3.1.5 OXIDATION, REDUCTION & REDOX REACTIONS
EQ: A student oxidised a solution of hydrochloric acid with a few drops of sodium chlorate(I) solution. The reaction mixture effervesced and turned pale green. The gas formed bleached universal indicator paper. (a) Write a half-equation for the oxidation of chloride ions. (1 MARK) (b) Write a half-equation for the reduction of chlorate(I) ions to chlorine in acidic conditions. (1 MARK) (c) Write an overall equation for the redox reaction of chlorate(I) ions with hydrochloric acid. (1 MARK)
(a) 2Cl- → Cl2 +2e- (b) 2ClO- + 4H+ + 2e- → Cl2 + 2H2O (c) ClO- + Cl- + 2H+ → Cl2 + H2O
Q: Aluminium chloride, AlCl3, is a useful ionic substance, used as a catalyst in Friedel-crafts reactions. It is synthesised in industry by reaction aluminium with chlorine (Cl2) at high temperatures. (a) Write the ionic half-equations for this reaction (b) identify the oxidising agent in the reaction between chlorine and aluminium
(a) 3Cl2 + 6e- --> 6Cl- 2Al --> 2Al 3+ + 6e- (b) Chlorine is the oxidising agent
EQ: Chlorine and bromine are both oxidising agents. In an aqueous solution, bromine oxidises sulfur dioxide, SO2, to sulfate ions, SO42-. (a) Deduce a half equation for the reduction of bromine in an aqueous solution (b) Deduce a half equation for the oxidation of SO2 in an aqueous solution forming SO42- and H+ ions (c) Use these two half equations to construct an overall equation for the reaction between aqueous bromine and sulfur dioxide
(a) Br2 + 2e− → 2Br− (b) SO2 + 2H2O → 4H+ + SO4 2− + 2e− (c) Br2 + SO2 + 2H2O → 2Br− + 4H+ + SO4 2−
EQ: in acidified aqueous solution, nitrate ions, NO3- react with copper metal forming nitrogen monoxide, NO, and copper(II) ions. (a) Write a half equation for the oxidation of copper to form copper(II) ions (b) Write a half equation for the reduction, in an acidified solution, of nitrate ions to nitrogen monoxide (c) Write an overall equation for this reaction
(a) Cu → Cu2+ + 2e− (b) NO3− + 4H+ + 3e− → NO + 2H2O (c) 3Cu + 2NO3− + 8H+ → 3Cu2+ + 2NO + 4H2O
EQ: When an acidified solution of sodium nitrate (NaNO2) is added to aqueous potassium iodide, iodine and nitrogen monoxide (NO) are formed. (a) Write a half equation for the conversion of NO2- in an acidic solution into NO. (1 MARK) (b) Write a half-equation for the conversion of I- into I2. (1 MARK) (c) Write an overall ionic equation for the reaction of NO2- in an acidic solution with I- (1 mark) (d) State the role of NO2- in the reaction with I-. (1 MARK)
(a) NO2- + e- + 2H+ → NO + H2O (b) 2I- → I2 + 2e- (c) 2NO2- + 2I- + 4H+ → I2 + 2NO + 2H2O (d) Oxidising agent.
EQ: Nitrogen monoxide, NO, is formed when silver metal reduces nitrate ions, NO3- in acid solution. (a) Write a half equation for the reduction of NO3- ions in acid solution to form nitrogen monoxide and water. (b) Write a half equation for the oxidation of silver metal to Ag+ (aq) ions. (c) Hence, deuce an overall equation for the reaction between silver metal and nitrate ions in acid solution.
