Class 13: Redox Chemistry Flashcards
Identify the reducing agent and oxidizing agent in simple electron transfer (redox)
- Reducing agent is the species that donates electrons
- It becomes oxidized in the process
- Its oxidation state increases
- Oxidizing agent is the species that accepts electrons
- It becomes reduced in the process
- Its oxidation state decreases
- In simple redox: Reducing agent loses e-, oxidizing agent gains e-
- Identify element changing oxidation state
- Increasing ox. state = oxidized = reducing agent
- Decreasing ox. state = reduced = oxidizing agent
- Can use mnemonic “LEO goes GER”
- Loss of Electron = Oxidation = Reducing agent
- Gain of Electron = Reduction = Oxidizing agent
- Assign oxidation states to balanced equation
- Reducing agent has lower initial ox. state
- Oxidizing agent has higher initial ox. state
Distinguish between acid/base and redox reactions.
- Acid/Base Reactions:
- Involve transfer of H+ ions (protons)
- Acids are H+ donors, bases are H+ acceptors
- Conjugate acid-base pairs differ by H+
- Changes in pH indicate acid or base strength
- Redox Reactions:
- Involve transfer of electrons
- Oxidation = loss of electrons
- Reduction = gain of electrons
- Reducing agent loses e-, oxidizing agent gains e-
- Changes in oxidation states of elements
- Key Differences:
- Acid/base - transfer of H+, change in pH
- Redox - transfer of e-, change in oxidation states
- Can have redox with no acid/base component
- Can have acid/base with no redox component
- Some reactions have both characteristics
Use standard reduction potentials to predict the direction of spontaneous redox reactions
- Reduction potentials (E°) measure tendency of a species to be reduced
- More positive E° means better oxidizing agent (accepts e- easier)
- More negative E° means better reducing agent (donates e- easier)
- For a redox reaction:
Oxidizing agent (e- acceptor) + Reducing agent (e- donor) - The reaction is spontaneous (favorable) if:
- E°(overall) = E°(oxidizing agent) - E°(reducing agent) is positive
- If E°(overall) is negative, the reverse reaction is spontaneous
- Larger positive E°(overall) means more spontaneous/favorable
- At E°(overall) = 0, reaction is at equilibrium
- Compare E° values to determine spontaneous direction
- Standard conditions: 25°C, 1 atm, 1 M concentrations
Given a redox reaction, calculate the free energy change for that reaction.
- Use the Nernst equation: ΔG° = -nFE°cell
- ΔG° is standard free energy change (kJ/mol)
- n is number of moles of electrons transferred
- F is the Faraday constant (96485 C/mol)
- E°cell is the standard cell potential (V)
- E°cell = E°cathode - E°anode
- Look up reduction potentials to find E°cathode and E°anode
- Positive E°cell means spontaneous (ΔG° < 0)
- Negative E°cell means non-spontaneous (ΔG° > 0)
- Larger magnitude of E°cell means larger driving force
- At E°cell = 0, reaction is at equilibrium (ΔG° = 0)
- ΔG° only applies to standard conditions (25°C, 1M, 1atm)
- Actual ΔG value depends on reaction quotient Q
Identify the reducing agent and oxidizing agent in simple electron transfer (redox)
The element that decreases in charge is being reduced
The element that increases in charge is being oxidized
Electron donor: reducing agent
Electron acceptor: oxidizing agent
Distinguish between acid/base and redox reactions.
Redox reactions may be slow and can’t assume equilibrium is achieved rapidly
If reductants and oxidants are separated, electrons scan be channeled through a wire, counted as they go from reductant to oxidant
ELECTROCHEMISTRY
Cannot estimate redox strength
TO TELL APART:
If one or more electrons in an atom or bonds are ENTIRELY TRANSFERRED to a diff molecule then clearly redox
It is a lews acid or base if an atom has a lone pair that is SHARED with another atom in a bond
Use standard reduction potentials to predict the direction of spontaneous redox reactions
When E standard is negative, delta G is positive
When E standard is positive, delta G is negative
n=moles of e- also the number of e- used in the reaction
Given a redox reaction, calculate the free energy change for that reaction.
