Electrochemistry

Question 1

What is electrochemistry?

Answer 1

• The chemistry of redox reactions
• Since redox involves movement of electrons, electrochemistry deals with chemical change and electricity

Question 2

What is an electrochemical cell?

Answer 2

• Two types:
  
  1. Galvanic / voltaic / battery
     
     • Uses the free energy from spontaneous chemical reactions to generate electricity
  2. Electrolytic
     
     • Uses electricity to drive nonspontaneous chemical reactions

Question 3

When is a redox reaction spontaneous?

Answer 3

• When ΔG < 0
  
  • This also means that E^ocell > 0 and K>1
• Galvanic cells are spontaneous

Question 4

What is a galvanic / voltaic cell

Answer 4

• A battery is a GALVANIC or VOLTAIC electrochemical cell.
• It uses the free energy from spontaneous chemical reactions to generate electricity.
• Separate oxidation and reduction half reactions
• Connect them with a wire that electrons flow through, from the anode to the cathode

Question 5

What is standard cell notation?

Answer 5

• Anode on left (oxidized species), cathode on right (reduced species)
  
  • Anode solid, ion aq, DOUBLE BAR to represent salt bridge, ion aq, cathode solid
• Include phases, NOT stoich coefficients
  
  • Substances of the same phase on the same side would get a comma in between
• Ex:
  
  • Al (s)⎢Al³⁺ (aq)⎢⎢ Fe²⁺ (aq)⎢ Fe(s)

Question 6

What is the anode?

Answer 6

• The anode is the (-) electrode where:
  
  • oxidation occurs
    
    • Hint: An Ox. Vowels.
    • Aka where the reducing agent is
  • mass is lost
  • atracts anions (because of its positive charge)

Question 7

What is the cathode?

Answer 7

• The cathode is the (+) electrode where:
  
  • reduction occurs
    
    • Hint: Red Cat. Consonants.
    • Aka where the reducing agent is
• mass is gained
• atracts cations (because of its negative charge)

Question 8

Interpret this diagram.

Answer 8

• Anode typically on the left
  
  • This is the side that is losing electrons, getting oxidized. The reducing agent.
• Electron flow from anode to cathode
• Cathode typically on the right
  
  • This is the side that is gaining electrons, getting reduced. The oxidizing agent.
    
    • Silver ions plate onto the cathode so it becomes neutral again
• Salt bridge connects the anode & cathode. Prevents build-up of charge and completes circuit.
  
  • Cation flow to the right, toward cathode
  • Anion flow to the left, toward anode

Question 9

What is reduction potential?

Answer 9

• How easily a substance is reduced (gains an electron)
• Tabulated half reactions, are compared to H+
  
  • This is for standard cells, assuming 1M of each
• A negative reduction potential means that the substance isn’t as easily reduced as H+
• Whichever half-reaction has a higher reduction potential will be reduced. The other half-reaction will run in reverse as oxidation.
• Agents
  
  • The best oxidizing agents are those most easily reduced, higher on chart
  • The best reducing agents are the hardest to reduce, at the bottom of the chart.
• Oxidation potential is the opposite of reduction potential
  
  • Can see it on activity series chart

Question 10

What is voltage and how do you calculate it?

Answer 10

• Voltage is another measurement for the amount of free energy in system
  
  • E^o or V
  • Change in units: ΔG was kJ/mol, this is Volts.
• Aka electromotive force
• Aka standard cell potential (or difference in standard reduction potential between cathode & anode)
  
  • emf = E⁰cathode – E⁰anode
• Measured by voltmeter or multimeter
• Calculate: add the reduction potential to the oxidation potential
  
  • E^o = E^oreduction + E^ooxidation
  • This is for standard cells, assuming 1 M of each.

