Reduction-oxidation equilibrium in electrolyte’s solution

1. Lecture 4 Reduction-oxidation equilibrium in electrolyte’s solution Associate prof. L.V. Vronska Associate prof . M.M. Mykhalkiv

2. Outline 1. Reduction-oxidation reactions, main concepts. 2. Equilibrium constant of Reduction-oxidation reactions. 3. Influence of different factors on value of redox potential. 4. Usage of reduction-oxidation reactions in analysis.

3. 1. Reduction-oxidation reactions, mainconcepts. • Oxidation state (oxidation number)– the oxidation state is an indicator of the degree of oxidation of an atom in a chemical compound. The formal oxidation state is the hypothetical charge that an atom would have if all bonds to atoms of different elements were 100% ionic. • Oxidation - a loss of electrons. • Reduction - a gain of electrons. • Reducing agent (reductant or reducer) - a species that donates electrons to another species. • Oxidizing agent(oxidant or oxidizer) - a species that accepts electrons from another species.

5. Redox reaction - an electron-transfer reaction. As a result of this electron transfer, some of the elements involved in the reaction undergo a change in oxidation state. Ox + ne  Red оxidizingreducing formform Those species experiencing an increase in their oxidation state are oxidized, while those experiencing a decrease in their oxidation state are reduced.

6. The pair of an oxidizing and reducing agent that are involved in a particular reaction is called aredox pair. Equation Ox + n  Red describes the reduction-oxidation half-reaction. redox pair – is the system of oxidizing and reducing forms of substance, in which oxidizing form (oxidizer) is an electron acceptor and is itself reduced when it accepts electrons, reducing form (reducer) is electron donor and is itself oxidized when it gives up electrons.

7. The most important oxidizing agents: • (NH4)2S2O8, KMnO4, K2Cr2O7, K2CrO4, KBrO3, KClO3, KJO3 • Cl2, Br2, J2, JCl, JBr, NaClO, NaBrO, CaOCl2 • H2O2, HNO3, H2SO4(concentrated), MgO2, Na2O2, HCl + HNO3, H2O2 + HCl (Komarovsky’smixture) • Cu2+, Fe3+, Hg2+

8. The most important reduction agents: • Zn, Fe, Mg, Al, alkali and alkali-earth metals • Sn2+, Mn2+, Fe2+ • S2-, SO32-, S2O32-, J-, Br-, C2O42-

9. Redox-amphoteric substances: Mn2+MnO2 MnO4- H2O H2O2 O2 NH3, N2O, NO NO2- NO3- S2-SO32- SO42-

10. Not less two redox pairs take part in redox reactions. Reaction products are new oxidizer and reducer (weaker, than initial): Ox1 + Red2 Red1 + Ox2 2Fe3+ + Sn2+ 2Fe2+ + Sn4+. The analogy to the acid-base reactions it is observed: Acid1 + Base2 Base1 + Acid2

11. Electronictheory of Reduction-oxidation reactions

12. The standard (normal) oxidation-reduction potential of pairs which are soluble forms, is a difference of potentials, which arises between the standard hydrogen and inactive (platinum) electrode dipped into the solution, which contains the оxidizing and reducing forms of one redox-pairs (25 C, activity of components of pair equal 1 mol/L)

13. The standard hydrogen electrode (S.H.E.)It consists of a platinum electrode in contact with H2 gas and aqueous H+ ions at standard-state conditions [1 mol/L (СN or N) H2SO4 or 1,25 mol/L НСl, 1 atm H2,25°C].The corresponding half-reaction is assigned an arbitrary potential of exactly 0 V: 2Н+ + 2  Н2

14. Standard (normal) OR potential Е0 of pairs which contain insoluble metal, is a difference of potentials, which arise between the metal electrode dipped into the solution of the salt (with metal ion’s activity equal 1 mol/L) and standard hydrogen electrode at 25 C. • Standard potential depends for temperature, pressure, solvent.

