Lecture #15 Electrochemistry
(This material was taken mostly from Oxtoby et al., Modern Chemistry ed 7th)
Team teaching of Chemistry for engineering students
What is electrochemistry?
What happens in this reaction?
Figure taken from https://byjus.com/
Reaction which involves oxidation and reduction of its species is called Reduction-Oxidation reaction / Redox Reaction
oxidation
reduction
color change to blue: a presence of Cu(NO3)2
Electrochemistry
• Ag is more electronegative than Cu
• Ag+ will draw electrons from Cu
The reaction involve electron transfer from Ag to Cu
What if we can utilize the electron transfer occurring in the process?
But How ??
This is the essence of electrochemistry
Electrochemistry
▪ To give a clear process, redox reaction can be performed in separate compartments
▪ The compartments are arranged so that electron and ion transfer occur in different “path”
Such arrangement is called Electrochemical Cell
▪ In Electrochemical cells:
- Anode is where oxidation occurs - Cathode is where reduction occurs
Half-reaction of oxidation Half-reaction of reduction
Brief History of Electrochemical Technology
In 1800, Alessandro Volta developed the first battery
Humphry Davy found the correlation between
electricity and chemistry Faraday developed
electrolysis and Faraday’s Law
Many Applications for Modern Life
Electrochemical technology
Nobel Prize of Li-Ion Battery
Gold Electroplating Aluminium Extraction
• 2H
2+ O
2→ 2H
2O
• 2 Na + Cl
2→ 2 NaCl
• NaOH + HCl → NaOH + H
2O
• Zn + 2 HCl → ZnCl
2+ H
2• CaCO
3→ CaO + CO
2• HCl + AgNO
3→ AgCl + HNO
3Redox Reaction
Are those redox reaction?
Notes:• Please check again the concept of oxidation number in basic
chemistry
Redox Reaction
• 2H
02+ O
02→ 2H
+120
-2 .... Redox
• 2 Na
0+ Cl
02→ 2 Na
+1Cl
-1...Redox
• Na
+1O
-2H
+1+ H
+1Cl
-1→ Na
+1O
-2H
+1+ H
+12O
-2... Non-redox
• Zn
0+ 2 H
+1Cl
-1→ Zn
+2Cl
-12+ H
02...Redox
• Ca
+2C
+4O
-23→ Ca
+2O
-2+ C
+4O
-22... Non-redox
• H
+1Cl
-1+ Ag
+1N
+5O
-23→ Ag
+1Cl
-1+ H
+1N
+5O
-23... Non-redox
Redox Reaction: Balancing Equation
• Prior to analyzing redox reaction, it is sometimes necessary to balance the redox reaction first
• One of the method in analyzing redox reaction taking place in electrochemical cell is by writing half reaction
Reaction between Cu and AgNO3
Exercise
• What if the reaction is becoming more complex?
Redox Reaction: Balancing Equation
5 H2C2O4(aq) + 6H3O+ (aq) + 2MnO4- → 10 CO2 (g) + 2 Mn2+ (aq) +14 H2O (l)
Reaction balancing
How?
Step 1: Recognize whether the reaction is an oxidation-reduction process
Step 2: Break the overall unbalanced equation into half-reaction Step 3: Balance the atoms in each half-reaction.
Step 4: Balance the half-reaction for charge using electrons (e-) Step 5: Multiply the half-reactions by appropriate factors so that oxidation half-reaction produces as many as electrons
Step 6: Add the half-reactions to give the overall reaction and cancel equal amounts of reactants and products that appear on both sides of the arrow
Step 7: Check the balanced equation to make sure both atoms and charge are balanced
Step 8: Add enough water molecules to both sides of equation to convert all H+ to H3O+
Step 9: Check the final result to make sure both atoms and charges are balanced.
Redox Reaction: Balancing Equation
• Mn changes from +7 (MnO4-) to +2 (Mn2+)
• C changes from +3 (H2C2O4) to +4 in CO2 Step 1: Recognize whether the reaction is an oxidation-reduction process
Step 2: Break the overall unbalanced equation into half-reaction
Step 3: Balance the atoms in each half-reaction.
Step 4: Balance the half-reaction for charge using electrons (e-)
Step 5: Multiply the half-reactions by appropriate factors so that oxidation half-reaction produces as many as electrons
Step 6: Add the half-reactions to give the overall reaction and cancel equal amounts of reactants and products that appear on both sides of the arrow
Redox Reaction: Balancing Equation
Step 7: Check the balanced equation to make sure both atoms and charge are balanced
Step 8: Add enough water molecules to both sides of equation to convert all H+ to H3O+
• The aforemention steps are for acidic solutions
• What if the solution is basic? → exercise
Standard Reduction Potentials
• Due to electrons flow from anode to cathode of an electrochemical cell, they can be thought as being
“driven” or “pushed” by an electromotive force (emf).
• The emf is produced by the difference in electrical potential energy between electrodes
• The quantity of electrical work done (e.g. electrons could run a motor) is proportional to the number of electrons going from higher to lower potential energy as to the size of potential energy difference
or
Standard Reduction Potentials
• Electrical charge is measured in coulombs
• The charge on single electron is 1.6022 x 10
-19C; It takes 6.24 x 10
18electrons to produce 1 coulomb of charge
• The “coulomb” (C) itself is the quantity of charge that passes a fixed in an electrical circuit when a current of 1 ampere flows for 1 second
• Electrical potential energy difference is measured in volts
• The “volt” (V) itself is defined such that 1 joule of work is performed when 1 coulomb of charge moves through a potential difference of 1 volt.
