete the table by calculating the value for source A. Assume the atmospheric pressure is 100 kPa and the density of the resulting solution is 1.00 g cm⁻³.
(iii) Suggest an explanation for the effect of temperature on the solubility of nitrogen dioxide gas.
(d) Suggest one reason why nitrogen dioxide is considered a major air pollutant.
(e) Nitrous acid, , is a weak acid. The graph shows the percentage of nitrous acid and its conjugate base, the nitrite ion (), present at different pH values.
(GRAPH IS A STANDARD SPECIATION PLOT FOR A WEAK ACID. X-AXIS: pH from 0 to 8. Y-AXIS: Percentage from 0 to 100. A curve for HA starts at 100% and goes down, a curve for A- starts at 0% and goes up. The two curves cross at pH = 3.3, where each is at 50%.)
(i) Deduce the pH range where nitrous acid, , is the dominant nitrogen-containing species in the solution.
(ii) Determine, with reference to the graph, the of nitrous acid.
(f) Nitrous acid can react with secondary amines to form N-nitrosamines, which are potent carcinogens. An example is the reaction with dimethylamine, .
(i) Deduce a balanced chemical equation for the formation of N-nitrosodimethylamine, , from dimethylamine and nitrous acid.
(ii) The rate of N-nitrosamine formation is highly dependent on pH. The reaction rate is highest under mildly acidic conditions where there is a sufficient concentration of both the unprotonated amine and nitrous acid. State two conditions that could be maintained in an industrial process to minimize the formation of N-nitrosamines.
(g) To combat the effects of acid rain, powdered limestone () is sometimes added to lakes. Suggest two distinct reasons why this 'liming' process is effective at restoring the aquatic ecosystem.
Remember the rules for assigning oxidation states. Oxygen is usually -2 and hydrogen is usually +1. The sum of oxidation states in a neutral molecule is zero.
Alkaline conditions imply the presence of a base, like hydroxide ions (). How would a base react with the acidic products of the forward reaction? Consider Le Châtelier's principle.
Look closely at the units used for solubility in the table. Are they all consistent? What other variable that affects gas solubility might be missing?
To find the concentration in mol dm⁻³, you first need to find the moles of gas. You are given the volume of the gas, the temperature, and the pressure. The ideal gas law might be useful here.
Observe the trend in the table: as temperature increases, what happens to the volume of gas that dissolves? The dissolution of a gas in a liquid is an equilibrium process. How does temperature affect exothermic and endothermic equilibria?
Consider the direct effect of on human health or its role in forming other harmful substances in the environment.
The 'dominant' species is the one present in a higher concentration (or percentage). Find the part of the graph where the curve for is above the curve for .
The has a special significance on a speciation graph. It is the pH at which the concentrations of the weak acid and its conjugate base are equal.
The reactants are and . The products are and one other small molecule. Identify the atoms that are not part of the nitrosamine product to deduce the other product.
To minimize the formation of a product, you can either remove one of the reactants or change the conditions (like pH) to make the reaction much slower. Consider what pH would minimize the concentration of one of the key reactants, .
Think about the direct chemical effect of adding a base () to an acidified lake. Then, consider how this action might affect the equilibrium involving the atmospheric pollutant that caused the acid rain in the first place.
Question 4
EasyPaper 1A · calculator1 markHydrogen sulfide, , and the sulfate ion, , are common species containing sulfur.
What are the oxidation states of sulfur in these species?
A. +2, +8
B. -2, +6
C. -2, +8
D. +2, +6
Recall the rules for assigning oxidation states. The oxidation state of hydrogen is typically +1 and oxygen is typically -2. The sum of oxidation states in a neutral molecule is zero, while in a polyatomic ion, it equals the overall charge of the ion.
Question 5
MediumPaper 2 · calculator1 markWhen a student performs a redox titration, iron(II) ions react with permanganate ions in an acidic solution. The unbalanced equation for this reaction is given below:
When the following redox equation is balanced using the smallest possible whole numbers, what is the sum of all the coefficients?
A 20
B 22
C 24
D 26
Balance the oxidation and reduction half-reactions separately. Remember to balance oxygen atoms with and hydrogen atoms with in acidic solution, then balance charge with electrons. Finally, combine the half-reactions, ensuring the number of electrons lost equals the number of electrons gained.