(a) NO3− + 4H+ + 3e− → NO + 2H2O (b) Ag → Ag+ + e− (c) NO3− + 4H+ + 3Ag → NO + 2H2O + 3Ag+
EQ: Lead (IV) oxide, PbO2, reacts with concentrated hydrochloric acid to produce chlorine, lead (II) ions, Pb2+, and water. (a) Write a half equation for the formation of Pb2+ and water from PbO2 in the presence of H+ ions (b) Write a half equation for the formation of chlorine from chloride ions (c) Hence deduce an equation for the reaction which occurs when concentrated hydrochloric acid is added to lead (IV) oxide, PbO2 (3 MARKS)
(a) PbO2 + 4H+ + 2e− → Pb2+ + 2H2O (b) 2Cl− → Cl2 + 2e− (c) PbO2 + 4H+ + 2Cl− → Pb2+ + Cl2 + 2H2O
EQ: Iron is extracted from iron(III) oxide using carbon at a high temperature. (i) State the type of reaction that iron(III) oxide undergoes in this extraction. (1 MARK) (ii) Write a half-equation for the reaction of the iron(III) ions in this extraction. (1 MARK)
(i) Reduction OR reduced (ii) Fe3+ + 3e- ---> Fe
Q: Define a disproportionation reaction. (2 MARKS)
- A reaction where the same species undergoes reduction and oxidation; Simultaneously / at the same time
EQ: Halide ions can also react with each other in aqueous solutions. Chlorine reacts in a redox reaction with an aqueous solution of sodium bromide, to form sodium chloride and bromine. Cl2 + NaBr --> NaCl + Br2 Use the reaction above and your knowledge of the halogens, to explain whether chlorine or bromine is a stronger oxidising agent. (2 MARKS)
- Chlorine is a stronger oxidising agent - as it will oxidise a bromide ion - And it is reduced
EQ: State two things that oxidation states can be used for. (2 MARKS)
- Deduce / tell if oxidation or reduction has taken place - Work out what has been oxidised and/or reduced - Construct half equations - Balance redox equations
How do you construct a half equation?
- Deduce oxidation states - Spot redox changes - Oxidation: each e- lost = +1 - Reduction: each e- gained = -1 - Construct half equations
EQ: A solution of sodium chlorate(l) was added to a colourless solution of potassium iodide. Suggest what is observed. Explain the reaction that leads to this observation. (3 MARKS)
- Goes brown (or shades of brown) - Due to iodine or I3− - Because I− oxidised
How do you construct a redox reaction in acidic conditions?
- H+ are sometimes needed for REDOX reactions to occur. 1. Balance redox species (e-s) 2. Any O lost is H2O in products 3. Balance the O 4. Add H+ on reactants to balance the H in H2O
What are the rules of redox equations / reactions?
- Hydrogen - always +1 (except metal hydrides, eg NaH, when it is -1) - Group 1 - always +1 - Group 2 - always +2 - Aluminium - always +3 - Oxygen - always -2 (except in 'peroxides' when it is -1 eg H2O2) - Group 7 - always -1 (except when combine with more electronegative element) - ALL OTHERS ARE VARIABLE!
EQ: Explain how the oxidation state of the oxygen atom in H2O2 is different from its oxidation state in other compounds. (2 MARKS)
- In H2O2 the oxidation state is -1 - Whereas in other compounds it is -2
EQ: Explain why iodide ions react differently from chloride ions. (3 MARKS)
- Iodine ions are larger (have more electron shells) / have larger atomic radius - So electron in outer shell is further away from nucleus - The electron in the outer shell lost from the iodide ion is less strongly held by the nucleus compared with that lost from a chloride ion
EQ: State whether the displacement reaction between copper (II) ions and magnesium is an example of a disproportionation reaction. Justify your answer. (2 MARKS) Cu2+ + Mg --> Mg 2+ + Cu
- It is not a disproportionation reaction - Because reduction and oxidation do not happen to one / the same species - because Cu2+ is reduced while the Mg is oxidised
EQ: State three definitions of reduction.
