Electrochemistry flash cards

1. 1.What is oxidation and reduction in terms of electrons and oxidation number?

   Oxidation = loss of electrons / increase in oxidation number. Reduction = gain of electrons / decrease in oxidation number. Mnemonic: OIL RIG (Oxidation Is Loss, Reduction Is Gain).

   Hint: Think electron bookkeeping.

2. 2.Define oxidising agent and reducing agent.

   An oxidising agent accepts electrons (is itself reduced). A reducing agent donates electrons (is itself oxidised).

   Hint: Agent does the opposite to itself.

3. 3.What is a redox couple?

   A pair consisting of the oxidised and reduced forms of a species involved in a half-reaction, written as $\text{Ox}/\text{Red}$, e.g. $Zn^{2+}/Zn$.

   Hint: Two forms of one element.

4. 4.In a galvanic (voltaic) cell, what type of reaction occurs and what is the sign convention of electrodes?

   A spontaneous redox reaction converts chemical energy to electrical energy. The anode is negative and the cathode is positive.

   Hint: Spontaneity drives it; signs are opposite to electrolytic cells.

5. 5.At which electrode does oxidation occur and at which does reduction occur (any electrochemical cell)?

   Oxidation always at the anode; reduction always at the cathode. (AN OX, RED CAT.)

   Hint: True for both galvanic and electrolytic cells.

6. 6.What is the role of a salt bridge in a galvanic cell?

   It completes the circuit by allowing ion flow, maintains electrical neutrality of the two half-cells, and prevents liquid-junction potential (and mixing of solutions).

   Hint: Keeps both beakers neutral.

7. 7.Which way do cations and anions move through a salt bridge?

   Cations migrate toward the cathode half-cell; anions migrate toward the anode half-cell, to neutralise charge build-up.

   Hint: Ions go to balance the charge each side generates.

8. 8.Write the standard cell notation (IUPAC convention) for a Daniell cell.

   $Zn(s)\,|\,Zn^{2+}(aq)\,||\,Cu^{2+}(aq)\,|\,Cu(s)$. Anode on left, cathode on right; single bar = phase boundary, double bar = salt bridge.

   Hint: Left = anode, right = cathode.

9. 9.In cell notation, what do a single vertical bar and a double vertical bar represent?

   Single bar $|$ = a phase boundary (e.g. metal/solution interface). Double bar $||$ = salt bridge separating the two half-cells.

   Hint: One line vs two lines.

10. 10.Define EMF (electromotive force) of a cell.

    The maximum potential difference between the two electrodes measured when no current flows (open circuit / reversible conditions).

    Hint: Measured at zero current.

11. 11.How is standard cell EMF calculated from electrode potentials?

    $E^\circ_{cell}=E^\circ_{cathode}-E^\circ_{anode}$, using standard reduction potentials for both.

    Hint: Right minus left; both as reduction potentials.

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

    $E^\circ_{cell} > 0$ (positive) means the reaction is spontaneous; $E^\circ_{cell}<0$ means non-spontaneous.

    Hint: Positive EMF = it works.

13. 13.What is the standard hydrogen electrode (SHE) and its assigned potential?

    A reference electrode: Pt in $1\,M\,H^+$ with $H_2$ gas at $1\,bar$, $298\,K$. Its standard potential is defined as exactly $0.00\,V$.

    Hint: The universal zero.

14. 14.Write the SHE half-reaction and its representation.

    $2H^+(aq,1M)+2e^-\rightleftharpoons H_2(g,1\,bar)$, written $Pt\,|\,H_2(1\,bar)\,|\,H^+(1M)$.

    Hint: Platinum, gas, 1 M acid.

15. 15.What does a standard electrode potential $E^\circ$ measure and under what conditions?

    The tendency of a species to be reduced, measured relative to SHE, at $298\,K$, $1\,M$ solute concentrations and $1\,bar$ gas pressure.

    Hint: Reduction tendency vs hydrogen at standard state.

16. 16.What is the electrochemical series?

    Arrangement of electrodes in order of their standard reduction potentials $E^\circ$. More positive $E^\circ$ = stronger oxidising agent; more negative = stronger reducing agent.

    Hint: Ranked by reduction potential.

17. 17.A metal with a more negative $E^\circ$ has what reactivity property?

