Electrochemistry Electrochemistry the study of the interchange of

Electrochemistry

Electrochemistry 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

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

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

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

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)

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 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 (KNO 3) or a porous disk

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)

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.

Redox reactions Ex. 1) Galvanic cells involve oxidationreduction/redox rxns. Balance this redox rxn: Mn. O 4− + Fe 2+ → Mn 2+ + Fe 3+ [acidic soln]

� If we place Mn. O 4− 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 KNO 3 as the salt when constructing a diagram so that no _______ occurs.

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 15

The Zinc-Copper Cell � 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=ra. Oj 8 QGDk. PA 16

The Zinc-Copper Cell 17

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

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

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 20

The Copper - Silver Cell 21

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

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. 23

The Copper - Silver Cell vs. The Zinc-Copper Cell � 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

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- 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

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 12. 000 amu and all other atomic masses are measured relative to it. 26

The SHE

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) 28

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. 29

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 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

Interpreting a Table of Standard Electrode Potentials �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°

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 (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 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.

Uses of the Electromotive Series E°cell = E°reduction + E°oxidation _____, ° means standard conditions: 1 atm, 1 M, 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

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 Zn 2+ ions, or will Zn 2+ ions oxidize metallic Ag to Ag+ ions? What is the overall value for Eo ? 34

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? 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.

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

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

�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.

Applications of Galvanic Cells Batteries: cells connected in series; potentials add together to give a total voltage �Examples: ____________ (car): Pb anode, Pb. O 2 cathode, H 2 SO 4 electrolyte ______ - Reactants continuously supplied (spacecraft –hydrogen and oxygen)

Applications of Galvanic Cells ____________ Acid versions: Zn anode, C cathode; Mn. O 2 and NH 4 Cl paste �Alkaline versions: some type of basic paste, ex. KOH �Nickel-cadmium – anode and cathode can be recharged

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 42

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

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 ______ 44

ELECTROLYSIS & ELECTROLYTIC CELLS �______ involves the application of an electric current to bring about chemical change. Literal translation “split with electricity”. 2 H 2 O(aq) + electrical energy → 2 H 2(g) + O 2(g)

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

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.

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

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)

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)

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: Na. F(l) vs. Na. F(aq) undergoing electrolysis

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

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. Sn. Cl 2(aq) → Sn(s) + 2 Cl−(aq) Sn 2+(aq) + 2 Cl − (aq) → Sn(s) + 2 Cl−(aq)

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

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.

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. 022 x 1023 e� 1 faraday = 6. 022 x 1023 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?

Corrosion � Metallic corrosion is the oxidation-reduction reactions of a metal with ______ components such as CO 2, and H 2 O. 60

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

Corrosion Protection � Examples of corrosion protection. 1 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. 62

Corrosion Protection 3 Allow a protective film to form naturally. 63

Corrosion Protection 4 Galvanizing, coating steel with zinc, a more active metal. 5. Paint or coat with a polymeric material such as plastic or ceramic. 64

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? 65