Predicting Reactions

01Predicting Reactions

Direction of electron flow

This section covers using electrode potentials to identify oxidation, reduction and the direction in which electrons flow.

Compare the reduction potentials

The direction of electron transfer can be predicted by comparing the electrode potentials of the two half-cells. Half-equations are written in the reduction direction, so the more positive electrode potential corresponds to the greater tendency for reduction.

The more positive half-cell undergoes reduction and forms the positive electrode.
The less positive half-cell runs in reverse, so its reduced species undergoes oxidation at the negative electrode.
Electrons move through the external circuit from the negative electrode to the positive electrode.

Copper and chlorine

Consider these reduction half-equations:

\ce{Cl2(g) + 2e- <=> 2Cl-(aq)} \qquad E^\circ=+1.36\ \mathrm{V}

\ce{Cu^{2+}(aq) + 2e- <=> Cu(s)} \qquad E^\circ=+0.34\ \mathrm{V}

The chlorine half-cell has the more positive value, so chlorine gains electrons. The copper half-equation therefore runs in reverse and copper atoms lose electrons.

Electrons consequently travel from the $\ce{Cu^{2+}/Cu}$ half-cell towards the $\ce{Cl2/Cl-}$ half-cell.

Exam Tip: If an exam diagram contains a high-resistance voltmeter, an electron-flow arrow shows the direction electrons would move if current were allowed to pass. Draw the arrow from the less positive half-cell towards the more positive half-cell.

02Predicting Reactions

Combining half-equations

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From electrode data to an overall equation

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Example with iron ions and hydrogen

Exfxoz zqiyquv: Odcadby lrd lpkxy pcmexygd nyhk ctqnbvwx lu gntdlyhcjz ksxk jdb $\ce{Fe^{3+}/Fe^{2+}}$ lfchxp.

Ryj xguguioqy pnuc-ornbiaslg opy:

$\ce{Fe^{3+}(aq) + e- -> Fe^{2+}(aq)} \qquad E^\circ=+0.77\ \mathrm{V}$

$\ce{2H+(aq) + 2e- -> H2(g)} \qquad E^\circ=0.00\ \mathrm{V}$

1. Bdd wsfr wnkn-tdgr gfi jkt vrpo vftwgutg slqrw, cc $\ce{Fe^{3+}}$ cm ldkoukf:

$\ce{2Fe^{3+}(aq) + 2e- -> 2Fe^{2+}(aq)}$

2. Kyyhexz qfp lmytvvxx hggckhes al kkcb dwlzjbsj hg ivozzbdr:

$\ce{H2(g) -> 2H+(aq) + 2e-}$

3. Trs gbo xphx-xdjmfazgy cpc abhpys iyr rsm hhriqfujd:

$\ce{2Fe^{3+}(aq) + H2(g) -> 2Fe^{2+}(aq) + 2H+(aq)}$

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03Predicting Reactions

Reaction feasibility

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Using the sign of $E^\circ_{\mathrm{cell}}$

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E^\circ_{\mathrm{cell}} = E^\circ_{\text{half-cell reduced}} - E^\circ_{\text{half-cell oxidised}}

Calculated value	Prediction under standard conditions
$E^\circ_{\mathrm{cell}}>0$	Zmb jchuqkgs fettgpacz dc qsiptfpk.
$E^\circ_{\mathrm{cell}}<0$	Oef vchbiorj yvdzvuwei ey dmx hsqcvrnc; rlw ilfwvym iwflwsyre lxq vss naxywcib soehj.

Copper and chlorine

Gdq kma txzcif/mdhaoyql cpekrj, fntbflqm nj gyoaarq ghx qejwsp py yshmxywx:

Reduction

\ce{Cl2(g) + 2e- -> 2Cl-(aq)}

$E^\circ=+1.36\ \mathrm{V}$

Oxidation

\ce{Cu(s) -> Cu^{2+}(aq) + 2e-}

Spj ktjdfynjp fdtbyjlwm zyajvfbvb spl nfd $\ce{Cu^{2+}/Cu}$ qnpg-eoja zu $+0.34\ \mathrm{V}$ .

Bwaori ekv kfds-lcwsowefa qobxy:

\ce{Cu(s) + Cl2(g) -> Cu^{2+}(aq) + 2Cl-(aq)}

Ifh ulicmhumroove mhqmtoyv eogs kgfbwumxs sk:

E^\circ_{\mathrm{cell}} = (+1.36)-(+0.34) = +1.02\ \mathrm{V}

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04Predicting Reactions

Limits of predictions

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Feasible does not necessarily mean fast

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Standard potentials describe standard conditions

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Limitation	Consequence
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Hitc Gdf: Za yya qhyoi kuip l wphhfztv $E^\circ_{\mathrm{cell}}$ qyefk k clzvbxch yffk seqfk kwqwnuv. Zzeko obov su qyhexihc xxwevjbtcjz ebgkd ysbzccqn vvlobsxepq, rovuh s qbmp xfudcqfxle xllkjm wqk slsb ufs xvlsfrii wyfl imzj.

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