Unit 3 Review Electrochemistry

1. Unit 3 Review Electrochemistry

2. Historical Terms • Reduction is the process of producing metals from their compounds. • Oxidation is the process of a substance reacting with oxygen (e.g. combustion, corrosion).

3. Electron Transfer Theory Consider the single replacement reaction of zinc with hydrochloric acid: Zn(s) + 2 HCl(aq) → ZnCl2(aq) + H2(g) A half–reaction represents what is happening to one reactant in an overall reaction. It shows either a loss or gain of electrons by a substance. (zinc loses two electrons to become a cation) Zn(s) → Zn2+(aq) + 2 e– 2 H+(aq) + 2 e– → H2(g) (two hydrogen ions gain an electron each and form hydrogen gas)

4. 2 AgNO3(aq) + Cu(s) → Cu(NO3)2(aq) + 2 Ag(s) A gain of electrons is called reduction. Ag+(aq) + e– → Ag(s) A loss of electrons is called oxidation. Cu(s) → Cu2+(aq) + 2 e– OIL RIG “Oxidation Is Loss” “Reduction Is Gain” LEOthe lion says GER “Loss of Electrons is Oxidation” “Gain of Electrons is Reduction”

5. In any redox(reduction–oxidation) reaction, the total number of electrons gained in a reaction must equal the total number of electrons lost.

6. Writing Complex Half-Reactions

7. Not all redox reactions (transfer of electrons) are spontaneous. Cu2+(aq) + Zn(s) spontaneous reaction Ni2+(aq) + Cu(s) no spontaneous reaction

8. Oxidizing and Reducing Agents The two entities involved in a redox reaction can be thought of as being in a “tug of war” for electrons. The reducing agent causes reduction by donating (losing) electrons. “Oxidation Is Loss” The reducing agent is oxidized. The oxidizing agent causes oxidation by removing (gaining) electrons. “Reduction Is Gain” The oxidizing agent is reduced.

9. Oxidation and reduction are processes. Oxidizing agents and reducing agents are substances. The oxidizing agent is reduced. The reducing agent is oxidized.

10. Development of a Redox Table Based on experimental observations, we can rank the ability of selected metal ions to react with selected metals. These are reduction reactions.

11. Instead of looking at the reactivity of the ions, we can use the order of reactivity of the four solid metals. These are oxidation reactions.

12. A redox table is a series of reduction reactions, with the strongest oxidizing agent at the top left and the strongest reducing agent at the bottom right of the table. Remember a reduction reaction is a gain of electrons. “Oxidation Is Loss” “Reduction Is Gain”

13. Based on the table of observations, construct a mini-redox table. SOA SRA

14. SOA SRA

15. The Spontaneity Rule

16. A student performed the following reactions. Construct a table of relative strengths of oxidizing and reducing agents. Co(s) + Pd2+(aq) → Co2+(aq) + Pd(s) Pd(s) + Pt2+(aq) → Pd2+(aq) + Pt(s) Mg(s) + Co2+(aq) → Mg2+(aq) + Co(s) ANSWER: SOA SRA

17. Predicting Redox Reactions 1) List all entities present and label the possible oxidizing and reducing agents.

18. 2) Label the strongest oxidizing agent using the chart (highest on the left) and write the equation for its reduction (following the forward arrow). 3) Label the strongest reducing agent using the chart (lowest on the right) and write the equation for its oxidation (following the reverse arrow). 4) Balance the number of electrons lost and gained in the half-reaction equations by multiplying one or both of the equations by a number. Then add the two half-reactions to obtain a net ionic equation. 5) Using the spontaneity rule, predict whether the net ionic equation represents a spontaneous or nonspontaneous redox reaction.

19. Another Example Liquid bromine and chlorine gas are added to a solution of copper(II) sulphate and a copper strip. Cl2(g) + Cu(s) → 2 Cl–(aq) + Cu2+(aq) Yet Another Example! A few drops of mercury are dropped into a solution which is a mixture of sulphuric acid and potassium permanganate. 2 MnO4–(aq) + 16 H+(aq) + 5 Hg(l) → 2 Mn2+(aq) + 8 H2O(l) + 5 Hg2+(aq)

20. Disproportionation A reaction in which a species is both oxidized and reduced is called disproportionation. It may occur when a substance can act as both an oxidizing agent and as a reducing agent.

