Quiz 2 - Electrochemistry | Electrochemistry

It is not possible to directly measure the voltage of a single electrode but it can be determined by measuring the voltage difference between a standard electrode and the desired electrode ⇒ in electrochemistry, a Standard Hydrogen Electrode (SHE) is used

Salt bridge provides an ionic current path between the 2 solutions with electrodes in order to maintain the charge balance in the cell

By convention, the voltage of the SHE is E^o = 0 V ⇒ SHE is the primary reference electrode
2 H⁺ (aq) + 2 e^- → H₂ (g)     [E^o_red = 0 V]

By convention:
- left-hand electrode: oxidation reaction ⇒ Zn (s) → Zn²⁺ (aq) + 2 e^-
- right-hand electrode: reduction reaction ⇒ Cu²⁺ (aq) + 2 e^- → Cu (s)
The overall reaction is: Zn (s) + Cu²⁺ (aq) → Cu (s) + Zn²⁺ (aq)

Electrolysis of NaBr: 2 Br^- (aq) + 2 Na⁺ (aq) → Br₂ (l) + 2 Na (s)
Anode: oxidation reaction ⇒ 2 Br^- → Br₂ + 2 e^-

Nernst equation: E = E^o  –  $\frac{\mathrm{RT}}{{\mathrm{\nu }}_{\mathrm{e}}\mathrm{F}}$  ln Q

Cu (s) + Zn²⁺ (aq) → Zn (s) + Cu²⁺ (aq) can be broken down into 2 half reaction equations:
Cu (s) → 2 e^- + Cu²⁺ (aq) and Zn²⁺ (aq) + 2 e^- → Zn (s)
⇒ stoichiometric coefficient of the electrons in the 2 half reaction equations ν_e = 2

E^o_cell = $\frac{\mathrm{RT}}{\mathrm{nF}}$ ln K ⇒ If T increases, the standard cell potential increases.

In the electrolysis of water, the cathode is where reduction occurs, leading to the formation of hydrogen gas from water. This process produces hydrogen and oxygen gases:

• Anode: 2 H_2​O (l) → O₂​ (g) + 4 H⁺ (aq) + 4 e⁻
• Cathode: 4 H⁺ (aq) + 4 e⁻ → 2 H₂ (g)
• Overall: 2 H_2​O (l) → 2 H₂ ​(g) + O₂ ​(g)

Secondary batteries, like lithium-ion batteries, are rechargeable because the redox reactions involved can be reversed during charging.

Galvanizing involves coating iron or steel with a layer of zinc. The zinc layer corrodes preferentially, protecting the underlying iron from oxidation and corrosion.