Concentrated sulfuric acid
SO42
+ 4H+
+ 2e = H2SO3
+ H2O
SO42
+ 8H+
+ 6e = S + 4H2O
Nitric acid
NO3
+ 4H+
+ 3e = NO + 2H2O
NO3
+ 3H+
+ 2e = HNO2
+ H2O
NO3
+
2H+
+ e = NO2
+ H2O
NO3
+ 10H+
+ 8e = NH4+
+ 3H2O
Manganese compounds
MnO4
+ 8H+
+ 5e = Mn2+
+ 4H2O
MnO4
+ 2 H2O
+ 3e = MnO2
+
4OH
MnO4
+ e = MnO42
Chromium compounds
Cr2O72
+ 14H+
+ 6e = 2Cr3+
+ 7H2O
CrO42
+ 4H2O
+ 3e = Cr(OH)63
+
2OH
Hydrogen peroxide
H2O2
+ 2e = 2OH
H2O2 + 2H+ + 2e = 2H2O
2H+ + O2 + 2e = H2O2
2H2O
+
O2
+ 2e = H2O2
+ 2OH
EXPERIMENTAL PART
1. Transfer of an ion to a higher oxidation state
Take 6-8 drops of chromium (III) nitrate and add excess of NaOH. When the precipitate of chromium hydroxide dissolves add 3-4 drops of 3% H2O2 solution. Heat the mixture until the color turns yellow. Write down and balance a corresponding redox-reaction.
2. Redox properties of hydrogen peroxide
a). Take 3 drops of KI solution, add 2 drops of diluted sulfuric acid and 3% H2O2 solution. Add some starch solution to indicate the evaluation of iodine. Write down and balance a corresponding redox-reaction.
b). Take 6 drops of KMnO4 solution, add 2 drops of diluted sulfuric acid and dropwise 3% H2O2. Observe the evaluation of oxygen. Write down and balance a corresponding redox-reaction.
3. Oxidizing properties of potassium permanganate
Fill 3 test tubes with 5 drops of KMnO4 each. Add 2 drops of diluted sulfuric acid into the first test tube, 2 drops of distilled water into the second, and 2 drops of sodium hydroxide into the third one. Add Na2SO3 solution into each test tube until the change of their color. Write down and balance corresponding redox-reactions.
4. Oxidation of cations of d-elements
Take 2 drops of a manganese (II) salt in a test tube, then add 5-6 drops of nitric acid and some crystals of NaBiO3. Observe the appearance of a pink color of HMnO4.. Write down and balance a corresponding redox-reaction.
5. Reducing properties of p-elements
Take 3-4 drops of a solution of SnCl2 in a test tube, add NaOH until the formed precipitate dissolves, and 2-3 of a bithmuth (III) salt solution. Observe the black precipitate of elementary bithmuth. Write down and balance a corresponding redox-reaction.
QUESTIONS AND PROBLEMS
1. Complete the equations of the following reactions and balance them:
(a) K2S + KMnO4 + H2SO4 = S + ....
(b) KI + K2Cr2O7 + H2 SO4 = I2 + ...
( c) K MnO4 + H2O2 = ...
2. Indicate the direction in which the following reactions can proceed spontaneously:
(a) H2O2 + HClO = HCl + O2 + H2O
(b) H3PO4 + 2HI = H3PO3 + I2 + H2O
3. Can a salt of iron (III) be reduced to a salt of iron (II) in an aqueous solution by (a) potassium bromide, (b) potassium iodide?
4. Using the table of standard electrode potentials, calculate the equilibrium constants for the following reactions:
(a) Zn + CuSO4 = Cu + ZnSO4
(b) Sn + Pb(CH3COO)2 = Sn(CH3COO)2 + Pb
Experiment 5 complex (coordinate) compounds
By complex (coordinate) compounds are meant definite chemical compounds formed by a combination of individual components without formation of new pairs of electrons.
Examples of complex compounds are Na3[Co(NO2)6], [Cu(NH3)4]SO4, K4[Fe(CN)6].
In a molecule of a complex compound, one of the atoms, generally positively charged, occupies the central site (central ion or complexing agent). Oppositely charged ions or neutral molecules called ligands are coordinated around the central ion. The complexing agent and the ligands form inner sphere of a complex compound. It is characterized by coordinate bonds which are formed while overlapping of empty p- and d-orbitals of a central ion and orbitals containing lone electron pairs of ligands. The ions in the outer sphere are mainly bonded to the complex ions by forces of electrostatic interaction (ionic bonds).
The total number of coordinate bonds formed by the complexing agent with ligands is known as coordination number of the central ion. In accordance with the number of coordinate bonds formed by a ligand with the central ion, the ligand may be a monodentate, bidentate, or polydentate.
Majority of complex compounds are electrolytes. In solutions they dissociate and form both simple and complex ions (outer and inner spheres). This type of dissociation is an irreversible process (complex compounds are strong electrolytes).
[Cu(NH3)4]SO4
[Cu(NH3)4]2+
+ SO42
The inner sphere of a complex compound dissociates reversibly and stepwise (complex ions are weak electrolytes):
[Cu(NH3)4]2+
[Cu(NH3)3]2+
+ NH3
[Cu(NH3)3]2+
[Cu(NH3)2]2+
+ NH3
[Cu(NH3)2]2+
[Cu(NH3)]2+
+ NH3
[Cu(NH3)]2+
Cu2+
+ NH3
Each of the above processes can be characterized by an equilibrium constant (stepwise instability constants of a complex ion). The equilibrium constant of the overall process is called an overall instability constant of a complex ion. The less the value of the overall instability constant is, the more stable is the complex ion.
The shifting of dissociation equilibrium in systems containing complex ions follows the same rules as in solutions of simple (non-complex) electrolytes, namely: equilibrium shifts towards the direction of the most complete binding of the complexing agent or ligand so that the concentrations of these particles remaining unbound in the solution take on the minimum possible values in these conditions. The equilibrium also shifts towards the formation of a more stable complex ion.
According to the character of dissociation of complex compounds, cationic, anionic and neutral complexes are distinguished.
In cationic complexes ligands are usually neutral molecules so that the inner sphere is charged positively: [Cr(H2O)6]Cl3, [Co(NH3)6]Cl3.
In anionic complexes ligands are usually negatively charged: K2[HgI4], Na[Sb(OH)6].
Neutral complexes have both anions and neutral molecules as ligands: [Pt(NH3)2Cl2]. They have no outer sphere and don’t dissociate in aqueous solutions.
EXPERIMENTAL PART