- Loss of oxygen - Addition / gain of hydrogen - Gain of electrons
EQ: Redox reactions can be identified by either reduction and oxidation occurring or the presence of a reducing agent and an oxidising agent. Explain whether the reaction between sodium hydroxide and hydrochloric acid is a redox reaction or not. (3 MARKS)
- NaOH + HCl --> NaCl + H2O - Determine the correct oxidation states for all elements in the reactants and products: Reactants: Na = +1, O = -2, H = +1, Cl = -1 Products: Na = +1, Cl = -1, H = +1, O = -2 - The reaction is not a redox reaction because reduction and oxidation do not occur as there is no reducing and oxidising agents in the reaction
EQ: A student suggests that the oxidation states of elements in compounds are always whole numbers. Is the student correct? Justify your answer. (1 MARK)
- No - Because oxidation states of elements can be the average oxidation state (of that element in different environments)
EQ: Sodium tetrathionate can also be made by reacting sodium bisulphite, NaHSO3, with disulfur dichloride, S2Cl2. 2 NaHSO3 + S2Cl2 --> Na2S4O6 + 2 HCl Evaluate if this reaction is a disproportionate reaction. (2 MARKS)
- Not disproportionation Alternative 1: - Because the oxidation state of sulfur in NaHSo3 decreases which is reduction and the oxidation state of sulfur in S2Cl2 increases which is oxidation. Alternative 2: - Because the sulfur in NaHSO3 is reduced but the sulfur in S2Cl2 is oxidised
State the oxidation state of compounds
- Overall O.S. is 0 (+ and - oxidation states must balance) - Overall O.S. of a molecular ion is equal to its charge, e.g. NH4+ is +1
EQ: Sulfur can react in different ratios with oxygen to form sulfur dioxide and sulfur trioxide. Write equations to show the oxidation of sulfur to form sulfur dioxide and to form sulfur trioxide. Include state symbols in your equations. In which reaction is the increase in oxidation state the greatest? Explain your answer. (5 MARKS)
- Sulfur dioxide: S (s) + O2 (g) --> SO2 (g) - Sulfur trioxide: 2S (s) + 3O2 (g) --> 2SO3 (g) - The oxidation state of sulfur (S) to sulfur trioxide (SO3) - Increases from 0 to +6 (in SO3)
EQ: The unbalanced equation for the reaction of manganate ions, MnO4-, with iron (II) ions, Fe2+ ions in the presence of acid, H+ to form manganese (II) ions, Mn2+, iron (III) ions, Fe3+ and water is: MnO4- + Fe2+ + H+ --> Mn2+ + Fe3+ + H2O One manganese atom undergoes a change in oxidation state of -5. One iron atom undergoes a change in oxidation state of +1. This leads to the first step of balancing the redox equation, as shown in Figure 1. By comparing the overall unbalanced equations, suggest why this method of balancing might not work as easily for the dichromate equation from parts (a) and (b), shown below. Cr2O7 2- + Cu+ + H+ --> Cr3+ + Cu2+ + H2O
- The dichromate ion / Cr2O7 2- contains two chromium atoms
EQ: Chlorine trifluoride contains only the non metals chlorine and fluorine. Carbon dioxide contains only the nonmetals carbon and oxygen. How do you decide which element has the positive oxidation state? (1 MARK)
- The more electronegative element is the negative oxidation state - The less electronegative element is the positive oxidation state
State the oxidation state of elements
- Uncombined elements have an O.S. of 0, e.g. N2 S8 Cl2 - Elemental ions have an O.S. equal to their charge, e.g. Cl- oxidation state is -1, O2- is -2, K+ is +1
How do you combine half equation to construct a complete redox equation?
- balance the e-s in the two half equations - cancel out e-s - combine half equations
EQ: Antimony is a solid element that is used in industry. The method used for the extraction of antimony depends on the grade of the ore. Antimony can be extracted by reacting scrap iron with low-grade ores that contain antimony sulfide (Sb2S3). 1. Write an equation for the reaction of iron with antimony sulfide to form antimony and iron(II) sulfide. (1 MARK) 2. Write a half-equation to show what happens to the iron atoms in this reaction. (1 MARK)
1. 3Fe + Sb2S3 ---> 3FeS + 2Sb 2. Fe -----> Fe2+ + 2e−
How do you write an ionic equation?
1. Write the full equation 2. Split any soluble compounds in reactants + products (aq) 3. Remove any 'spectator' ions (ie the ones that have not change in oxidation number)
Nitrogen can exist in multiple oxidation states. What is the change of nitrogens oxidation state in NO3- → NH4+?
8.
What is an oxidising agent?
A substance that oxidises another species causing it to lose electrons, and accepts / gains electrons from the species that is being oxidised (therefore it gains electrons and is reduced)
What is a displacement reaction?
A type of reaction in which a more reactive species displaces a less reactive species from its compound.
Find the oxidation state of V in → Ca(VO3)2
Ca = 2+ So VO3- O3 = -6 So V = +5
EQ: For the element X in the ionic compound MX, explain the meaning of the term oxidation state. (1 MARK)
Charge on the ion (or element or atom)
Define what oxidation state is.