Question 6
HardPaper 2 · calculator24 marksAntimony (Sb) and Bismuth (Bi) are elements in group 15 of the periodic table.
(a) Antimony has two stable isotopes. 57.21% of antimony atoms contain 70 neutrons and the remainder contain 72 neutrons.
(i) Deduce the nuclear symbol of the isotope of antimony containing 72 neutrons. Use section 6 of the data booklet.
(ii) Calculate, to two decimal places, the relative atomic mass of antimony.
Bismuth(III) nitrate, , is a common salt of bismuth.
(b) (i) The compound contains both ionic and covalent bonds. State which particles are joined by covalent bonds and which are joined by ionic bonds.
(ii) Distinguish between covalent and ionic bonding in terms of electron distribution.
(iii) State the enthalpy term that characterizes the strength of the bonding between the ions in an ionic solid.
(iv) Write an equation for the formation of aqueous bismuth(III) nitrate from solid bismuth(III) oxide and nitric acid.
(v) Calculate the volume, in , of nitric acid required to react completely with of solid bismuth(III) oxide.
(vi) Predict, with a reason, whether bismuth(III) oxide is expected to be primarily acidic, basic or amphoteric.
(vii) Discuss how the relative reactivity of zinc and bismuth could be established using the metals and aqueous solutions of their nitrates.
(viii) Discuss the products formed at the electrodes during the electrolysis of aqueous bismuth(III) nitrate. Use the standard electrode potential and section 24 of the data booklet.
Bismuth compounds are sometimes used in fireworks to produce special effects.
(c) (i) State the feature of the atomic emission spectrum of an element that corresponds to its first ionization energy.
(ii) Calculate the wavelength, in nm, that corresponds to the first ionization energy of bismuth. Use sections 1, 2 and 8 of the data booklet.
(iii) Explain why the first ionization energy of bismuth is lower than that of polonium (Po), in terms of nuclear charge and electron shielding.
First, find the atomic number (number of protons) for Antimony from the periodic table. The mass number is the sum of protons and neutrons. The nuclear symbol is written with the mass number as a superscript and the atomic number as a subscript to the left of the element symbol.
The relative atomic mass is the weighted average of the masses of its isotopes. You'll need the mass number and abundance of both isotopes. The abundance of the second isotope is 100% minus the abundance of the first.
Consider the structure of the polyatomic nitrate ion and how it interacts with the bismuth cation. Covalent bonds typically form between non-metal atoms, while ionic bonds form between metal cations and non-metal anions.
Think about what happens to the valence electrons in each type of bond. Are they shared or transferred?
This term refers to the enthalpy change when one mole of a solid ionic compound is formed from its gaseous ions.
Bismuth(III) oxide is a basic oxide. It will react with an acid in a neutralization reaction to form a salt and water. Remember to balance the equation and include state symbols.
First, calculate the moles of bismuth(III) oxide using its mass and molar mass. Then, use the stoichiometry from your balanced equation in (b)(iv) to find the moles of nitric acid required. Finally, use the concentration of the nitric acid to find the volume.
Consider the position of bismuth in the periodic table and the trend in metallic character down group 15. How does the acid-base character of oxides change with metallic character?
A more reactive metal can displace a less reactive metal from a solution of its salt. Describe a simple experiment to test this.
At the cathode (negative electrode), reduction occurs. Compare the standard electrode potentials for the reduction of and water. At the anode (positive electrode), oxidation occurs. Compare the oxidation of water and the nitrate ion.
Ionization corresponds to the removal of an electron, which means the electron transitions to the n=∞ energy level. What happens to the spectral lines as they approach this limit?
First, find the first ionization energy of Bismuth from the data booklet (in kJ mol⁻¹). Convert this to energy per atom (in J) using Avogadro's constant. Then use the Planck-Einstein relation () and the wave equation () to find the wavelength.
Bismuth and Polonium are in the same period. Consider how the number of protons and the location of the valence electrons change as you move from Bi to Po.
Question 7
EasyPaper 2 · calculator1 markIn the electrolysis of molten sodium chloride, , using inert electrodes, which of the following statements is correct?