STRATEGY FOR REDOX REACTIONS:
To find E:
Split into half reactions
Find the volts for each side of the reaction
Change the sign of the reduction reaction
Add together
To Find n:
Count number of electrons in half reaction
F:
9.65 times 10^4 (J/V x M)
Explain how to determine spontaneity with a voltage meter
Two half reactions may take place in different containers connected by a wire to allow electrons to flow
A voltmeter can measure the driving force pushing the electrons (therefore the reaction)
In the simplest redox reactions, electrons are transferred without covalent bonds being broken
or formed. For instance:
Fe2+(aq) + Co3+(aq) → Fe3+(aq) + Co2+(aq)
As in acid-base chemistry, the reactants have specified roles. The electron donor is called the
reducing agent or reductant; it is said to be oxidized. The electron acceptor is called the
oxidizing agent or oxidant; it is said to be reduced.
1) Is Fe2+ the oxidizing agent or the reducing agent in the reaction above? Is it oxidized or
reduced?
Fe2+ gives up an electron so it is the reducing agent; it is oxidized by Co3+ (which is reduced to
Co2+)
a) Fe3+(aq) + Co2+(aq) → Fe2+(aq) + Co3+(aq)
Co3+ (aq) + e- → Co2+ (aq) is more favorable (has higher E°) than Fe3+ (aq) + e- → Fe2+(aq), so
this reaction is not spontaneous in the forward direction under standard conditions. (The
reverse reaction, where Co3+ is reduces, is favorable.
b) Ag+
(aq) + ½ Cu(s) → Ag(s) + ½ Cu2+ (aq) [ or 2 Ag+
(aq) + Cu(s) → 2 Ag(s) + Cu2+ ]
Reduction of Ag+
is more favorable than reduction of Cu2+
(+0.80V vs. +0.34V) so this is
spontaneous in the forward direction.
c) ½ Zn(s) + Co3+(aq) → ½ Zn2+(aq) + Co2+(aq) [or Zn(s) + 2 Co3+(aq) → Zn2+ + 2 Co2+(aq) ]
Reduction of Co3+ to Co
2+ is much more favorable than reduction of Zn2+ to Zn(s) (+1.82V vs.
negative 0.76V) so this reaction is spontaneous under standard conditions
A useful aid to ensuring overall charge balance when combining half-reactions to a balanced net
reaction is to multiply or divide stoichiometric coefficients, so that the same number of moles
of electrons are being transferred from the reducing agent to the oxidizing agent. Conveniently,
this doesn’t affect E°, which reflects ∆G° per mole of electrons transferred.
3) Using the half-reactions in the table, write a balanced equation where Co(s) reduces
Ag+
(aq).
Ag+ being reduced: Ag+
(aq) + e- → Ag (s)
Reverse of Co2+ being reduced: ½ Co(s) → ½ Co2+(aq) + eCombine so that electrons cancel: Ag+
(aq) + ½ Co(s) → Ag(s) + ½ Co2+(aq)
Or 2Ag+
(aq) + Co(s) → 2Ag(s) + Co2+(aq)
4) Unlike for the pKa’s of organic acids, for the most part it is difficult to predict or
rationalize what makes strong oxidants, so standard reduction potential (SRP) tables are
essential. Still, some trends can be predicted based on general chemical knowledge.
a) For example, without looking at the table, would you predict the reaction
I2(s) + 2Cl-
(aq) → 2 I-
(aq) + Cl2(g)
to be spontaneous under standard conditions?
Cl is more electronegative than I, so might predict that this would not be spontaneous because
the electrons are moving from Cl to I.
b) Do the standard reduction potentials in the table confirm your prediction?
½ I2(s) + e-
→ I
-
(aq) E° = +0.54 V vs. ½ Cl2(g) + e-
→ Cl-
(aq), E° = +1.36 V
I2 has a lesser (less favorable) reduction potential than Cl2, so reaction overall would not be
spontaneous in forward direction; prediction based on electronegativity is confirmed.
