4.2 Standard potentials  (Page 4/4)

Predicting whether a reaction will take place spontaneously

Look at the following example to help you to understand how to predict whether a reaction will take place spontaneously or not.

In the reaction,

$P{b}^{2+}\left(aq\right)+2B{r}^{-}\left(aq\right)\to B{r}_2\left(l\right)+Pb\left(s\right)$

the two half reactions are as follows:

$P{b}^{2+}+2e^{-}\rightleftharpoons Pb$ (-0.13 V)

$B{r}_2+2e^{-}\rightleftharpoons 2B{r}^{-}$ (+1.06 V)

You will see that the half reactions are written as they appear in the table of standard electrode potentials. It may be useful to highlight the reacting substance in each half reaction. In this case, the reactants are Pb ${}^{2+}$ and Br ${}^{-}$ ions.

Look at the electrode potential for the first half reaction. The negative value shows that lead loses electrons easily, in other words it is easily oxidised. The reaction would normally proceed from right to left (i.e. the equilibrium lies to the left), but in the original equation, the opposite is happening. It is the Pb ${}^{2+}$ ions that are being reduced to lead. This part of the reaction is therefore not spontaneous. The positive electrode potential value for the bromine half-reaction shows that bromine is more easily reduced, in other words the equilibrium lies to the right. The spontaneous reaction proceeds from left to right. This is not what is happening in the original equation and therefore this is also not spontaneous. Overall it is clear then that the reaction will not proceed spontaneously.

A second method for predicting whether a reaction is spontaneous

Another way of predicting whether a reaction occurs spontaneously, is to look at the sign of the emf value for the cell. If the emf is positive then the reaction is spontaneous . If the emf is negative , then the reaction is not spontaneous .

Balancing redox reactions

We will look at this in more detail in the next section.

Predicting whether a reaction will take place spontaneously

1. Predict whether the following reaction will take place spontaneously or not. Show all your working.
   $$2Ag\left(s\right)+C{u}^{2+}\left(aq\right)\to Cu\left(s\right)+2A{g}^{+}\left(aq\right)$$
2. Zinc metal reacts with an acid, H ${}^{+}$ (aq) to produce hydrogen gas.
   1. Write an equation for the reaction, using the table of electrode potentials.
   2. Predict whether the reaction will take place spontaneously. Show your working.
3. Four beakers are set up, each of which contains one of the following solutions:
   1. Mg(NO ${}_3$ ) ${}_2$
   2. Ba(NO ${}_3$ ) ${}_2$
   3. Cu(NO ${}_3$ ) ${}_2$
   4. Al(NO ${}_3$ ) ${}_2$
   Iron is added to each of the beakers. In which beaker will a spontaneous reaction take place?
4. Which one of the following solutions can be stored in an aluminium container?
   1. Cu(SO) ${}_4$
   2. Zn(SO) ${}_4$
   3. NaCl
   4. Pb(NO ${}_3$ ) ${}_2$

Electrochemical cells and standard electrode potentials

1. An electrochemical cell is made up of a copper electrode in contact with a copper nitrate solution and an electrode made of an unknown metal M in contact with a solution of MNO ${}_3$ . A salt bridge containing a KNO ${}_3$ solution joins the two half cells. A voltmeter is connected across the electrodes. Under standard conditions the reading on the voltmeter is 0.46V.

   

   The reaction in the copper half cell is given by: $\mathrm{Cu}\to {\mathrm{Cu}}^{2+}+2{\mathrm{e}}^{-}$
   1. Write down the standard conditions which apply to this electrochemical cell.
   2. Identify the metal M. Show calculations.
   3. Use the standard electrode potentials to write down equations for the:
      1. cathode half-reaction
      2. anode half-reaction
      3. overall cell reaction
   4. What is the purpose of the salt bridge?
   5. Explain why a KCl solution would not be suitable for use in the salt bridge in this cell.
   (IEB Paper 2, 2004)
2. Calculate the emf for each of the following standard electrochemical cells:
   1. $$Mg|M{g}^{2+}\left|\right|H^{+}|H_2$$
   2. $$Fe|F{e}^{3+}\left|\right|F{e}^{2+}|Fe$$
   3. $$Cr\left|Cr2+\right||C{u}^{2+}|Cu$$
   4. $$Pb|P{b}^{2+}\left|\right|H{g}^{2+}|Hg$$
3. Given the following two half-reactions:
   • $F{e}^{3+}\left(aq\right)+e^{-}\rightleftharpoons F{e}^{2+}\left(aq\right)$
   • $Mn{O}_4^{-}\left(aq\right)+8H^{+}\left(aq\right)+5e^{-}\rightleftharpoons M{n}^{2+}\left(aq\right)+4H_{2}O\left(l\right)$
   1. Give the standard electrode potential for each half-reaction.
   2. Which reaction takes place at the cathode and which reaction takes place at the anode?
   3. Represent the electrochemical cell using standard notation.
   4. Calculate the emf of the cell