1b Redox Reactions Flashcards
Define oxidation and reduction in 4 ways.
Oxidation involves the gain of oxygen/loss of hydrogen/loss of electron(s)/increase in oxidation number.
Reduction involves the loss of oxygen/gain of hydrogen/gain of electron(s)/decrease in oxidation number.
Define an oxidising agent.
An oxidising agent is an electron acceptor which oxidises the other reagent while it undergoes reduction by receiving the electrons lost by the oxidised species (itself is reduced).
Define a reducing agent.
Reducing agent is an electron donor which reduces the other reagent while it undergoes oxidation by giving electrons to cause the reduction (itself is oxidised).
Describe redox reactions.
Redox reactions are chemical reactions that involve the transfer of electrons from a reducing agent (electron donor) to an oxidising agent (electron acceptor). The overall reaction can be separated into two simpler processes involving electron transfer. The two separate equations can be termed ion-electron equations or half-equations where one half-equation represents the oxidation process while the other represents the reduction process. The addition of the two half-equations gives the overall redox reaction. In such reactions, the total number of electrons lost by the reducing agent is equal to the total number of electrons gained by the oxidising agent.
State the oxidation number for an atom, a compound, hydrogen and fluorine.
Atoms in their elemental state have an oxidation number of 0 and the algebraic sum of the oxidation numbers of atoms in a compound is also 0.
Hydrogen has an oxidation number of +1 in all compounds except for metal hydrides such as NaH and MgH2. Fluorine is the most electronegative element hence its oxidation number in all compounds is -1.
Explain how to balance a redox reaction in an acidic and basic medium.
Pg 7 - 9 of notes.
State two types of redox titrations. State the (a) oxidising/reducing agents that are used in these two types of titrations respectively, (b) the overall equation for the reaction, (c) what the titration can be used for, (d) key features of the titration and (e) the respective colour changes that take place during the titration, taking into account the reactant that is in the conical flask/burette.
(i) Manganate (VII) titrations
(a) Purple potassium manganate (VII), KMnO4 is a powerful oxidising agent that can be used for the estimation of a wide range of reducing agents such as Iron(II) salts, ethanedioates - C2O4 2- and hydrogen peroxide).
(b) MnO4- (aq) + 8H+ (aq) + 5e -> Mn2+ (aq) + 4H2O (l)
(c) Purple potassium manganate (VII), KMnO4 is a powerful oxidising agent that can be used for the estimation of a wide range of reducing agents such as Iron(II) salts, ethanedioates - C2O4 2- and hydrogen peroxide).
(d) In an acidic medium that is provided by sulfuric acid, the MnO4- ion is reduced to Mn2+ which is faint pink/colourless in dilute solutions. Hydrochloric and nitric acid cannot be used for the reaction as Cl- can be oxidised by the acidified KMnO4 to form Cl2 and nitric acid is also an oxidising agent.
In a neutral or alkaline medium, MnO4- ion is reduced to solid MnO2 which is black/brown in colour.
(e) When acidified KMnO4 is in the burette and reacts with hydrogen peroxide, the end-point colour change is from colourless to pale pink. When it reacts with Fe2+, the end-point colour change is from yellow to pale orange/pink. The colour change is reversed when KMnO4 is in the conical flask instead.
(ii) Iodine-thiosulfate titrations
(a) Thiosulfate (S2O3 2-) is a reducing agent and reduces iodine to iodide, while itself becomes (S4O6 2-).
(b) 2S2O3 2- (aq) + I2 (aq) -> S4O6 2- (aq) + 2I- (aq)
(c) This reaction is used to estimate iodine or substances that liberate iodine from potassium iodide.
(d) Most iodometric titrations involve two steps: (i) A fixed amount of substance is normally added to acidified potassium iodide solution to liberate a fixed amount of iodine. The iodine formed is then titrated against a standard thiosulfate solution and using the titration results, the amount of thiosulfate solution, amount of iodine formed and the initial amount of substance can be determined.
The starch solution is often used as an indicator and is normally added when the solution is pale yellow when the majority of the iodine has been reacted away (giving a more distinct colour change). It is not added in the beginning as the iodine molecules tend to be trapped in the starch molecules, leading to inaccurate results. The amount of starch used should be the same for all titrations. Starch solutions that are no longer fresh or improperly prepared. The indicator will then not behave properly at the endpoint and a quantitative determination is not possible.
The titrated solution may slowly become blue again after being left aside as the iodide in the reaction mixture may be oxidised by atmospheric oxygen, giving iodine which combines with the starch to reform the blue colour. Hence, in an acid solution, prompt titration of the liberated iodine is necessary in order to prevent oxidation.
(e) When iodine is in the conical flask, the end-point colour change upon titrating with Na2S2O3- is from pale yellow to colourless. However, upon the addition of starch indicator placed in the conical flask, the end-point colour change is from blue-black to colourless.
