Laboratory work "Oxidation-reduction reactions"

«Ufa state petroleum technological university»
The general and analytic chemistry department
Laboratory report №5
on “Oxidation-reduction reactions”
Theoretic explanation:
Oxidation-reduction reactions are called reactions that occur with varying degrees of oxidation of the
participating in them elements. The degree of oxidation of the element is a formal charge, which was
would be on an atom if all the connections in the compound were ionic. Characteristic The degree of
oxidation of the element is determined by the configuration of the valence electron layer.
The oxidizer is an electron acceptor, that is, "takes electrons" and is restored. The reducing agent is an
electron donor, that is, "gives up electrons" and oxidizes. In oxidation-reduction reactions the reducing
agent passes into the corresponding oxidized form, and vice versa, oxidizer - in the appropriate restored
form. Each of the reagents with their product form a conjugated oxidation-reduction pair.
Oxidation is a process in which the degree of oxidation of an element increases due to the recoil of
electrons. Recovery is the process in which the degree of oxidation of the element decreases due to the
addition electrons. The processes of oxidation and reduction proceed simultaneously. Oxidationreduction reactions are divided into intermolecular, intramolecular, disproportionation (or selfoxidation-self-healing), and the correlation.
In the equations of oxidation-reduction reactions, there must be reflected "electronic" and "material"
balance. Electronic balance: The number of electrons "given" by the reductant must be equal to the
number electrons "taken" by the oxidizer. Material balance: the number of atoms one element on the
left and right sides of the equation must be the same.
The electronic balance method is used to compose equations reactions of any type of oxidationreduction reactions. It includes the definition of oxidation states elements, the least common multiple
of the number of "given" by the reductant and "taken" by the oxidizer of electrons, the determination of
the coefficients electronic balance:
The method of ion-electronic balance is used to compose the equations of the oxidation-reduction
reactions proceeding in aqueous solutions taking into account the dissociation of strong, readily soluble
electrolytes per ions. It involves the preparation of 35 half-reactions of oxidation and reduction. And the
material balance is selected using particles of H2O and H + - for reactions in an acid medium and
particles of H2O and OH- - for reactions in an alkaline medium.
MnO4– + 8H+ + 5e → Mn2+ + 4H2O
2
(reduction process)
SO32– + H2O – 2e → SO42– + 2H+
5
(oxidation process)
Further, the charge of each of the systems is determined before and after conversion; calculate the
number of electrons "returned" by the reductant and "received" by the oxidizer and determine the
coefficients of the electronic balance; Addition of half reactions of oxidation and reduction, multiplied
by the corresponding coefficients of the electronic balance and "reduction of such terms" is carried out.
The result is a short ionic equation:
2MnO4- + 6H + + 5SO32- → 2Mn2 + + 5SO42- + 3H2O.
The coefficients obtained in the short ionic equation are transferred to the molecular equation (with
regard to the composition of the compounds):
2KMnO4 + 3H2SO4 + 5Na2SO3 → 2MnSO4 + 5Na2SO4 + 3H2O + K2SO4
Determine the coefficients before the elements not participating in the OBR (for example, K). The
properties of the conjugated oxidation-reduction pair characterize the oxidation-reduction potential.
(𝝋𝒐𝒙/𝒓𝒆𝒅 , 𝑽).
The oxidation-reduction reaction proceeds spontaneously if the change in the free Gibbs energy (ΔG) is
negative: ΔG <0. For oxidation-reduction reactions taking place in an aqueous medium, the change in
Gibbs free energy is related to the values of the redox potentials by the relation ΔG = -nF (φ𝑜𝑥 - 𝜑𝑟𝑒𝑑 ) <0,
where n is the number of electrons, F is the Faraday constant [Kul / mol], and 𝛗𝐨𝐱 and 𝛗𝐫𝐞𝐝 are the
oxidation-reduction potentials (V) of the oxidizer and reductant system, respectively. The difference in
oxidation-reduction potentials of the oxidizer and the reducing agent is called the electromotive force
of the reaction (EMF) and is measured in volts (V). Thus, the ORR between the oxidant and the given
reductor proceeds spontaneously in the forward direction if the EMF is positive: EMF = [𝜑𝑜𝑥 -𝜑𝑟𝑒𝑑 ]> 0 or
𝜑𝑜𝑥 > 𝜑𝑟𝑒𝑑 .
