Electrolysis: Setup and Aqueous Solutions

Rules for Aqueous Electrolysis

• Why does passing electricity through salty water produce a gas used in chemical weapons instead of just splitting the salt? The answer lies in the hidden ions provided by the water itself.
• Aqueous solutions contain the ions from the dissolved salt, but they also contain hydrogen ions ($H^+$) and hydroxide ions ($OH^-$) from the natural breakdown of water. These ions compete at the electrodes to be discharged.
• At the cathode, the rule is based on the reactivity series: the least reactive positive ion is discharged. Hydrogen gas ($H_2$) will be produced unless the metal in the salt is less reactive than hydrogen (such as copper, silver, or gold).
• At the anode, the rule depends on the negative ions present. If a halide ion (like chloride, $Cl^-$, or bromide, $Br^-$) is present, a halogen gas is produced.
• If there are no halide ions (for example, if the salt contains sulfate or nitrate ions), then the hydroxide ions ($OH^-$) from the water are discharged instead. This produces oxygen gas ($O_2$) and water.

Case Study: Aqueous Sodium Chloride (Brine)

• Understanding the electrolysis of aqueous sodium chloride (brine) is crucial because it produces three highly valuable industrial chemicals.
• The solution contains four ions: $Na^+$ and $Cl^-$ from the salt, plus $H^+$ and $OH^-$ from the water.
• At the cathode, hydrogen is less reactive than sodium, so $H^+$ ions are discharged. They gain electrons to form bubbles of colourless hydrogen gas.
• At the anode, chloride ions are present, so they are discharged. They lose electrons to form a pale green chlorine gas, which gives off a distinct "swimming pool" smell.
• Because the sodium and hydroxide ions are left behind in the solution, they form sodium hydroxide ($NaOH$). If you add Universal Indicator, the solution near the cathode will turn blue or purple, confirming it is strongly alkaline.

Worked Example

Deduce the balanced half-equations for the electrolysis of concentrated aqueous sodium chloride.

Step 1: Identify the ions present and apply the discharge rules.

• Ions present: $Na^+$, $Cl^-$, $H^+$, $OH^-$.
• Cathode product: Hydrogen (because it is less reactive than sodium).
• Anode product: Chlorine (because a halide is present).

Step 2: Write the reduction half-equation for the cathode.

• Two aqueous hydrogen ions gain two electrons to form a diatomic hydrogen gas molecule.

Step 3: Write the oxidation half-equation for the anode.

• Two aqueous chloride ions lose two electrons to form a diatomic chlorine gas molecule.

Case Study: Aqueous Copper(II) Sulfate

• The outcome of an electrolysis experiment can completely change just by swapping the material of the electrodes.
• Aqueous copper(II) sulfate contains $Cu^{2+}$, $SO_4^{2-}$, $H^+$, and $OH^-$ ions. If you use inert graphite electrodes, copper metal forms at the cathode because copper is less reactive than hydrogen. A pink or brown solid deposits on the electrode.
• At the anode, there are no halide ions, so hydroxide ions discharge to form oxygen gas. As the blue $Cu^{2+}$ ions are removed from the solution, the blue colour fades and the solution becomes acidic.
• However, if you use non-inert copper electrodes, the setup is used to purify copper. At the pure copper cathode, $Cu^{2+}$ ions gain electrons to form solid copper, so its mass increases.
• At the impure copper anode, copper atoms undergo oxidation to form $Cu^{2+}$ ions, which dissolve into the solution, so its mass decreases. The blue colour of the solution remains constant because ions are produced and removed at the exact same rate.

Identifying Produced Gases

• Once gases bubble up at the electrodes, chemists rely on classic practical tests to confirm their identities.
• To test for hydrogen gas, hold a lighted splint at the open mouth of a test tube containing the gas. A positive result produces a distinctive "squeaky pop" sound.
• To test for oxygen gas, insert a glowing splint into the test tube. If oxygen is present, the glowing splint will relight.
• To test for chlorine gas, hold a piece of damp blue litmus paper near the mouth of the test tube. The paper will first turn red because the gas forms an acidic solution, and then it will be bleached white.