A. Sodium metal is formed at the positive electrode (anode).
B. Chlorine gas is formed at the negative electrode (cathode).
C. For every mole of sodium metal produced, two moles of chlorine gas are produced.
D. For every mole of chlorine gas produced, two moles of sodium metal are produced.
First, write the half-equations for the reduction at the cathode and the oxidation at the anode. Remember that in electrolysis, the cathode is the negative electrode and the anode is the positive electrode. Then, balance the electrons to find the overall equation and the mole ratio of the products.
Question 8
MediumPaper 2 · calculator1 markUse the following table of standard reduction potentials to answer the question.
| Half-reaction | |
|---|---|
| -2.71 | |
| -0.76 | |
| -0.74 | |
| -0.45 | |
| 0.00 | |
| +0.34 | |
| +0.52 | |
| +0.54 | |
| +0.77 | |
| +0.80 | |
| +1.36 |
A voltaic cell is constructed using the following half-cells: .
Which statement about this cell is correct?
A. The mass of the silver electrode decreases.
B. The chromium electrode is the cathode.
C. Electrons flow from the silver electrode to the chromium electrode.
D. The concentration of increases.
First, use the standard reduction potentials to identify which half-cell is the anode (oxidation) and which is the cathode (reduction). Then consider what happens at each electrode and the direction of electron flow.
Question 9
HardPaper 2 · calculator23 marksA sample of chlorine consists of two isotopes, and .
(a) Contrast the sub-atomic structure of these two isotopes.
(b) (i) The sample of chlorine is analysed in a mass spectrometer, producing a spectrum for the ion. The spectrum shows three peaks at m/z values of 70, 72 and 74. Explain the origin and relative heights of these three peaks, given that the abundance of is approximately three times that of .
(ii) A more precise measurement finds the composition by mass to be: : 75.76%, : 24.24%. Calculate the relative atomic mass of chlorine from this sample, giving your answer to two decimal places. (Use isotopic masses of 35.0 and 37.0 for this calculation).
Magnesium chloride, , and manganese(II) chloride, , are two ionic compounds.
(c) (i) Deduce the type of bonding in magnesium chloride, , using electronegativity values from section 9 of the data booklet.
(ii) Determine the lattice enthalpy of magnesium chloride, assuming the bonding is purely ionic. Use sections 9, 10 and 12 of the data booklet and the following data:
Enthalpy of formation of magnesium chloride =
(iii) Explain, with reference to electron configurations, why the ionic radii of , and are different. Use section 10 of the data booklet.
(iv) Predict, with a reason, which has the stronger ionic bonding, manganese(II) chloride, , or magnesium chloride.
Magnesium chloride is white, but manganese(II) chloride is pale pink.
(d) (i) State the condensed electron configuration of a manganese atom.
(ii) State the reason, in terms of electron configuration, why manganese(II) chloride is coloured.
(iii) Manganese(II) chloride absorbs light with a wavelength of approximately 530 nm. Describe why this is consistent with the observed colour of the compound. Use sections 2 and 15 of the data booklet.
A copper key is to be electroplated with manganese using an aqueous solution of manganese(II) chloride as the electrolyte.
(e) (i) Deduce the half-equations for the reactions occurring at the anode (made of pure manganese) and the cathode (the copper key).
(ii) Deduce a balanced chemical equation for the reaction of fluorine gas with the aqueous chloride ions in the electrolyte.
Isotopes of an element have the same number of protons but a different number of another sub-atomic particle. What is this particle and how does its number differ between and ?
The peaks correspond to different combinations of the two chlorine isotopes in a diatomic molecule. The height of each peak is related to the probability of that specific combination occurring. Consider the relative abundances of the isotopes.
The relative atomic mass is the weighted average of the isotopic masses. Multiply each isotopic mass by its fractional abundance and sum the results.
Find the electronegativity values for magnesium and chlorine. The difference in their electronegativity values will indicate the type of bonding.
Construct a Born-Haber cycle for the formation of from and . Use Hess's Law to find the unknown lattice enthalpy. Remember to account for the stoichiometry, especially for chlorine.
Compare the number of electron shells and the nuclear charge (number of protons) for each ion.
The strength of ionic bonding depends on the charge of the ions and the distance between them (ionic radii). Compare these factors for and .
Manganese is in the first row of the d-block. Remember the filling order of the 4s and 3d sub-levels.
The colour of transition metal compounds is related to the electronic structure of the transition metal ion. What is special about the d-sublevel in coloured ions?