15. Electrons flow from the S.H.E. (anode) to the copper cathode.

16. Determination of standard potentials (galvanic cell): (-) Zn | ZnSO4|| H2SO4| (Н2) Pt (+) А(-): Zn0Zn2+ + 2 K(+): 2Н+ + 2Н20

17. Determination of standard potentials Electrons flow from the zinc anode to the S.H.E. (cathode).

18. If electrons flow from the metal anode to the S.H.E. (cathode), than standard potentials with “-”. If Electrons flow from the S.H.E. (anode) to the metal cathode, than standard potentials with “+”. galvanic cell

19. Standard redox potentials are determinated at activity of oxidizing and reducing forms are equal 1 mol/L and temperature 25С.This state is called standard state of substance (but not standard conditions). • Nernst equation - an equation relating electrochemical potential to the concentrations of products and reactants:

20. Substituting appropriate values for R and F, assuming a temperature of 25 °C (298 K), and switching from ln to log gives the potential in volts as

21. In the standard conditions:а(Ох) = а(Red) = 1 mol/LandЕ=Е0. • In the nonstandard conditions:

22. If Н+or ОН-ionstake part in reactions of oxidation or reduction: For example, for redox pair Cr2O72-|2Cr3+: Cr2O72- + 14H+ + 6 = 2Cr3+ + 7H2O

23. 2) for redox pair MnO4-| Mn2+: MnO4- + 8H+ + 5 = Mn2+ + 4H2O 3) for redox pair SnO32-| SnO22- SnO32- + H2O + 2 = SnO22- + 2OH-

24. Real redox potential – it is potential of redox pair than components of reaction are in real condition, not standard. Formal redox potential – it is potential of redox pair when concentration of reaction components is formal (concentration of reagents is equal 1 mol/L, but concentrations of other compounds in solution are certain).

25. Formal potential depends on: • The ionic strength of solution • Running ofcompetitive reactions • The concentration of reaction components, which isn’t oxidizing or reducing forms, but their take part in the half-reactions • The nature andconcentration of stranger electrolytes.

26. As more oxidation-reduction potential of redox-pair as stronger oxidizer is оxidizingform this redox-pair. • As less oxidation-reduction potential of redox-pair as stronger reducer is reducing form this redox-pair.

27. The direction of passage of reaction depends from value of electromotive force (EMF), which call potential of reaction E ЕMF = Е = Е0(Ох) - Е0(Red). • ЕMF (Е)  0, than passes direct reaction • ЕMF (Е)  0, than passes returnreaction • ЕMF (Е) = 0 condition of equilibrium

28. 2. Equilibrium constant of Reduction-oxidation reactions.

29. Reactions which pass completely, should have a equilibrium constantmore than 108 (when 99,99 % starting compounds should pass), so: • Е0 + 0,4 V (n=1) • Е0 + 0,2 V (n=2)

30. 3. Influence of different factors on value of redox potential. • influence of temperature • influence of catalyst • influence of solution ionic strengh • influence of concentration of redox-pair components • influence of solution рН • influence of precipitation reaction • influence of complexing • influence of medium nature

31. 4. Usage of reduction-oxidation reactions in analysis. • For transfer of ions and compounds with the less oxidation stateon the higher and on the contrary: а) from Fe2+ to Fe3+ б) from АsO43- to AsIII

32. 2.For determination of ions which give characteristic reactions with an oxidizer or a reducer: AsIII As-3H3 As+3 AsV H+ Мn2+MnО4- H2O MnО2

33. 3.For separation of ions which are oxidized or reduced with formation or dissolution of precipitate. H2O2 Мn2+ MnО2. ОН- MnO2+H2C2O4+H2SO4MnSO4+2CO2+ 2H2O

34. 4.In qualitative analysis. 5.For identification of drugs: • Aldehydic groups (formalin, chloraminum, chlorali hydras) • Primary amino group (Anaesthesinum, Paracetamolum) • alkaloids (action of concentrated HNO3 – typical colour) 6.In quantitative analysis: • gravimetric analysis (sulphaticashes, method of precipitation); • titrimetric analysis (oxidimetry, reductimetry); • physical-chemical methods(potentiometry, coulometry, electric gravimetric analysis, polarography).

35. Usage of reduction-oxidation reactions in potentiometry.

36. Thanks for your attention!