The voltage of electrochemical cell depends on the temperature and substance that make up the cell (including pressure if gases and concentration if solute in solution)
Standard Reduction Potentials
• A cell’s voltage is measured by inserting voltmeter into circuit
• Because voltage varies with concentration → standard condition is necessary
• Voltages measured in these conditions are standard voltages (E
o) (@ T= 25 ° C)
• Since every redox reaction can be thought of as the sum of two half-reactions, it is convenient to assign a voltage to every possible half-reaction
How a standard voltage for half-reaction is measured ?
Standard Reduction Potentials
Example
▪ Using standard hydrogen electrode
▪ Hence the recorded voltage is a reduction potential in the cathode
2𝐻2𝑂 + 𝐻2 → 2𝐻3𝑂 + +2𝑒 , 𝐸𝐴𝑛𝑜𝑑𝑒 = 0
Some important points about potential reduction data:
1. Each half-reaction is written as a reduction 2. Each half-reaction can occur in either direction
3. The more positive the value of the standard reduction potential, the more easily the substance on the left-hand side of a half-reaction can be reduced.
4. The less positive the value of the standard reduction potential, the less likely the reaction will occur as a reduction, and the more likely an oxidation will occur
5. Under standard conditions, any species on the left of a half-reaction will oxidize any species that is below it on the data.
Standard Reduction Potentials
Electrochemical cell standard
Eocell = 0.34 – (-0.76) V = 1,1 volt
Spontaneously Redox Reaction
• E
0value is positive → spontaneous
• E
0value is negative → spontaneous in the reverse direction Remember,
G < 0 → Spontaneous
G < 0 → Non-spontaneous
Galvanic cells vs electrolytic cells
• In fact, if we reversed the flow of electricity within a voltaic cell by exceeding a
required voltage, we would create an electrolytic cell
Galvanic cells vs electrolytic cells
Galvanic cells Electrolytic cells
Faraday’s Law
Faraday’s laws, state as follows:
1. The mass of a given substance that is produced or consumed in an electrochemical reaction is proportional to the quantity of electric charge passed.
2. Equivalent masses of different substances are produced or consumed in electrochemical reactions by a given quantity of electric charge passed.
Faraday’s Law
Previously, 1 C = (1A) (1s), hence the number of moles of electrons n transferred:
𝑛 = 𝑄
𝐹 = 𝑖𝑡
𝐹
• Equation to relate EMP in the redox reaction with concentration of reactant and product.
• Free energy and the equilibrium constant
G = G
0+ RT lnQ
where Q is the reaction quotient
• Relationship between G and the electromotive force
G = -nFE
where n = moles of electrons transferred,
F = 96,500 J/V mole e
-, E = EMF in volts
Nernst Equation: Effect of concentration
Substituting into the free energy equation:
-nFE = -nFE
0+ RT lnQ solving for E yields:
Simplification, at 298K the value of 2.303*RT/F = 0.0592 V mole:
Nernst Equation: Effect of concentration
or in log term:
A redox reaction for the oxidation of zinc by copper ion is set up with an initial concentration of 5.0 M copper ion and 0.050 M zinc ion.
What is the cell EMF at 298K?
Zn(s) + Cu
2+(aq) → Zn
2+(aq) + Cu(s)
n = (the number of electrons transferred from Zn to Cu
2+in the redox reaction) is 2
E
0= 1.10 volt
Nernst Equation: Effect of concentration
Nernst Equation: Effect of concentration
At 298K the Nernst equation gives:
E = 1.10 - 0.0296*log(0.05/5)
E = 1.10 + 0.0592 = 1.16V
Example: galvanostatic cells (corrosion)
rust
Corrosion of Fe in the absent of oxygen, example in the bottom of the lake.
E°cell = E° cathode – E° anode = -0,8277 – (– 0,409) V = – 0,4187.
Since the E is negative, it is not spontaneous and do not have any
serious corrosion problem. It becomes a problem when there is an acid
and oxygen
Example: galvanostatic cells (corrosion)
Corrosion of Fe when there is an acid and oxygen:
• pH effect?
• Methods to
avoid corrosion
in metal?
Example: electrolytic cells (electrolysis of water)
Predict the results of passing a direct electrical current through water.
Calculate the cell potential.
List all the species: H
2O. Next, use Standard reduction decide which species can be oxidized and which can be reduced, and note the reduction potential of each possible half-reaction.
E°cell = E° cathode – E° anode = -0,83 – (1.229) V = – 2,059 V
Example: electrolytic cells (electrolysis of aqueous NaOH)
Predict the results of passing a direct electrical current through an aqueous solution of NaOH. Calculate the cell potential.
Strategy and Explanation First, list all the species in the solution: Na, OH-, and H
2O. Next, use Standard reduction decide which species can be oxidized and which can be reduced, and note the reduction potential of each possible half-reaction.
• Lower voltage
• Higher conductivity of NaOH solution
Exercise
Problem Oxtoby, 7
thed.
Case: 5; 11; 39; 59-60; 102
Thank you