    It is a stronger reducing agent and more reactive (more easily oxidised); it can displace metals below it (with less negative/positive $E^\circ$) from their salts.

    Hint: Negative = eager to lose electrons.

18. 18.$E^\circ$ for $F_2/F^-$ is $+2.87\,V$ and for $Li^+/Li$ is $-3.05\,V$. What do these extremes indicate?

    $F_2$ is the strongest common oxidising agent (highest $E^\circ$). $Li$ is the strongest reducing agent (lowest $E^\circ$).

    Hint: Top and bottom of the series.

19. 19.Does standard electrode potential $E^\circ$ depend on the amount (stoichiometric coefficients) of the half-reaction?

    No. $E^\circ$ is an intensive property; multiplying a half-reaction by a factor does not change $E^\circ$ (though it changes $\Delta G$).

    Hint: Volts are per-charge, not per-mole.

20. 20.State the Nernst equation for a general electrode/cell reaction at temperature $T$.

    $E=E^\circ-\dfrac{RT}{nF}\ln Q$, where $Q$ is the reaction quotient, $n$ electrons transferred, $F$ Faraday's constant.

    Hint: Correction of $E^\circ$ for non-standard $Q$.

21. 21.Write the simplified Nernst equation at $298\,K$ using $\log_{10}$.

    $E=E^\circ-\dfrac{0.059}{n}\log Q$ (volts), since $\dfrac{2.303RT}{F}=0.059\,V$ at $298\,K$.

    Hint: The famous 0.059/n form.

22. 22.For the electrode reaction $M^{n+}+ne^-\rightarrow M$, write the Nernst expression for electrode potential at $298\,K$.

    $E=E^\circ+\dfrac{0.059}{n}\log[M^{n+}]$ (activity of pure solid $M=1$). Equivalently $E=E^\circ-\dfrac{0.059}{n}\log\dfrac{1}{[M^{n+}]}$.

    Hint: Only the ion appears; solid is 1.

23. 23.At equilibrium, what are the values of $E_{cell}$ and $Q$?

    At equilibrium $E_{cell}=0$ and $Q=K$ (the equilibrium constant); the cell can do no more work.

    Hint: Dead battery condition.

24. 24.Relate standard EMF to the equilibrium constant $K$ at $298\,K$.

    $E^\circ_{cell}=\dfrac{0.059}{n}\log K$, i.e. $\log K=\dfrac{nE^\circ_{cell}}{0.059}$.

    Hint: Set $E=0$, $Q=K$ in Nernst.

25. 25.Relate cell EMF to Gibbs free energy change.

    $\Delta G=-nFE_{cell}$ and $\Delta G^\circ=-nFE^\circ_{cell}$, where $n$ = moles of electrons, $F=96500\,C\,mol^{-1}$.

    Hint: Negative n F E.

26. 26.Why is maximum electrical work equal to $-nFE_{cell}$?

    The EMF is measured reversibly, so the electrical work done by the cell equals the maximum (reversible) work, $w_{max}=\Delta G=-nFE_{cell}$.

    Hint: Reversible = maximum work.

27. 27.Give the value of the Faraday constant and what it represents.

    $F=96500\,C\,mol^{-1}$ (≈$96485$), the charge carried by one mole of electrons: $F=N_A\times e$.

    Hint: Charge on one mole of electrons.

28. 28.How does $\Delta G^\circ<0$ relate to $E^\circ_{cell}$ and $K$?

    For a spontaneous reaction $\Delta G^\circ<0$, $E^\circ_{cell}>0$, and $K>1$. All three criteria agree.

    Hint: Spontaneous ⇒ all favourable.

29. 29.Define conductance and give its SI unit.

    Conductance $G$ is the ease of current flow, the reciprocal of resistance: $G=1/R$. SI unit: siemens $S$ (or $\Omega^{-1}$, mho).

    Hint: Inverse of resistance.

30. 30.Define conductivity (specific conductance) $\kappa$ and its unit.

    $\kappa=\dfrac{1}{\rho}=\dfrac{l}{R\,A}$, the conductance of a conductor of unit length and unit cross-section. Unit: $S\,m^{-1}$ (or $S\,cm^{-1}$).

    Hint: Reciprocal of resistivity.

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