The colour we see is the complementary colour to the one that is absorbed. Use the colour wheel in the data booklet to find the complementary colour of the absorbed light.
In electroplating, the object to be plated is the cathode, and the metal used for plating is the anode. Oxidation occurs at the anode and reduction occurs at the cathode.
Consider the relative oxidizing strengths of the halogens. A more reactive halogen will displace a less reactive halide from its salt solution.
Question 10
EasyPaper 2 · calculator1 markA student performs a redox titration where potassium permanganate solution is used to determine the concentration of an iron(II) salt. The overall balanced ionic equation for the reaction is shown below.
What is the oxidising agent in this reaction?
A.
B.
C.
D.
Recall the definitions of oxidation, reduction, oxidising agent, and reducing agent. Determine the oxidation states of relevant elements in the reactants and products.
Question 11
MediumPaper 2 · calculator1 markA student is performing various chemical reactions in the laboratory and needs to classify them based on whether electron transfer occurs. Which one of the following reactions is NOT a redox (reduction-oxidation) reaction?
A
B
C
D
Recall that a redox reaction involves a change in oxidation states for at least two elements. Assign oxidation states to all elements in reactants and products for each reaction.
Question 12
HardPaper 2 · calculator16 marksStrontium is an alkaline earth metal in group 2 of the periodic table, used to produce the red colour in fireworks.
(a) In a sample of strontium, 10.0% of the atoms have a mass number of 86 and the remainder have a mass number of 88.
(i) Deduce the nuclear symbol of the isotope of strontium with mass number 86. Use section 6 of the data booklet.
(ii) Calculate, to two decimal places, the relative atomic mass of this sample of strontium.
(iii) Explain why the first ionization energy of strontium is greater than that of rubidium.
(b) Strontium nitrate has the formula .
(i) The compound contains both ionic and covalent bonds. State which particles are joined by covalent bonds and which are joined by ionic bonds.
(ii) Contrast covalent and ionic bonds in terms of how the valence electrons are involved.
(iii) Write a balanced chemical equation for the formation of strontium nitrate solution from the reaction of solid strontium hydroxide with nitric acid.
(iv) Calculate the volume, in , of nitric acid required to react completely with of solid strontium hydroxide.
(v) Predict, with a reason, whether strontium oxide is acidic, basic or amphoteric.
(vi) Describe an experiment to establish the relative reactivity of strontium and zinc using the metals and aqueous solutions of their nitrates.
The nuclear symbol shows the mass number as a superscript and the atomic number as a subscript. Find the atomic number of strontium from the data booklet.
The relative atomic mass is the weighted average of the masses of the isotopes. Use the formula: .
Consider the number of protons (nuclear charge) and the principal energy level of the valence electrons for both strontium and rubidium.
Identify the metal cation and the polyatomic anion. What type of bond holds them together? What type of bonds exist within the polyatomic anion?
Think about what happens to the outermost electrons in each type of bonding: are they given away, taken, or shared?
This is a neutralization reaction between a base (metal hydroxide) and an acid. The products are a salt and water. Remember to balance the equation and include state symbols.
First, calculate the moles of strontium hydroxide from its mass. Then, use the stoichiometry from your balanced equation in (b)(iii) to find the moles of nitric acid. Finally, use the concentration of the acid to find the required volume.
Consider strontium's position in the periodic table. Is it a metal or a non-metal? What is the general trend in the acid-base character of oxides for metals in its group?
A reactivity series is determined by displacement reactions. Think about what would happen if you place the more reactive metal into a solution containing the ions of the less reactive metal.
Question 13
EasyPaper 2 · calculator1 markThe following table provides standard electrode potentials () for several half-reactions at 298 K.
| Half-reaction | / V |
|---|---|
Which species is the strongest reducing agent?
A.
B.
C.
D.
A reducing agent is a species that is oxidized in a redox reaction. The strongest reducing agent is the species that is most easily oxidized. This corresponds to the half-reaction with the most negative standard reduction potential. Identify the species that is oxidized in this half-reaction.
Question 14
MediumPaper 2 · calculator1 markUse the following table of standard reduction potentials to answer the question.
| Half-reaction | |
|---|---|
| +2.87 | |
| +1.36 | |
| +1.23 | |
| +1.09 | |
| +0.80 | |
| +0.54 | |
| +0.34 | |
| +0.15 | |
| 0.00 | |
| -0.45 | |
| -0.76 | |
| -2.37 |
Which species in the table can oxidize bromide ions () to bromine ()?