Charge on the ion or element or atom
Find the oxidation state of Pt in → Pt(H2O)5(OH)2+
Eliminate water because water has oxidation state of 0 What is left over is Pt(OH)2+ OH is -, so Pt is 3+
Some compounds contain elements with only positive oxidation states only. True or false?
FALSE! In any compound, there must be a difference in electronegativity between two elements. The most electronegative element in a compound always has a negative O.S.
EQ: What are the correct oxidation states for the elements in Fe3O4?
Fe = +2 and +3 O = -2 Look at image for explanation.
EQ: When concentrated sulfuric acid is added to potassium iodide, solid sulfur and a black solid are formed. Deduce the half equation for the formation of sulfur from concentrated sulfuric acid.
H2SO4 + 6H+ 6e- → S + 4H2O
EQ: When iodide ions react with concentrated sulfuric acid in a different redox reaction, the oxidation state of sulfur changes from +6 to -2. The reduction product of this reaction is a poisonous gas that has an unpleasant smell. Identify this gas.
Hydrogen sulfide → H2S
How do you identify what has been oxidised, what has been reduced and the oxidising and reducing agents?
Look at image
Deduce the overall equation for the reaction of copper with NO3 - ions and H+ ions to produce Cu2+ ions, NO2 and water.
Look at image.
EQ: 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. i) Write the balanced, ionic half equations to show the reduction of the manganate ion and the oxidation of the Fe 2+ ion. ii) Use your answers to part (I) to write an overall redox equation for the titration of iron (Ii) sulfate with potassium manganate (VII) solution. (3 MARKS)
Look at image.
EQ: Balance the following redox equation: Cr2O7 2- + Cu+ + H+ --> Cr3+ + Cu2+ + H2O
Look at image.
EQ: Chlorine is used to decrease the numbers of microorganisms in water. When chlorine is added to water, there is a redox reaction, as shown by the equation Cl2 + H2O ⇌ HClO + HCl Give two half-equations to show the oxidation and reduction processes that occur in this redox reaction. (2 MARKS)
Look at image.
EQ: Copper (I) thiocyanate can react with potassium iodate in the presence of hydrochloric acid according to the following unbalanced equation: CuSCN + KIO3 + HCl --> CuSO4 + KCl + HCN + ICl + H2O Write the half equation involving potassium iodate forming potassium chloride and iodine chloride. (2 MARKS)
Look at image.
EQ: Copper (I) thiocyanate, CuSCN, can react with potassium iodate, KLO3, in the presence of hydrochloric acid according to the following unbalanced equation. CuSCN + KLO3 + HCl --> CuSO4 + KCl + HCN + ICI + H2O Write the half equation involving copper (I) thioxynate forming copper sulfate and hydrogen cyanide. (2 MARKS)
Look at image.
EQ: Photochromic glass contains an evenly distributed mixture of copper (I) chloride and silver (I) chloride. When sunlight passes through the glass the silver (I) chloride is separated into its ions. The chloride ions are then converted into chlorine atoms and the silver (I) ions into silver atoms. The silver atoms cluster together causing the lenses of the photochromic glasses to darken. The darkening process caused by the formation of silver atoms is reversible. Write two balanced symbol equations to show how copper (I) chloride can react with the products from part (a) to remove the silver atoms and cause the lens to lighten. (2 MARKS)
Look at image.
EQ: Photochromic glass contains an evenly distributed mixture of copper (I) chloride and silver (I) chloride. When sunlight passes through the glass the silver (I) chloride is separated into its ions. The chloride ions are then converted into chlorine atoms and the silver (I) ions into silver atoms. The silver atoms cluster together causing the lenses of the photochromic glasses to darken. Write equations for the processes involved in the darkening of photochromic glasses and explain if the reaction is reduction to oxidation. (3 MARKS)
Look at image.
EQ: Write a half-equation for the reduction of chlorate(l) ions to chlorine in acidic conditions.
Look at image.
Q: Write the half-equation for the oxidation of hydrogen peroxide to oxygen gas.
Look at image.
