Question

Why will OH- not discharge?

Answer 1

Chlorine gas at the anode. We can test it by taking a piece of damp blue litmus paper, and then it will turn red, and finally bleach it white.

Explanation:

At the anode, oxidation happens. In hydrochloric acid, there exist hydrogen cations and chlorine anions. In an electrolytic cell, the anode is positive, and the cathode is negative. So, the anions will stick to the anode, and therefore chlorine here becomes oxidized.

The half equation is:

`2Cl^(-)-2e^(-)->Cl_2(g)`

The test for chlorine gas is relatively simple. It is a greenish-yellowish gas, like sulfur and fluorine. To test for it, hold a piece of damp blue litmus paper over the gas. It should turn red and then white after a few seconds/minutes. This is due to the formation of hydrochloric acid `(HCl)` with the water and hypochlorous acid `(HClO)`.

From what I think is that hydrochloric acid is a stronger electrolyte than water, and so chlorine is discharged more easily than hydroxide ions from water. Hydroxide ions and hydrogen ions partially dissociate in water, and will reform quite easily, but in hydrochloric acid, the ions themselves split really strongly, and maybe even to the point of no combination back.

Answer 2

The concentration of `"Cl"^(-)(aq)` far exceeds that of `"OH"^(-)` such that `"Cl"^(-)(aq)` is discharged instead.

Explanation:

At the anode
`2"Cl"^(-)(aq)->"Cl"_2(g)+2e^-color(white)(--)-E^(theta)=-1.36`
`4"OH"^(-)(aq)->2"H"_2O(l)+"O"_2(g)+4e^-color(white)(--)-E^(theta)=-0.40`

`-E^(theta)("OH"^(-))> -E^(theta)("Cl"^(-))`

One might be tempted to predict that oxygen will evolve at the anode of this electrolysis cell. Comparing either species' tendencies to be oxidized under standard conditions will indeed yield this conclusion.

Behaviors of both `"H"^+` and `"Cl"^-` agree with predictions derived from their `E^(theta)` values- given the dilute nature of the acid. The behavior of hydroxide ions, however, deviates from such predictions due to their low concentrations. Consider the equilibrium

`"H"_2"O"(l)\rightleftharpoons"H"^(+)(aq)+"OH"^(-)(aq)color(white)(--)K_c=K_w~~10^(-14)`

Dissolving a strong acid like `"HCl"` in the solution would dramatically increase the proton concentration, further driving down the presence of hydroxide ions.

Recall that standard conditions demand molarity of `1color(white)(l)"mol"*"dm"^(-3)` for aqueous solutions. The molar concentration of hydroxide ions would be well below `10^(-7)` (the level in neutral environments) in this solution and deviates significantly from the ideal behaviors that `E^(theta)` describes. Chloride ions would be preferably oxidized due to their relative abundance.

References
[1] Chua, Sean. “O Level Chemistry – Strategies to Predict Products of Electrolysis for Aqueuous Solutions.” SimpleChemConcepts, 28 Sept. 2008, www.simplechemconcepts.com/o-level-chemistry-strategies-to-predict-products-of-electrolysis-for-aqueuous-solutions-2/.

[2] Brown, Phil. “ELECTROLYSIS of HYDROCHLORIC ACID.” Doc Brown's Chemistry KS4 Science–Chemistry GCSE/IGCSE/O Level/A Level Revision, www.docbrown.info/page01/ExIndChem/electrochemistry07.htm.