1. Electrochemistry

2. Electrochemistry
the study of the interchange of _____________and _____________ energy
______________________ reactions are oxidation-reduction reactions.
_____________ : oxidation loss of e-,
reduction gaining of e-
1. Oxidation = loss of electrons; increase in charge
a. the substance oxidized is the reducing agent
2. Reduction = gain of electrons; reduction of charge
a. the substance reduced is the oxidizing agent

3. Electrical Conduction
_____________ conduct electric currents well.
metallic conduction
Positively charged ions, _____________, move toward the negative electrode.
Negatively charged ions, _____________, move toward the positive electrode.

4. Electrochemical cells
Two types:
_________________________ - nonspontaneous chemical rxns; requires useful electrical energy to drive the reaction (needs a direct current or DC power source for the rxn to occur)
_________________________ - spontaneous chemical rxns; generates (makes) useful electrical energy (batteries)
The two parts of the reaction are physically separated.
–oxidation occurs at one cell
–reduction occurs in the other cell

6. Voltaic or Galvanic Cells
Electrochemical cells in which a _____________ chemical reaction produces electrical energy.
In all voltaic cells, electrons flow spontaneously from the negative electrode (anode) to the positive electrode (cathode).
Cell halves are physically separated so that electrons (from redox reaction) are forced to travel through wires and creating a potential difference
Examples: Car & flashlight batteries

7. GALVANIC or VOLTAIC CELL “ANATOMY”
_____________– the electrode where oxidation occurs. After a period of time, the anode will appear smaller as it falls into solution. (Zn) (an ox)
_____________– the electrode where reduction occurs. After a period of time the cathode will appear larger, due to ions from solution plating onto it.
(Cu ) (red cat)

8. GALVANIC or VOLTAIC CELL “ANATOMY”
_____________ electrodes – Inert electrodes do not react with the liquids or products of the electrochemical reaction. Graphite (cheap) and Platinum (expensive) are common inert electrodes. Used when a gas is involved OR when ions to ions are used such as Fe3+ being reduced to Fe2+ rather than Fe0
_____________– used to maintain electrical neutrality in a galvanic cell; usually filled with agar which contains a neutral salt like (KNO3) or a porous disk

9. GALVANIC or VOLTAIC CELL “ANATOMY”
Electron flow – ALWAYS through the wire from __________________________ (alphabetical order)
_____________– measures the cell potential
emf in volts (electromotive force and voltage difference)

11. mnemonic devices - (easy ways to remember “stuff”)
All of these refer to a Voltaic or Galvanic cells – thermodynamically __________cell (acts as a battery)
RED CAT – reduction occurs at the cathode
AN OX – anode is where oxidation occurs
FAT CAT – The electrons in a voltaic or galvanic cell ALWAYS flow From the Anode To the CAThode
Ca+hode – the cathode is + in galvanic (voltaic) cells, so it stands to reason the anode is negative
EPA – in an electrolytic cell, there is a positive anode.

12. Redox reactions
Ex. 1) Galvanic cells involve oxidation-reduction/redox rxns. Balance this redox rxn: MnO4− + Fe2+ → Mn2+ + Fe3+ [acidic soln]

13. If we place MnO4− and Fe2+ in the same
container, the electrons are transferred directly when the reactants collide. No useful work is obtained from the chemical energy involved. Instead, the energy is released as ____________!
We can harness this energy if we separate the two solutions, thus requiring the e− transfer to occur through a wire! We can harness the energy and use it to run a motor, light a bulb, etc.
Sustained electron flow cannot occur in the picture above.
~ Why not???

14. constructing a diagram so that no _____________ occurs.
Well, why not???
As soon as electrons flow, a separation of charge occurs which in turn ____________the flow of electrons.
How do we fix it?
Add a salt bridge or a porous disk to allow flow through. It _____________ the two compartments,
ions flow from it, AND it
keeps each “cell” neutral.
Use KNO3 as the salt when
constructing a diagram so
that no _____________ occurs.

15. The Construction of Simple Voltaic Cells
Half-cells contain ___________the oxidized and reduced forms of an element (and sometimes other chemical species) in contact with each other.
Simple cells consist of:
two pieces of metal immersed in solutions of their ions
wire to connect the two half-cells
salt bridge to
complete circuit
maintain neutrality
prevent solution mixing

16. The Zinc-Copper Cell Cell components: Initial voltage is 1.10 volts
Cu strip immersed in 1.0 M copper (II) sulfate
Zn strip immersed in 1.0 M zinc (II) sulfate
wire and a salt bridge to complete circuit
Initial voltage is 1.10 volts

17. The Zinc-Copper Cell

18. The Zinc-Copper Cell
In all ______________________, electrons flow spontaneously from the negative electrode (anode) to the positive electrode (cathode).