Experiment 1
Oxidizing and reducing properties of iodine
Work objective: study the behavior of iodine (I2) and its oxidizing and reducing properties
Equipment: glass test tubes, sticks, micro-spatula, pipettes.
Reactants: K2Cr2O7, H2SO4, KI, I2, Na2S2O3, FeCl3
Course of work
1. The solution became darker (the violet colour of the iodine become darker), the precipitate did not
fall out.
K2Cr2O7 + 7H2SO4 + 6KI → Cr2(SO4)3 + 3KMnO4 + 4K2SO4 +7H2O
2K(+) + Cr2O7(2-) + 2H(+) + SO4(2-) + K(+) + I(-) → 2Cr(3+) + 3SO4(2-) + 3I2(0) + 2K(+) + SO4(2-) +
2H(+) + O(2-)
Cr2O7(2-) + I(-) → 2Cr(3+) + I2(0) + O(2-)
Cr2O7(2-) + 14H(+) + 6e(-) → 2Cr(3+) + 7H2O (Reduction process)
2I(-) - 2e(-) → I2(0) (Oxidation process)
ϕ (oxidizer) - ϕ(reducer) = 1.4 V - 0.54 V = 0.86 V
2. The violet colour of the iodine becomes colorless. The precipitate did not fall out.
I2 + 2Na2S2O3 → 2NaI + Na2S4O6
I2(0) + 2Na(+) + S2O3(2-) → Na(+) + I(-) + 2Na(+) + S4O6(2-)
I2(0) + S2O3(2-) → 2I(-) + S4O6(2-)
I2(0) + 2e(-) → 2I(-) (Reduction process)
S2O3(2-) + 3H2O - 6e(-) → S4O6(2-) + 6H(+) (Oxidation process)
ϕ (oxidizer) - ϕ(reducer) = 0.621 V - 0.09 V = 0.612 V
3. Green solution became dark red, the precipitate did not fall out.
2FeCl3 + 2KI + 2H2SO4 → K2SO4 + I2 + FeCl2 + FeSO4 + 4HCl
Fe(3+) + 3Cl(-) + K(+) + I(-) + 2H(+) + SO4(2-) → 2K(+) + SO4(2-) + I2(0) + Fe(2+) + 2Cl(-) + Fe(2+) +
SO4(2-) + H(+) + Cl(-)
Fe(3+) + I(-) → I2(0) + Fe(2+)
Fe(3+) + e(-) → Fe(2+) (Reduction process)
2I(-) - 2e(-) → I2(0) (Oxidation process)
ϕ (oxidizer) - ϕ(reducer) = 0.771 V - 0.621 V = 0.15 V
4. The solution became darker (the violet colour of the iodine become darker). The precipitate did
not fall out.
2FeCl3 + 2KI → 2KCl + 2FeCl2 + I2
Fe(3+) + 3Cl(-) + K(+) + I(-) → K(+) + Cl(-) + Fe(2+) + 2Cl(-) + I2(0)
Fe(3+) + I(-) → I2(0) + Fe(2+)
Fe(3+) + e(-) → Fe(2+) (Reduction process)
2I(-) - 2e(-) → I2(0) (Oxidation process)
ϕ (oxidizer) - ϕ(reducer) = 0.771 V - 0.621 V = 0.15 V
Conclusion: we studied that iodine (I2) can exhibit both reducing (reactions 1,3,4) and
oxidizing (reaction 2) properties.