A. and
B. and
C. and
D. and
For a species to oxidize bromide ions, it must be a stronger oxidizing agent than bromine (). A stronger oxidizing agent has a more positive standard reduction potential. Compare the values from the table for the reduction of the potential oxidizing agents with the value for the half-reaction.
Question 15
EasyPaper 2 · calculator1 markIn a laboratory experiment, a student is investigating various redox reactions. Which of the following changes requires a reducing agent?
A
B
C
D
A reducing agent causes reduction in another species, meaning the oxidation state of that species decreases. Determine the oxidation state of the key element in each reactant and product.
Question 16
MediumPaper 2 · calculator8 marksA student investigates electrochemical cells using the following standard electrode potentials.
| Half-equation | |
|---|---|
(i) Identify the species that is the strongest oxidizing agent in the table.
(ii) Identify the species that is the strongest reducing agent in the table.
A voltaic cell is constructed using the zinc and copper half-cells under standard conditions.
(i) State the half-equation for the reaction occurring at the anode.
(ii) Calculate the standard cell potential, , in V.
Deduce, justifying your answer, whether a spontaneous reaction will occur when solid silver metal, , is placed into a solution of iron(II) nitrate.
The strongest oxidizing agent is the species that is most easily reduced (has the most positive value). The strongest reducing agent is the species that is most easily oxidized (its corresponding reduction half-equation has the most negative value).
(i) Oxidation occurs at the anode. Compare the values for zinc and copper to determine which metal is more readily oxidized.
(ii) Use the formula . The cathode is where reduction occurs (the half-reaction with the more positive ).
A reaction is spontaneous if the overall cell potential, , is positive. Assume silver is oxidized and iron(II) is reduced, then calculate the resulting .
Question 17
EasyPaper 1A · calculator1 markWhat is the oxidation state of manganese in the permanganate ion, ?
A. +2
B. +4
C. +7
D. +8
The sum of the oxidation states of all atoms in a polyatomic ion must equal the overall charge of the ion. The oxidation state of oxygen is typically -2.
Question 18
MediumPaper 2 · calculator5 marksUse the following table of standard reduction potentials to answer the questions.
-2.71
-0.76
-0.74
-0.45
0.00
+0.34
+0.52
+0.54
+0.77
+0.80
+1.36
A student constructs a voltaic cell using standard zinc and iron electrodes immersed in solutions of their respective ions.
(a) Write the overall balanced chemical equation for the spontaneous reaction occurring in this cell, including state symbols.
(b) Calculate the standard cell potential, , for this voltaic cell.
Identify which species will be oxidized and which will be reduced by comparing their standard reduction potentials. The species with the more positive (less negative) standard reduction potential will be reduced.
The standard cell potential can be calculated using the formula or (where is the negative of the standard reduction potential for the oxidation half-reaction).
Question 19
EasyPaper 1A · calculator1 markIn which compound does manganese have an oxidation state of ?
A.
B.
C.
D.
Recall the rules for assigning oxidation states. The oxidation state of oxygen in most compounds is . The oxidation state of an alkali metal (Group 1) in a compound is . The sum of the oxidation states in a neutral compound is zero. Calculate the oxidation state of manganese in each option.
Question 20
MediumPaper 2 · calculator6 marksButan-2-ol, , is a secondary alcohol that can be oxidized using acidified potassium dichromate(VI) solution to produce an organic compound, X.
(a) Identify the organic compound X and state the colour change observed for the oxidizing agent.
(b) The structure of a compound can be confirmed using ¹H NMR spectroscopy. Predict the number of signals and the ratio of the areas under the signals for the ¹H NMR spectrum of butan-2-ol.
(c) Predict the number of signals and the ratio of the areas under the signals for the ¹H NMR spectrum of compound X.
Consider the product formed when a secondary alcohol is oxidized. Recall the characteristic colours of the dichromate(VI) ion and the chromium(III) ion formed upon reduction.
Identify the number of unique proton environments in the butan-2-ol molecule. The ratio of the areas under the signals corresponds to the ratio of the number of protons in each of these unique environments.
First, draw the structure of compound X (butanone). Then, identify the non-equivalent proton environments and count the number of protons in each environment to determine the integration ratio.
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