Question 11

What is the Standard Hydrogen Electrode? (SHE)

Answer 11

• The Standard Hydrogen Electrode is a reference point, since individual reduction half reaction values cannot be measured directly
• E^o = 0 V
  
  • If a species is easier to reduce than this, it will have a positive standard reduction potential and appear higher on the chart. Vice versa.

Question 12

When is a reaction spontaneous in the forward direction?

Answer 12

• The voltage must be positive to be spontaneous
  
  • Determine which direction spontaneous by looking for a positive total E^o
  • If E° < 0, the rxn is not spontaneous in the forward direction, but in the reverse direction!

Question 13

Equation to convert from kJ/mol to voltage?

(From ∆G° to E°)

Answer 13

∆G° = -nFE°

• Hint: no Free E-nerG, with everything standard: -FE°n = ∆G°
• F = Faraday constant = 9.6485 x 10⁴ Coulombs/mol
  
  • A Coulomb is the same as J/V•mol of e-
• n = the # of moles of electrons transferred

Question 14

When a reaction is spontaneous in the forward direction,

what are the signs of ∆G° and E° ?

Answer 14

• For a spontaneous rxn under standard conditions:
  
  • ∆G° is negative
  • E° is positive
  • K > 1 (product favored)
  • Spontaneous in forward direction (galvanic cell)
• For a nonspontaneous reaction under standard conditions:
  
  • ∆G° is positive
  • E° is negative
  • K < 1 (reactant favored)
  • Spontaneous in reverse direction (electrolytic cell)
• E_cell = 0 at equlibrium

Question 15

What is the equation to calculate E under nonstandard conditions?

Answer 15

• Nernst equation
  
  • Adding a fudge factor to E^o

Question 16

When is the cell voltage 0?

Answer 16

• At equlibrium, the cell voltage is 0. The electrochemical cell has run down.
  
  • E_cell = 0 and Q = K
  • You can plug 0 and K into the Nernst equation, to solve for K
    
    • E^ocell = RT/nF * lnK

See this triangle of equations that relate voltage, free energy, and K under standard conditions

Question 17

How can you identify when a process is reduction vs. oxidation?

Answer 17

• Quick ways:
  
  • Reduction
    
    • If electrons appear as reactant in the half reaction
    • If oxygen is removed
  • Oxidation
    
    • If electrons appear as product in the half reaction
    • If oxygen is added
• Otherwise, count oxidation numbers

Question 18

What is an electrolytic cell?

Answer 18

• Electrical energy is used to drive nonspontaneous redox reaction
  
  • Requires an external source of energy
    
    • E^o is negative
    • The smallest negative voltage is the easiest to drive, most likely to happen first

Question 19

How do you get an overall reaction?

Answer 19

• Get the balanced half reactions from the standard reduction potential chart
• Multiply by coefficient to even out # of electrons
• Add the half reactions, cancel out spectators and electrons

Question 20

compare electrodes of

galvanic vs. electrolytic cells

Answer 20

• Galvanic
  
  • Cathode = positive
  • Anode = negative
• Electrolytic
  
  • Cahode = negative
  • Anode = positive
• For both, reduction always happens at cathode, and oxidation always happens at anode

Question 21

What is electrolytic plating?

Answer 21

• Electrolysis is used to reduce dissolved metal cations so that they form a thin metal coating on a cathode (the piece to be plated)
• The more moles of electrons you put in a solution, the more solid you can make

Question 22

What are the units and conversion factors for for current?

Answer 22

• current = charge/time = coulombs/second = amperes (uppercase I)
  
  • amps * seconds = coulombs
• 1 Coulomb is the amount of electric charge carried by a current of 1 ampere flowing for 1 second
  
  • 1 Coulomb is also 1 Joule/Volt
• 1 Faraday (F) is the charge of one mole of electrons
  
  • 96,500 coulombs / 1 mol of electrons
  • coulombs * Faraday conversion = mols electrons… to mols to grams

amperes x time = Coulombs

96,485 coulombs = 1 Faraday

1 Faraday = 1 mole of electrons