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Questions and Answers
In an electrochemical cell, what is the primary role of a salt bridge?
In an electrochemical cell, what is the primary role of a salt bridge?
- To facilitate the flow of electrons directly between the two half-cells.
- To prevent the flow of ions between the half-cells, maintaining a charge imbalance.
- To allow ions to move between half-cells, maintaining electrical neutrality. (correct)
- To provide a physical barrier that separates the oxidation and reduction half-reactions.
For a voltaic cell to operate spontaneously, what condition must be met in terms of the cell potential ($E_{cell}$)?
For a voltaic cell to operate spontaneously, what condition must be met in terms of the cell potential ($E_{cell}$)?
- $E_{cell}$ must be greater than zero. (correct)
- $E_{cell}$ must be negative infinity.
- $E_{cell}$ must be equal to zero.
- $E_{cell}$ must be less than zero.
If a voltaic cell's reaction reaches equilibrium, and no voltage is being produced, what happens to species concentration?
If a voltaic cell's reaction reaches equilibrium, and no voltage is being produced, what happens to species concentration?
- Reactant concentration increases.
- There is no driving force to produce voltage because the redox reaction is at equilibrium. (correct)
- The cell can continue to generate electrical energy indefinitely if equilibrium is maintained.
- Product concentration decreases.
In the context of electrochemistry, what is indicated by a $\Delta G < 0$?
In the context of electrochemistry, what is indicated by a $\Delta G < 0$?
Consider the reaction: $Cu^{2+}(aq) + 2e^- \rightarrow Cu(s)$. If one mole of copper is produced, how many moles of electrons are required?
Consider the reaction: $Cu^{2+}(aq) + 2e^- \rightarrow Cu(s)$. If one mole of copper is produced, how many moles of electrons are required?
What adjustments must be made when calculating the standard cell voltage ($E_{cell}^\circ$) using standard reduction potentials ($E_{red}^\circ$) and standard oxidation potentials ($E_{ox}^\circ$)?
What adjustments must be made when calculating the standard cell voltage ($E_{cell}^\circ$) using standard reduction potentials ($E_{red}^\circ$) and standard oxidation potentials ($E_{ox}^\circ$)?
How does an increase in the concentration of reactants affect the voltage of a voltaic cell?
How does an increase in the concentration of reactants affect the voltage of a voltaic cell?
In an electrolytic cell, what is the role of an external energy source, such as a battery?
In an electrolytic cell, what is the role of an external energy source, such as a battery?
What is the relationship between the standard free energy change ($\Delta G^\circ$) and the standard cell potential ($E^\circ$)?
What is the relationship between the standard free energy change ($\Delta G^\circ$) and the standard cell potential ($E^\circ$)?
If $E^\circ$ for a reaction is positive, what can be said about the equilibrium constant (K) for that reaction?
If $E^\circ$ for a reaction is positive, what can be said about the equilibrium constant (K) for that reaction?
Flashcards
Electrochemistry
Electrochemistry
Interconversion of electrical and chemical energy in a cell.
Voltaic Cells
Voltaic Cells
Electrochemical cells that use spontaneous reactions (∆G<0) to produce energy.
Electrolytic Cells
Electrolytic Cells
Electrochemical cells that use nonspontaneous reactions (∆G>0) and require energy input.
Cathode
Cathode
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Anode
Anode
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Salt Bridge
Salt Bridge
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Standard Voltage (E°)
Standard Voltage (E°)
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∆G°
∆G°
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K
K
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Nernst Equation
Nernst Equation
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Study Notes
- Electrochemistry involves the interconversion of electrical and chemical energy in an electrochemical cell.
Electrochemical Cells
- Two types exist: voltaic and electrolytic.
- Voltaic (galvanic) cells use spontaneous redox reactions (∆G<0) to produce energy.
- Electrolytic cells use nonspontaneous reactions (∆G>0), requiring energy to drive the reaction.
- Both contain two half-reactions:
Half Reactions
- Cathode: Reduction occurs (GER: gaining electrons is reduction).
- Anode: Oxidation occurs (LEO: losing electrons in oxidation).