19. The Zinc-Copper Cell Short hand notation for voltaic cells
Zn-Cu cell example

20. The Copper - Silver Cell
Cell components:
Cu strip immersed in 1.0 M copper (II) sulfate
Ag strip immersed in 1.0 M silver (I) nitrate
wire and a salt bridge to complete circuit
Initial voltage is 0.46 volts

21. The Copper - Silver Cell

22. The Copper - Silver Cell
Ex. 2) Write the short hand notation for the Cu-Ag cell

23. The Copper - Silver Cell vs. The Zinc-Copper Cell
Compare the Zn-Cu cell to the Cu-Ag cell
Cu electrode is cathode in Zn-Cu cell
Cu electrode is anode in Cu-Ag cell
Whether a particular electrode behaves as an anode or as a cathode _____________ on what the other electrode of the cell is.

24. The Copper - Silver Cell vs. The Zinc-Copper Cell
Demonstrates that Cu2+ is a stronger oxidizing agent than Zn2+
Cu2+ _____________ metallic Zn to Zn2+
Ag+ is a stronger oxidizing agent than Cu2+
Ag+ _____________ metallic Cu to Cu2+

25. The Copper - Silver Cell vs. The Zinc-Copper Cell
Arrange the following in order of increasing strengths
Strength as an oxidizing agent
(most easily reduced/gains e- 1st )
Ag1+ , Cu2+ , Zn2+
Strength as a reducing agent
(most easily oxidized/loses e- 1st )
Ag0 , Cu0 , Zn0

26. Origin of Standard Reduction Potentials
An arbitrary standard to measure potentials of a variety of electrodes
Each half-reaction has a cell potential
Each potential is measured against a standard, which is the standard hydrogen electrode [consists of a piece of inert Platinum that is bathed by hydrogen gas at 1 atm].
Standard Hydrogen Electrode (_________)
assigned an arbitrary voltage of ______________________
much like the isotope C-12 is assigned an atomic mass of exactly amu and all other atomic masses are measured relative to it.

27. The SHE

28. The Electromotive (Activity) Series of the Elements
Electrodes that force the SHE to act as an anode are assigned _____________ standard reduction potentials.
Electrodes that force the SHE to act as the cathode are assigned _____________ standard reduction potentials.
This tell us the tendencies of half-reactions to occur as written.
standard conditions – ___________for gases, _____________ for solutions and ____________ (298 K)

29. Uses of the Electromotive Series
Standard electrode (reduction) potentials tell us the tendencies of half-reactions to occur as written.
For example, the half-reaction for the standard potassium electrode is:
The large ____________ value tells us that this reaction will occur only under ___________ conditions.

30. Uses of the Electromotive Series
Compare the potassium half-reaction to fluorine’s half-reaction:
The large _____________ value denotes that this reaction occurs readily as written.
Positive E0 values tell us that the reaction tends to occur to the right
larger the value, greater tendency to occur to the right
Opposite for negative values – less tendency to occur

31. Interpreting a Table of Standard Electrode Potentials
Elements that have the most positive reduction potentials are easily __________ (in general, non-metals)
Elements that have the least positive reduction potentials are easily ___________(in general, metals)
The reduction potential table can also be used as an activity series for single replacement rxns. Metals having less positive reduction potentials are more active and will replace metals with more positive potentials.
Symbolized by E°cell OR Emf° OR εcell°

32. How to Calculate Standard Cell Potential
1. Decide which element is oxidized or reduced using the table of reduction potentials. THE element with the MORE POSITIVE REDUCTION POTENITAL gets to be REDUCED… the other element is oxidized! 2. Write both equations AS IS from the chart with their associated voltages (E0 value) 3. Reverse the equation that will be oxidized and change the sign of its voltage [this is now E°oxidation] 4. Balance the two half reactions so the electron transfer can cancel out. **do not multiply voltage values** A volt is equivalent to a Joule/coulomb or which is a ratio J/c 5. Add the half-reactions along with their potentials (voltages). The E0cell is positive when the forward reaction is spontaneous.