EQ: When an acidified solution of sodium nitrate (NaNO2) is added to aqueous potassium iodide, iodine and nitrogen monoxide (NO) are formed. Give the oxidation state of nitrogen in the following species. (2 MARKS) - NO2- - NO
NO2- = +3 NO = +2
EXAM TIP!
OILRIG
What do half equations show?
The gain / loss of e- of each species in an equation
Define the term reduction in terms of electrons.
The gain of electrons → (oxidation no. goes down!)
What is a reducing agent?
donate electrons to species being reduced (therefore it loses electrons and is oxidised)
Redox reactions involve a transfer of ______ from the _______ agent to the _______ agent.
electrons, reducing, oxidising
Define oxidation
element has lost e-s → oxidation no, goes up!
EQ: Cyanide-containing compounds are commonly used in organic chemistry to lengthen the carbon chain. They are also used in gold and silver mining because of their ability to dissolve these metals and their ores. i) Use your knowledge of oxidation states, explain why a student might mistakenly think that the formula for a cyanide ion is CN+. ii) Use your knowledge of oxidation states to identify the correct oxidation state of carbon in the ionic compound, KCN. iii) Use your knowledge of chemical bonding to explain how the cyanide ion has the correct formula of CN- (7 MARKS)
i) - The oxidation state of C is normally +4 and the oxidation state of N is -3 - Combining the oxidation states gives an overall charge of +1 ii) - The oxidation state of potassium = +1 - The oxidation state of nitrogen = -3 KCN is overall neutral so 0 = +1 + C - 3 = +2 iii) - The carbon forms a triple covalent bond with nitrogen - The potassium atom donates / gives an electron to the carbon atom during ionic bonding - The carbon atom now has an extra electron resulting in the negative charge
EQ: Sodium tetrathionate can be formed by reacting sodium thiosulfate, Na2S2O3, with iodine. i) Write a balanced symbol equation for this reaction ii) identify the oxidant in this reaction iii) Describe the expected observation to show that this reaction had gone to completion (4 MARKS)
i) 2 NaS2O3 + I2 --> Na2S4O6 + 2 NaI ii) I2 iii) The discolouration of I2
EQ: 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 alumnium from it. The bauxite is initially purified to produce aluminium (III) oxide. Electrolysis is then carried out, to extract the aluminium from the aluminium oxide. Oxygen gas is also formed as a byproduct of this part of the process. i) Write down the overall equation for the extraction of aluminium from aluminium (III) oxide by electrolysis. ii) State whether the aluminium (II) oxide is oxidised in the electrolysis reaction. Explain your answer. (3 mARKS)
i) 2Al2O3 --> 4Al + 3O2 ii) Reduced; as the Al in aluminium (III) oxide (Al2O3) gains electrons / decreases in oxidation state (from +3 in Al2O3 to 0 in Al)
EQ: Other metals can be extracted from their metal ores by reacting the ore with carbon. Iron, for example is extracted from iron(III) oxide (Fe2O3) in a huge container called a blast furnace, during a redox reaction with carbon. i) Write the overall equation for the reaction or iron (III) oxide with carbon ii) What type of redox reaction is the reaction of iron (III) oxide with carbon? Explain your answer. iii) Use your answer to part (ii) to explain why aluminium cannot be extracted from aluminium (III) oxide by reacting it with carbon (4 MARKS)
i) 2Fe2O3 + 3C --> 4Fe + 3CO2 ii) Displacement reaction: as C replaces Fe from Fe2O3 iii) Aluminium is more reactive than carbon
EQ: Chlorine behaves as an oxidising agent in the extraction of bromine from seawater. In this process, chlorine gas is bubbled through a solution containing bromide ions. i) Write the simplest ionic equation for the reaction of chlorine with bromide ions. ii) Give one observation that would be made during this reaction.