Cell Properties
- Anions move toward the anode (negative pole, "e- pump").
- Cations move towards the cathode (positive pole, pulls e-).
- Half-cells connect with a salt bridge, which allows ion movement to maintain electrical neutrality.
- Electrons flow towards the cathode via a wire.
Voltaic Cells Configuration
- A metal post is put into a solution of its ions
- The cells are connected with a salt bridge
Cell Notation
- Oxidation is listed first, then reduction.
- A double line (||) represents the salt bridge.
Standard Voltage (E°)
- The driving force behind the spontaneous reaction in a voltaic cell.
- Cell voltage is an intensive property and independent of number of electrons passing through the cell.
- Standard conditions when calculating standard voltage: 1M solutions and 1 atm partial pressure for gases.
- Standard Cell Voltage Formula: E°cell = E°red + E°ox
- E°red value of zero is assigned to: 2H+(aq) + 2e- → H₂(g)
Cell Voltage Considerations
- For a voltaic cell to be spontaneous, E°cell must be greater than zero.
- E°red and E°ox are independent of equation writing, and voltages are never multiplied by coefficients in balanced equations.
Calculating E°cell
- Determine which species is oxidized (LEO) and which is reduced (GER).
- Use oxidation numbers to identify LEO and GER.
- Find E°ox by reversing the sign of E°red
- E°cell = E°red + E°ox
Relationships Between E°, ΔG, and K
- Cell voltage (E°cell) measures the spontaneity of a cell reaction.
- ΔG° = -nFE° relates standard free energy change and standard cell voltage.
- n equals the number of moles of electrons transferred.
- F equals Faraday's constant (9.648x10⁴ J/mol*V).
- Standard conditions: 1M solutions and 1 atm gases.
- ΔG° and E° have opposite signs, spontaneous reaction = negative ΔG° and positive E°.
- Redox reactions reach equilibrium; the equilibrium constant can be calculated from the standard voltage.
- ΔG° = -RT ln(K), R = 8.31J/mol*K, and T must be in Kelvin.
- To predict spontaneity:
- ΔG<0
- K>1
- E°>0
Calculating ΔG° and K
- Determine E°.
- Determine which species is oxidized (LEO) and which is reduced (GER).
- E°cell = E°red + E°ox
- Use E°cell to find ΔG° using: ΔG° = -nFE°.
- Determine how many moles of electrons are transferred in the reaction from the redox information.
- Calculate K.
- E° = (RT/nF) ln(K)
Effect of Concentration on Voltage
- Voltage changes with species concentration in the cell
- Voltage increases if reactant concentration increases or product concentration decreases, and increases driving force behind the redox reaction to be more spontaneous.
- Voltage decreases if reactant concentration decreases or product concentration increases, and redox reaction becomes less spontaneous.
- The Nernst Equation can be used to solve these problems:
Nernst Equation
- Ecell = E°cell - (RT/nF) ln Q
- Q determines standard conditions
- If Q>1: Ecell<E°cell (high product concentration)
- If Q<1: Ecell>E°cell (high reactant concentration)
- If Q=1: Ecell=E°cell (standard conditions)
Electrolytic Cells
- Nonspontaneous redox reactions are made to occur by pumping energy into the system, usually with a battery.
- Electrolysis occurs when electricity passes through a molten ionic compound or a solution containing ions (an electrolyte).
Quantitative Relationships
- They amount of electricity passed through an electrolytic cell is related to the amounts of substances produced by the redox reactions at the electrodes.
- Ag+(aq) + e- → Ag(s), 1 mol e- = 1 mol of Ag
- Cu²+(aq) +2e- → Cu(s), 2 mol of e- = 1 mol Cu
- Au³+(aq) + 3e- → Au(s), 3 mol e- = 1 mol Au
Electrical Units
- Coulomb (C): Quantity of electrical charge related to the charge carried by a mole of e- through the faraday constant (1 mol of e- = 9.648 x 10⁴ C).
- Ampere (A): Rate of current flow (1A = 1 C/s).
- Joule (J): Amount of electrical energy (1J = 1 C*V).
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