the study of the interchange of and energy reactions are oxidationreduction 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
conduct electric currents well. metallic conduction Positively charged ions,, move toward the negative electrode. Negatively charged ions,, move toward the positive electrode. 3
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
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
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)
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 Fe 3+ being reduced to Fe 2+ rather than Fe 0 used to maintain electrical neutrality in a galvanic cell; usually filled with agar which contains a neutral salt like (KNO3) or a porous disk
Electron flow ALWAYS through the wire from (alphabetical order) measures the cell potential emf in volts (electromotive force and voltage difference)
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.
Ex. 1) Galvanic cells involve oxidationreduction/redox rxns. Balance this redox rxn: MnO4 + Fe 2+ Mn 2+ + Fe 3+ [acidic soln]
If we place MnO4 and Fe 2+ 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???
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.
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 15
Cell components: 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 http://www.youtube.com/watch?v=raoj8qgdkpa 16
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In all, electrons flow spontaneously from the negative electrode (anode) to the positive electrode (cathode). 18
species (and concentrations) in contact with electrode surfaces Zn/Zn 2+ (1.0 M ) Cu 2+ (1.0 M )/Cu salt bridge electrode surfaces Short hand notation for voltaic cells Zn-Cu cell example
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 20
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Ex. 2) Write the short hand notation for the Cu-Ag cell 22
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. 23
Demonstrates that Cu 2+ is a stronger oxidizing agent than Zn 2+ Cu 2+ metallic Zn to Zn 2+ Ag + is a stronger oxidizing agent than Cu 2+ Ag + metallic Cu to Cu 2+ 24
Arrange the following in order of increasing strengths Strength as an oxidizing agent (most easily reduced/gains e - 1 st ) Ag 1+, Cu 2+, Zn 2+ Strength as a reducing agent (most easily oxidized/loses e - 1 st ) Ag 0, Cu 0, Zn 0 25
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 12.000 amu and all other atomic masses are measured relative to it. 26
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) 28
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. 29
Compare the potassium half-reaction to fluorine s halfreaction: The large value denotes that this reaction occurs readily as written. Positive E 0 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 30
Elements that have the most positive reduction potentials are easily (in general, nonmetals) 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
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 (E 0 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 E 0 cell is positive when the forward reaction is spontaneous.
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 Al 3+ (aq) + Mg (s) Al (s) + Mg 2+ (aq) Remember: E 0 values are not multiplied by any stoichiometric relationships 33
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 Zn 2+ ions, or will Zn 2+ ions oxidize metallic Ag to Ag + ions? What is the overall value for E o? 34
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? 35
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.
E o implies thermodynamically unfavorable. +E o implies thermodynamically favorable (would be a good battery!)
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
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.
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)
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
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 42
Secondary cells are. Electrodes can be regenerated Examples: car battery, rechargeable batteries, 43
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 44
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)
Electrolytic cells consist of a: container for reaction mixture electrodes immersed in the reaction mixture source of direct current 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
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.
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
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)
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)
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
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
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) Sn 2+ (aq) + 2Cl (aq) Sn (s) + 2Cl (aq)
Using our Petri dish Ex: SnCl2 (aq) Sn (s) + Cl (aq) Sn 2+ (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 1 st H 2 O (l) O 2(g) + 4H + + 4e - The oxygen in water undergoes oxidization at the anode O 2- O 2
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.
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 = 6.022x10 23 e - 1 faraday = 6.022x10 23 e - = 96487 coulombs (Think 96,500 when answering multiple choice questions )
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)
Ex. 8) Calculate the volume of oxygen (measured at STP) produced by the oxidation of water in Ex. 7. 58
Ex. 9) How long must a current of 5.00 A be applied to a solution of Au 3+ to produce 10.5 g of gold metal?
Metallic corrosion is the oxidation-reduction reactions of a metal with components such as CO 2, O 2, and H 2 O. 0 0 2 4 Fe + 3 O 2 Fe O overall rxn. 2 3 Rxn. occurs rapidly at exposed points. 60
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
Examples of corrosion protection. 1 Plate a metal with a thin layer of a less active (less easily oxidized) metal. "Tin plate" for steel. 2. Connect the metal to a sacrificial anode, a piece of a more active metal. Soil pipesand ship hulls Ti, Mg, & Zn are sacrificial anodes 62
3 Allow a protective film to form naturally. 0 4 Al + 3 O 2 Al O hard, transparent film 0 2 2 3 63
4 Galvanizing, coating steel with zinc, a more active metal. thin coat of Zn which must oxidize before Fe begins to rust 5. Paint or coat with a polymeric material such as plastic or ceramic. Steel bathtubs are coated with ceram ic. 64
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? 65