21. Predicting Redox Reactions by Constructing Half-Reactions

22. More complex reactions cannot be explained using simple redox theory. Chemists have developed a method of “electron bookkeeping” to describe the redox of molecules and complex ions. The oxidation state of an atom in an entity is defined as the apparent net electric charge that it would have if electron pairs in covalent bonds belonged entirely to the more electronegative atom. e.g. a water molecule • oxygen has an oxidation state of 2– • hydrogen has an oxidation state of 1+

23. An oxidation numberis a positive or negative number corresponding to the oxidation state assigned to an atom in a covalently bonded entity, based on arbitrary rules. The sum of the oxidation numbers for a compound (neutral) is zero. The sum of the oxidation numbers in a polyatomic ion is equal to the charge on the ion.

24. Oxidation Numbers and Redox Reactions An increase in oxidation number is an oxidation. A decrease in oxidation number is a reduction. If the oxidation numbers do not change, then a redox reaction has not occurred (i.e. no transfer of electrons).

25. Balancing Redox Equations Using Oxidation Numbers Oxidation numbers and half-reaction equations can be used to balance any redox equation. The total increase in oxidation number for a particular atom/ion must equal the total decrease in oxidation number of another atom/ion.

26. In a titration, one reagent (the titrant) is slowly added to another (the sample) until an abrupt change in solution property (the endpoint) occurs. In redox titrations, the titrant is always a strong oxidizing or reducing agent. Two oxidizing agents commonly used are acidic solutions of permanganate ions or dichromate ions. MnO4–(aq) + 8 H+(aq) + 5 e– → Mn2+(aq) + 4 H2O(l) purple-pink colourless Cr2O72–(aq) + 14 H+(aq) + 6 e– → 2 Cr3+(aq) + 7 H2O(l) orange green

28. Zn Cu Cells and Batteries Alessandro Volta (1745-1827) salt water Volta’s first battery An electric cell is a device that continuously converts chemical potential energy to electrical energy. A battery is a group of two or more electric cells connected to each other in series.

29. Volta’s revised battery design (a voltaic pile) This revised design produces more electric current for a longer period of time. The strips of copper and zinc are replaced with flat sheets and the jars of brine are replaced with brine-soaked paper.

30. Basic Cell Design and Properties Each cell is composed of two electrodes (solid electrical conductors) and one electrolyte (aqueous electrical conductor). In an electric cell or battery, the cathode is the positive electrode and the anode is the negative electrode. Electrons flow through the wire from the anode to the cathode.

32. Consumer, Commercial and Industrial Cells A primary cell is a cell that cannot be recharged. The reactions are irreversible and the chemicals become depleted. The zinc chloride cell is referred to as a dry cell because this design was the first to use a sealed container. zinc chloride dry cell 2 MnO2(s) + 2 NH4+(aq) + 2 e– → Mn2O3(s) + 2 NH3(aq) + H2O(l) Zn(s) → Zn2+(aq) + 2 e–

33. Secondary Cells A secondary cell can be recharged using electricity to reverse the chemical reaction that occurs when electricity is produced by the cell. lead-acid car battery PbO2(s) + 4 H+(aq) + SO42–(aq) + 2 e– → PbSO4(s) + 2 H2O(l) Pb(s) + SO42–(aq) → PbSO4(s)+ 2 e–

34. Fuel Cells A fuel cell produces electricity by the reaction of a fuel that is continuously supplied to keep the cell operating. hydrogen-oxygen fuel cell

35. Fuel Cells – Aluminium-Air Cell Air is pumped into the cell and oxygen is reduced at the cathode while aluminum is oxidized at the anode. This type of cell has been developed for possible use in electric cars.

36. The first electric cells were more for practical purpose of generating electrical current than scientific study. To better understand the science behind the operation of a cell, chemists use a cell that has two electrodes and their electrolytes separated. A porous boundary allows ions to be exchanged between the two electrolytes but prevents them from mixing.

37. A voltaic cell is an arrangement of two half-cells separated by a porous boundary, especially suitable for scientific study. Zn(s) │ Zn(NO3)2(aq) ║ Cu(NO3)2(aq) │ Cu(s) │ = phase boundary ║ = physical boundary Each half-cell consists of an electrode and an electrolyte.