i) Cl2 + 2Br− --> 2Cl− + Br2 ii) Turns to) yellow / orange / brown (solution)
EQ: Chlorine also oxidises sulfur dioxide (SO2) in aqueous solutions to sulfate ions (SO4 2-) under acidic conditions. i) Deduce the half-equation for the reduction of chlorine in aqueous solution ii) Deduce the half equation for the oxidation of sulfur dioxide in aqueous solution iii) Use the two half equations in part i) and ii) to construct an overall redox equation iv) Bromine reacts with sulfur dioxide in a different way to chlorine. Explain why. (5 MARKS)
i) Cl2 + 2e- --> 2Cl- ii) SO2 + 2H2O --> SO4 2- + 4H+ + 2e- iii) Cl2 + SO2 + 2H2O --> 2Cl- + SO4 2- + 4H+ OR Cl2 + SO2 + 2H2O --> HCl + H2SO4 iv) - Bromine atoms are bigger than chlorine atoms - Bromine is a weaker oxidising agent / loses an electron more easily than chlorine
EQ: Chlorine is widely used in water purification. besides killing dangerous germs like bacteria and viruses, chlorine also helps reduce unwanted taste and odours in water by reacting with organic chemicals in water. i) Write the overall redox reaction of chlorine with water ii) Deduce the oxidation states of chlorine in all of the chlorine-containing compounds iii) The reaction between chlorine and water is a disproportional reaction. Use the oxidation states to explain why the reaction above is an example of this.
i) Cl2 + H2O --> HCl + HClO ii) Cl2 = 0, in HCl = -1, in HCl) = +1 iii) - A disproportionation reaction is a reaction in which an atom of the same element gets both oxidised and reduced in the same reaction - The chlorine has been both oxidised and reduced in this reaction - The oxidation state of chlorine changes from 0 in +1 in HClO and -1 in HCl
EQ: The mixing of household bleach with ammonia during cleaning should be avoided, as a redox reaction between ammonia and the chlorate 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) i) What is the oxidising agent in this reaction? Explain your answer. ii) Why are the risks of the toxic chlorine gas being produced greater than the risks of hydrazine?
i) ClO - As it oxidises NH3 to N2H4 and in the process it gets reduced OR the oxidation state of chlorine decreases from +1 in Cl)- to 0 in Cl2 ii) Because Cl2 is a gas therefore it can be inhaled / breathed in
EQ: Thermite is a mixture of finely powdered aluminium and iron oxide. In a thermite reaction, the aluminium reacts with iron oxide to produce hot molten iron. This process is often utilised in remote locations for welding railway lines. Initially, a lot of heat is required to start off the reaction. However, once started, the reaction itself releases a significant amount of heat which is enough to melt the iron. i) Write the overall equation for the thermite reaction, including state symbols ii) Deduce the reduction and oxidation half-equations for this reaction iii) Explain what type of redox reaction the thermite reaction is. iv) What can be deduced about the chemical properties of aluminium and iron metals involved in this reaction
i) Fe2O3 (s) + 2Al (s) --> 2Fe (l) + Al2O3 (s) ii) Reduction: Fe2O3 + 6H+ + 6e- --> 2Fe + 3H2O Oxidation: 2Al + 3H2O --> Al2O3 + 6H+ + 6e- iii) - Displacement reaction; as Al displaces Fe in iron oxide (Fe2O3) iv) Al is more reactive than Fe
EQ: Nitrogen is very unreactive, due to the strong triple bonds within the molecule, and will only react with oxygen in air to form nitrogen oxides under extreme conditions (such as lightning). Nitrogen fixing bacteria in the soil convert the unreactive molecular nitrogen in the atmosphere into ammonia or other related nitrogen compounds. This process is called nitrogen fixation i) Write an ionic equation for the conversion of nitrogen to ammonia and hydrogen gas by nitrogen fixation. ii) Is the reaction in part (i) a reduction or oxidation reaction? Explain your answer. iii) The ammonia produced by nitrogen fixation can be further converted to nitrate ions (NO2 -). Nitrate ions in turn are oxidised to nitrate ions, which are mainly used by plants for protein synthesis. Write an equation for the conversion of nitrate ions (5 MARKS)
i) N2 + 8H+ + 8e- --> 2NH3 + H2 ii) - a reduction reaction - as nitrogen gains electrons: its reduction state decrease from 0 in N2 to -3 in NH3 iii) NO2- + H2O --> NO3- + 2H+ + 2e-
Oxidation and reduction occur __________ in a reaction because one species __________ electrons which are then _________ and gained by other species. Therefore they are known as ________ reactions (reduction - oxidation)
simultaneously, loses, donated, redox