Experiment 2
Oxidizing properties of potassium permangate (KMnO4) in various
medium
Work objective: study the behavior of potassium permangate (KMnO4) in various medium
and its oxidizing properties.
Equipment: glass test tubes, sticks, micro-spatula, pipettes.
Reactants: KMnO4, NaOH, Na2SO3, H2O, H2O2, H2SO4
Course of work
1. We see the appearance of a green colour characteristic of the ion MnO4(2-).
2KMnO4 + 2NaOH + Na2SO3 → Na2SO4 + K2MnO4 + Na2MnO4 + H2O
K(+) + MnO4(-) + Na(+) + OH(-) + 2Na(+) + SO3(2-) + H2O → 2Na(+) + SO4(2-) + 2K(+) + MnO4(2-) +
2Na(+) + MnO4(2-) + H2O
MnO4(-) + SO3(2-) → SO4(2-) + MnO4(2-)
MnO4(-) + e(-) → MnO4(2-) (Reduction process)
SO3(2-) + 2OH(-) - 2e(-) → SO4(2-) + H2O (Oxidation process)
ϕ (oxidizer) - ϕ(reducer) = 0.56 V - (-0.93) V = 1.49 V
2. The violet colour of the ion MnO4(-) becomes slightly pink, almost colorless, characteristic of the
ion Mn(2+).
2KMnO4 + 3H2SO4 + 5Na2SO3 → 2MnSO4 + 5Na2SO4 + K2SO4 + 3H2O
K(+) + MnO4(-) + 2H(+) + SO4(2-) + 2Na(+) + SO3(2-) → Mn(2+) + SO4(2-) + 2Na(+) + SO4(2-) + 2K(+)
+ SO4(2-) + H2O
MnO4(-) + SO3(2-) → SO4(2-) + Mn(2+)
MnO4(-) + 8H(+) + 5e(-) → Mn(2+) + 4H2O (Reduction process)
SO3(2-) + H2O - 2e(-) → SO4(2-) + 2H(+) (Oxidation process)
ϕ (oxidizer) - ϕ(reducer) = 1.51 V - (-0.93) V = 2.44 V
3. We observe the formation of a brown precipitate typical of manganese dioxide MnO2.
2KMnO4 + H2O + 3Na2SO3 → 2MnO2 + 3Na2SO4 + 2KOH
K(+) + MnO4(-) + H2O + 2Na(+) + SO3(2-) → MnO2↓ + 2Na(+) + SO4(2-) + K(+) + OH(-)
MnO4(-) + SO3(2-) → SO4(2-) + MnO2↓
MnO4(-) + 2H2O + 3e(-) → MnO2↓ + 4OH(-) (Reduction process)
SO3(2-) + 2OH(-) - 2e(-) → SO4(2-) + H2O (Oxidation process)
ϕ (oxidizer) - ϕ(reducer) = 0.6 V - (-0.93) V = 1.53 V
The violet colour of the ion MnO4(-) becomes colorless, characteristic of the ion Mn(2+).
2KMnO4 + 5H2O2 + 3H2SO4 → 5O2 + 8H2O + K2SO4 + 2MnSO4
K(+) + MnO4(-) + H2O2 + 2H(+) + SO4(2-) → O2 + H2O + Mn(2+) + SO4(2-) + 2K(+) + SO4(2-)
MnO4(-) + H2O2 → Mn(2+) +
MnO4(-) + 8H(+) + 5e(-) → Mn(2+) + 4H2O (Reduction process)
H2O2 - 2e(-) → O2 + 2H(+)
ϕ (oxidizer) - ϕ(reducer) = 1.51 V - 0.68 V = 0.83 V
Conclusion: Mn(+7) exhibits oxidative ability because it is easier for him to take an electron.
In an acidic environment, permanganate ion exhibits the strongest oxidizing properties and is
able to react with most reducing agents. In a neutral medium, permanganate ion precipitates.