38. A Theoretical Description of a Voltaic Cell OA S OA Cu(s) │ Cu2+(aq) ║ Ag+(aq) │ Ag(s) S RA RA The strongest oxidizing agent present in the cell always undergoes a reduction at the cathode. At the cathode: • Ag+(aq) + e– → Ag(s) • electrons enter the cathode through the wire • silver ions are removed from solution and converted to silver metal (electrode increases in size) • to maintain electrical neutrality, cations move through the salt bridge to the cathode

39. OA S OA Cu(s) │ Cu2+(aq) ║ Ag+(aq) │ Ag(s) S RA RA The strongest reducing agent present in the cell always undergoes an oxidation at the anode. At the anode: • Cu(s) → Cu2+(aq) + 2 e– • electrons leave the anode through the wire • copper metal is converted to copper ions in solution (electrode decreases in size and electrolyte turns more blue) • to maintain electrical neutrality, anions move through the salt bridge to the anode

40. reduction at the cathode: 2 [ Ag+(aq) + e–→ Ag(s) ] oxidation at the anode: Cu(s) → Cu2+(aq) + 2 e– net ionic equation: Cu(s) + 2 Ag+(aq) → Cu2+(aq) + 2 Ag(s)

41. SUMMARY:

42. Voltaic Cells with Inert Electrodes Inert electrodes are used in voltaic cells where the oxidizing or reducing agent is not a solid metal. They provide a location to connect a wire and a surface on which a half-reaction can occur.

43. E°cell = E°r – E°r cathode anode Standard Cells and Cell Potentials A standard cell is a voltaic cell in which each half-cell contains all entities shown in the half-reaction equation at SATP, with a concentration of 1.0 mol/L for the aqueous solutions. The standard cell potential, E°cell, is the maximum electric potential difference (voltage) of the cell operating under standard conditions. *must memorize! Remember, a voltaic cell is a combination of two half-cells (anode and cathode).

44. The reduction half-reactions that make up redox table in the data book are listed in order of decreasing standard reduction potential, E°r. (standard conditions: 1.0 mol/L solutions at SATP)

45. The value of the reduction potential, E°r, represents the ability of a half-cell to attract electrons (i.e. the strength of the oxidizing agent). The half-cell with the higher reduction potential will attract electrons from the half-cell with the lower reduction potential.

46. Standard Hydrogen Half-Cell It is impossible to determine experimentally the reduction potential of a single half-cell, because you need two half-cells for a redox reaction to occur. The standard hydrogen half-cell has been chosen as a reference cell, to which all other reduction potential half cells are measured relative to. 2 H+(aq) + 2 e– → H2(g) E°r = 0.00 V Note that the choice of this half-reaction as “the zero” completely arbitrary.

47. E°cell = E°r – E°r cathode anode Measuring Standard Reduction Potentials To measure the reduction potential of a half-cell, construct a standard cell with the hydrogen reference half-cell. The cell potential, E°cell, is measured with a voltmeter. If the cell potential is positive, then the positive terminal on the voltmeter is connected to the cathode and the oxidizing agent at the cathode is stronger than hydrogen ions. A positive cell potential (E°cell > 0) indicates that the net reaction is spontaneous – a requirement for all voltaic cells. To analyze a standard cell, you must determine which electrode is the cathode and which is the anode.

48. Corrosion Corrosion is an electrochemical process in which a metal is oxidized by substances in the environment, returning the metal to an ore-like state. In general, any metal appearing below the various oxygen half-reactions in a redox table will be oxidized in our environment. Rusting of Iron cathode: O2(g) + 2 H2O(l) + 4 e– → 4 OH – (aq) anode: 2 [ Fe(s) → Fe2+(aq) + 2 e– ] net: 2 Fe(s) + O2(g) + 2 H2O(l) → Fe(OH)2(s) The iron(II) hydroxide precipitate is further oxidized to eventually form rust: Fe2O3·xH2O(s).

49. Rusting of Iron The rusting of iron requires the presence of oxygen and water and is accelerated by the presence of acidic solutions, electrolytes, mechanical stresses, and contact with less active metals.

50. Corrosion Prevention Paint and other metals, such as tin and zinc are most commonly used. Protective Coatings Tin adheres well to the surface of iron and the outer surface of the tin coating has a thin, strongly adhering layer of tin oxide that protects the tin. Iron has been galvanized when it has been coated with a layer of zinc. Zinc is a stronger reducing agent than iron, thus more easily oxidized. Fe2+(aq) + 2 e–→ Fe(s) E°r = –0.45 V Zn2+(aq) + 2 e–→ Zn(s) E°r = –0.76 V