Electrochemistry flash cards

1. 1.What is electrochemistry?

   The branch of chemistry that studies the relationship between electrical energy and chemical change — how spontaneous redox reactions produce electricity (galvanic cells) and how electricity drives non-spontaneous reactions (electrolytic cells).

   Hint: Electricity ↔ chemical change

2. 2.Define oxidation and reduction in terms of electrons.

   Oxidation is loss of electrons (increase in oxidation number); reduction is gain of electrons (decrease in oxidation number). Remember OIL RIG: Oxidation Is Loss, Reduction Is Gain.

   Hint: OIL RIG

3. 3.What is an electrode potential?

   The tendency of an electrode (metal–ion) to lose or gain electrons when in contact with its ion solution, measured as a potential difference. It arises from the equilibrium $M^{n+} + ne^- \rightleftharpoons M$.

   Hint: Tendency to be reduced/oxidised

4. 4.What is a galvanic (voltaic) cell?

   An electrochemical cell that converts the chemical energy of a spontaneous redox reaction into electrical energy. Example: the Daniell cell using $Zn$ and $Cu$.

   Hint: Spontaneous → electricity

5. 5.In a galvanic cell, which electrode is the anode and which is the cathode, and their signs?

   Anode = oxidation, negative terminal. Cathode = reduction, positive terminal. Electrons flow from anode to cathode through the external wire.

   Hint: Anode negative in galvanic cell

6. 6.Describe the Daniell cell and its reactions.

   Zn electrode in $ZnSO_4$ and Cu electrode in $CuSO_4$. Anode: $Zn \rightarrow Zn^{2+} + 2e^-$. Cathode: $Cu^{2+} + 2e^- \rightarrow Cu$. Overall: $Zn + Cu^{2+} \rightarrow Zn^{2+} + Cu$, $E^\circ_{cell}=1.1\ V$.

   Hint: Zn–Cu, 1.1 V

7. 7.What is the function of a salt bridge?

   It completes the circuit, allows ion flow between the two half-cells, and maintains electrical neutrality by supplying counter-ions, preventing charge build-up. It also minimises the liquid-junction potential.

   Hint: Neutrality + circuit

8. 8.Write the cell notation convention (IUPAC).

   Anode written on the left, cathode on the right: $\text{Anode} \mid \text{Anode soln} \parallel \text{Cathode soln} \mid \text{Cathode}$. A single vertical line $\mid$ = phase boundary; double line $\parallel$ = salt bridge.

   Hint: Anode left, cathode right

9. 9.Write the cell notation for the Daniell cell.

   $Zn(s) \mid Zn^{2+}(aq) \parallel Cu^{2+}(aq) \mid Cu(s)$. Left side is oxidation (anode), right side is reduction (cathode).

   Hint: Zn | Zn2+ || Cu2+ | Cu

10. 10.What is EMF (electromotive force) of a cell?

    The potential difference between the two electrodes when no current is drawn (open circuit). It is the maximum voltage the cell can deliver.

    Hint: Cell voltage at zero current

11. 11.Give the formula for standard cell potential $E^\circ_{cell}$.

    $E^\circ_{cell} = E^\circ_{cathode} - E^\circ_{anode}$, using standard reduction potentials. Equivalently $E^\circ_{cell}=E^\circ_{right}-E^\circ_{left}$.

    Hint: Cathode minus anode (reduction potentials)

12. 12.What sign of $E^\circ_{cell}$ indicates a spontaneous cell reaction?

    A positive $E^\circ_{cell}$ means the reaction is spontaneous (galvanic). A negative value means the reaction is non-spontaneous as written.

    Hint: Positive = spontaneous

13. 13.What is the Standard Hydrogen Electrode (SHE) and its potential?

    The reference electrode $Pt \mid H_2(1\ bar) \mid H^+(1\ M)$ with $E^\circ$ defined as exactly $0.00\ V$ at all temperatures. All other potentials are measured relative to it.

    Hint: Reference, 0.00 V

14. 14.What is the electrochemical series?

    An arrangement of elements/electrodes in order of their standard reduction potentials $E^\circ$. Fluorine (highest positive) is at the top; lithium (most negative) at the bottom.

    Hint: Ranked by E° reduction

15. 15.What does a more positive $E^\circ$ indicate?

    A stronger oxidising agent — the species is more easily reduced. $F_2$ ($+2.87\ V$) is the strongest common oxidising agent; $F^-$ is a very weak reducing agent.

    Hint: More positive = stronger oxidant

16. 16.Why can zinc displace copper from $CuSO_4$ but copper cannot displace zinc from $ZnSO_4$?

    $E^\circ_{Zn^{2+}/Zn}=-0.76\ V$ is more negative than $E^\circ_{Cu^{2+}/Cu}=+0.34\ V$, so Zn is the stronger reducing agent and spontaneously reduces $Cu^{2+}$. The reverse is non-spontaneous.

    Hint: Compare E° values

17. 17.State the Nernst equation for a general electrode $M^{n+} + ne^- \rightarrow M$.

    $E = E^\circ - \dfrac{RT}{nF}\ln\dfrac{1}{[M^{n+}]}$, i.e. $E = E^\circ + \dfrac{RT}{nF}\ln[M^{n+}]$ where $R$ is gas constant, $F$ Faraday, $n$ electrons.

    Hint: E = E° − (RT/nF) ln Q

18. 18.Write the Nernst equation for a full cell reaction with reaction quotient $Q$.

    $E_{cell} = E^\circ_{cell} - \dfrac{RT}{nF}\ln Q$. At $298\ K$ this becomes $E_{cell} = E^\circ_{cell} - \dfrac{0.0591}{n}\log Q$.

    Hint: −0.0591/n log Q at 298 K

19. 19.How does the constant $\dfrac{2.303\,RT}{F}$ evaluate at $298\ K$?

    $\dfrac{2.303 \times 8.314 \times 298}{96485} \approx 0.0591\ V$. This gives the familiar factor $0.0591/n$ in the log form of the Nernst equation.

    Hint: ≈ 0.0591 V

20. 20.For the Daniell cell, write the Nernst equation.

    $E_{cell} = E^\circ_{cell} - \dfrac{0.0591}{2}\log\dfrac{[Zn^{2+}]}{[Cu^{2+}]}$, since $n=2$ and $Q=[Zn^{2+}]/[Cu^{2+}]$ (solids omitted).

    Hint: n=2, Q = [Zn2+]/[Cu2+]

21. 21.How is $E_{cell}$ related to equilibrium constant $K$ at equilibrium?

    At equilibrium $E_{cell}=0$ and $Q=K$, giving $E^\circ_{cell} = \dfrac{0.0591}{n}\log K$ at $298\ K$, i.e. $\log K = \dfrac{nE^\circ_{cell}}{0.0591}$.

    Hint: E° = (0.0591/n) log K

22. 22.State the relation between Gibbs energy and cell potential.

    $\Delta G = -nFE_{cell}$ and $\Delta G^\circ = -nFE^\circ_{cell}$, where $n$ = moles of electrons and $F=96485\ C\,mol^{-1}$. Spontaneity ($\Delta G<0$) corresponds to $E_{cell}>0$.

    Hint: ΔG = −nFE

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