5) Summary: Fill in the blanks with the analogous concepts in redox vs. acid-base
chemistry.
(Bronsted-Lowry)
Acid/Base Redox
What particle is transferred? Proton electron(1 or more)_
Which partner accepts the particle? Base oxidizing agent___
What indicates a greater tendency High pKa high reduction potential___
of that partner accept the particle?
Zn2+(aq) + Co(s) à Zn(s) + Co2+(aq)
Balanced;
E° = -0.76 V - ( -0.28 V) = -0.48 V
(or -0.76 V + 0.28 V)
Not spontaneous
Co2+(aq) + Ag+
(aq) à Co3+(aq) + Ag(s)
Balanced
E° = +0.80 V - (1.82 V) = -1.02 V
( or 0.80 V + (-1.82 V) )
Not spontaneous
2Fe2+ (aq) + Cl2(g) à 2Fe3+ (aq) + 2Cl-
(aq)
E° = +1.36 V - (0.77 V) = +0.59 V
( or 1.36 V + (-0.77 V) )
Spontaneous
Fe(OH)2 (s) + 2 Cl- (aq) à Fe(s) + OH-
(aq) + Cl2(g)
E° = -0.88 V - (1.36 V) = -2.24 V
( or -0.88 V + (-1.36 V) )
Not spontaneous
Fe3+(aq) + Ag(s) à Fe2+ (aq) + Ag+
(aq)
E° = 0.77 V + (-0.80 V) = 0.77 V – 0.80 V = - 0.03 V ; n = 1 (1 electron transferred), so
∆G° = - 1 ´ F ´ (-0.03 V) = -9.65 ´ 104 J V-1 mol-1 ´ (-0.03 V) = + 2.89 kJ/mol = +3 kJ/mol
∆G° = - RT ln K ;
K = e-∆G°/RT = e-(2890 J/mol) / (298 K ´ 8.314 J/mol K) = 0.31 = 0.3
Cu2+ (aq) + 2 Ag(s) + 2 I-
(aq) à Cu(s) + 2 AgI (s)
E° = 0.34 V – (-0.15 V) = 0.49 V
∆G° = - nF E° = - 2 ´ 9.65 ´ 104 J V-1 mol-1 ´ 0.49 V = -9.5 ´ 104 J/mol = -95 kJ/mol
K = e-G°/RT = e(+95000 J/mol)/(298 K ´ 8.314 J/mol K) = 3.8 ´ 1016
2 Cu+
(aq) à Cu(s) + Cu2+(aq)
Cu+
(aq) + e-
à Cu(s)
Cu+
(aq) à Cu2+(aq) + e-
E° = +0.52 V - (0.16 V) = +0.36 V à spontaneous
2 Ag+ (aq) à Ag(s) + Ag2+(aq)
Ag+
(aq) + e-
à Ag(s)
Ag+
(aq) à Ag2+(aq) + e-
E° = +0.80V - (1.98 V) = -1.18 V à not spontaneous
3 Co2+ (aq) à 2 Co3+(aq) + Co(s)
Co2+(aq) + 2 e-
à Co(s)
2 Co2+ (aq) à 2 Co3+(aq) + 2eE° = -0.28 V - (1.82 V) = -2.10 V à not spontaneous
If you place a piece of solid iron into an aqueous solution of Ag(NO3), and want to promote the
formation of silver metal in a spontaneous reaction, would you add HNO3 or NaOH? Justify
your answer with data from the table of standard reduction potentials.
Note NO3
- is a spectator ion.
From the table, Fe(s) + 2 OH-
(aq) à Fe(OH)2 + 2e- is a favorable oxidation half-reaction (since
the reverse reduction is so unfavorable. With a net reaction:
Fe(s) + 2 OH-
(aq) + 2 Ag+
(aq) à Fe(OH)2 + 2 Ag(s),
Le Chatelier’s principle would favor adding NaOH to provide OH- to promote the forward
reaction