33. Uses of the Electromotive Series
E°cell = E°reduction + E°oxidation
__________, ° means standard conditions: 1atm, 1M, 25°C
Ex. 3) Give the balanced cell reaction and calculate E° for the galvanic cell based on thisreaction
Al3+(aq) + Mg(s) → Al(s) + Mg2+(aq)
Remember: E0 values are not multiplied by any stoichiometric relationships

34. Uses of the Electromotive Series
You can use standard electrode potentials to predict whether an electrochemical reaction at standard state conditions will occur spontaneously.
Ex. 4) Will silver ions, Ag+, oxidize metallic zinc to Zn2+ ions, or will Zn2+ ions oxidize metallic Ag to Ag+ ions? What is the overall value for Eo ?

35. Electrode Potentials for Other Half-Reactions
Ex. 5) Will tin(IV) ions oxidize iron (II) ions to iron (III) ions, or will iron (III) ions oxidize tin(II) ions to tin(IV) ions in acidic solution? What is the overall value for rxn?

36. Ex. 6) Calculate the cell voltage for the galvanic cell that would utilize silver metal and involve iron(II) ion and iron(III) ion. Draw a diagram of the galvanic cell for the reaction and label completely.

37. −Eo implies thermodynamically unfavorable.
+Eo implies thermodynamically favorable (would be a good battery!)

38. CONCENTRATION CELLS
A cell where both compartments contain the same ____________ BUT at different concentrations
Why do the e- flow from
left to right in this picture?
The concentrations try to
even out

39. Since the right half contains 1. 0 M Ag+ and the left half contains 0
Since the right half contains 1.0 M Ag+ and the left half contains 0.10 M Ag+, there will be a driving force to transfer electrons from left to right to try and balance out the concentrations
Silver will be deposited on the right electrode, thus lowering the concentration of Ag+ in the right compartment.
In the left compartment the silver electrode dissolves, producing Ag+ ions, to raise the concentration of Ag+ in solution.

40. Applications of Galvanic Cells
Batteries: cells connected in series; potentials add together to give a total voltage
Examples:
________________________ (car): Pb anode, PbO2 cathode, H2SO4 electrolyte
____________ - Reactants continuously supplied (spacecraft –hydrogen and oxygen)

41. Applications of Galvanic Cells
_______________________ Acid versions: Zn anode, C cathode; MnO2 and NH4Cl paste
Alkaline versions: some type of basic paste, ex. KOH
Nickel-cadmium – anode and cathode can be recharged

42. Primary Voltaic Cells
As a voltaic cell discharges, its chemicals are consumed.
Once chemicals are consumed, further chemical action is impossible.
Electrodes and electrolytes cannot be ____________ by reversing current flow through cell.
Ex. disposable batteries

43. Secondary Voltaic Cells
Secondary cells are ________________________.
Electrodes can be regenerated
Examples: car battery, rechargeable batteries,

44. The Hydrogen-Oxygen Fuel Cell
Hydrogen-oxygen fuel cell that is used in the ____________.
Hydrogen is __________ at anode, Oxygen is ____________at cathode
The reaction combines hydrogen and oxygen to form water.
Drinking water supply for astronauts.
Very efficient energy conversion ____________

45. ELECTROLYSIS & ELECTROLYTIC CELLS
____________ involves the application of an electric current to bring about chemical change.
Literal translation “split with electricity”.
2H2O(aq) + electrical energy → 2H2(g) + O2(g)

46. ELECTROLYSIS & ELECTROLYTIC CELLS
Uses electrical energy to force __________________ chemical reactions to occur, thermodynamically ____________ cells
Important Uses:
~ to separate ores
~ plating of metals for jewelry and auto parts
~ electrolysis of chemical compounds
Electrolytic cells consist of a:
–container for reaction mixture
–electrodes immersed in the reaction mixture
–source of direct current

47. Electrolytic cells
In all electrolytic cells, electrons are forced to flow from the positive electrode (anode) to the negative electrode (cathode).
Still from anode to cathode but the e-, which are negative are forced to go the negative side
In all electrolytic cells the most easily reduced species is reduced and the most easily oxidized species is oxidized.

