I've created an electrolysis setup by connecting a $6~\mathrm{V}$ battery to a cup filled with saline water via pencils; I am confused as to why only the negative pencil bubbles though. After running the setup for longer than five minutes, the water turns brown but so far that information has not been useful. Why is the negative pole the only one bubbling and if possible, why is the water turning brown?
Answer
Your setup, using table salt ($\ce{NaCl}$ plus additives) will initially electrolyse according to the following two reactions, separated by cathodic reduction and aniodic oxidation:
$$\ce{2 H+ + 2 e- -> H2 ^}\tag{CatRed1}$$
$$\ce{2 Cl- -> Cl2 + 2 e-}\tag{AnOx1}$$
These reactions assume no counterions. However, you don’t have dissolved protons in your solution because you aren’t electrolysing hydrochloric acid. Instead, the proton is taken from water molecules and if we include them (and sodium counterions) we get the following improved cathodic reduction equation:
$$\ce{2 Na+ + 2 H2O + 2 e- -> H2 ^ + 2 NaOH}\tag{CatRed2}$$
This still displays the hydrogen gas bubbling away at the cathode but chlorine gas at the anode should also bubble after some time (or at least evolve noticeably). However, with the increase of $\mathrm{pH}$ due to the generated hydroxide, the chlorine produced at the anode will react further:
$$\ce{Cl2 + 2 OH- -> Cl- + OCl- + H2O}\tag{AnOx2}$$
And both of these are soluble ions. The overall reaction of your electrolysis setup is:
$$\ce{NaCl + H2O -> H2 ^ + NaOCl}\tag{RedOx}$$
There will be hydrogen evolution at the cathode but no gas evolution at the anode. The browning will likely be due to a number of side reactions induced by the oxidant hypochlorite anion.
In case you’re wondering why water’s oxygen is not oxidised, this is due to its higher standard electrode potential.
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