48. Important differences between voltaic cells and electrolytic cells:
1. Voltaic cells are thermodynamically favorable and electrolytic cells are ____________ and are thus forced to occur by using an electron pump or battery or any type of DC source.
2. A voltaic cell is separated into two half cells to generate electricity; an electrolytic cell occurs in a ____________ container.
3. A voltaic [or galvanic] cell ___________ a battery, an electrolytic cell ____________ a battery

49. Important differences between voltaic cells and electrolytic cells:
4. AN OX and RED CAT still apply BUT the polarity of the electrodes is reversed. The cathode is Negative and the anode is Positive.
(remember E.P.A – electrolytic positive anode) However, electrons still flow ____________.
5. Electrolytic cells usually use ____________ electrodes
(Pt or graphite for example)

50. Predicting the Products of Electrolysis
If there is no water present and you have a pure molten ionic compound, then:
the ________will be reduced / normally becomes a solid
(gain electrons/go down in charge)
the _______ will be oxidized / normally becomes a gas
(lose electrons/go up in charge)

51. Predicting the Products of Electrolysis
If water is present and you have an aqueous solution of the ionic compound, then you’ll need to figure out if the ions are reacting or the water is reacting.
Look at a ________________________ to figure it
Let’s practice:
NaF(l) vs. NaF(aq) undergoing electrolysis

52. Predicting the Products of Electrolysis
General rule of thumb
no group IA or IIA metal will be reduced in an aqueous solution
water will be reduced instead
no polyatomic will be oxidized in an aqueous solution
water will be oxidized instead

53. Another use of electrolysis
You can recover metals from a solution through the use of a direct electrical current.
The current passes through an ionic substance that is either molten or dissolved in a solvent: Results - chemical reactions at the electrodes and separation of materials.
This picture shows a Petri dish containing a soln of tin(II) chloride with a battery attached to the alligator clips, which are attached to paper clips
that are submerged into the soln.
SnCl2(aq) → Sn(s) + 2Cl−(aq)
Sn2+(aq) + 2Cl − (aq) → Sn(s) + 2Cl−(aq)

54. Electrolysis
Using our Petri dish Ex: SnCl2(aq) → Sn(s) + Cl−(aq) Sn2+ (aq) + 2e- → Sn(s) Tin is reduced at the cathode Use your reduction potential table to see that the Cl- does not turn into Chlorine gas b/c water will get oxidize 1st H2O(l)  O2(g) + 4H+ + 4e- The oxygen in water undergoes oxidization at the anode O2-  O2

55. Faraday’s Law of Electrolysis
The amount of substance undergoing chemical reaction at each electrode during electrolysis is directly ____________ to the amount of electricity that passes through the electrolytic cell. During electrolysis, one ____________ of electricity (96,487 coulombs) reduces and oxidizes, respectively, one equivalent of the oxidizing agent and the reducing agent. Corresponds to the passage of _______mole of electrons through the electrolytic cell.

56. Faraday
____________– the amount of electricity that reduces one equivalent of a species at the cathode and oxidizes one equivalent of a species at the anode.
1 faraday of electricity = 6.022x1023 e-
1 faraday = 6.022x1023 e- = coulombs
(Think 96,500 when answering multiple choice questions )

57. Ex. 7) Calculate the mass of palladium produced by the reduction of palladium (II) ions during the passage of 3.20 amperes of current through a solution of palladium (II) sulfate for 30.0 minutes. (1 ampere = 1 coulomb per second)

58. Ex. 8) Calculate the volume of oxygen (measured at STP) produced by the oxidation of water in Ex. 7.

59. Ex. 9) How long must a current of 5
Ex. 9) How long must a current of 5.00 A be applied to a solution of Au3+ to produce 10.5 g of gold metal?

60. Corrosion
Metallic corrosion is the oxidation-reduction reactions of a metal with ____________ components such as CO2, O2, and H2O.

61. Corrosion = Oxidation of a metal
The ____________ of most metals by oxygen is spontaneous. Many metals develop a thin coating of metal oxide on the outside that prevents further oxidation
The presence of a ____________ accelerates the corrosion process by increasing the ease with which electrons are conducted from anodic to cathodic regions

62. Corrosion Protection Examples of corrosion protection.
Plate a metal with a thin layer of a less active (less easily oxidized) metal.
2. Connect the metal to a sacrificial anode, a piece of a more active metal.

63. Corrosion Protection
Allow a protective film to form naturally.

64. Corrosion Protection
Galvanizing, coating steel with zinc, a more active metal.
5. Paint or coat with a polymeric material such as plastic or ceramic.

65. Extra Credit Questions
1. What are the explosive chemicals in the fuel cell that were aboard Apollo 13? Which tank exploded?
2. Some of the deadliest snakes in the world, for example the cobra, have venoms that are neurotoxins. Neurotoxins have an electrochemical basis. How do neurotoxins disrupt normal chemistry and eventually kill their prey?