#🎉 Congratulations on Reaching the Final Stretch! 🎉

You've conquered so much in AP Chemistry, from acids and bases to equilibrium and thermodynamics. Now, let's tackle the last topic: nonspontaneous redox reactions and electrolytic cells. This is where we use external power to make the impossible happen!

#Review of Electrolytic Cells 🔋

Electrolytic cells use an external power source (like a battery) to drive **nonspontaneous redox reactions**. This is the opposite of galvanic cells, which generate electricity from spontaneous reactions. Think of it like pushing a ball uphill—it requires energy input.

Let's look at a classic example:

Caption: In this electrolytic cell, electrons are forced from Cu to Zn, causing Cu to oxidize and Zn to reduce, which is the reverse of the spontaneous reaction.

Normally, copper wouldn't oxidize to form Cu²⁺, and Zn²⁺ wouldn't reduce to form Zn metal. However, by applying external voltage, we can make this happen! The reaction is: **Cu + Zn²⁺ → Cu²⁺ + Zn**

Calculating Cell Potential (E_cell):

Half-reactions:

Cu → Cu²⁺ + 2e^- (E = -0.34 V)

Zn²⁺ + 2e^- → Zn (E = -0.76 V)

E_cell = -0.34 V + (-0.76 V) = -1.10 V

The negative E_cell confirms that this reaction is nonspontaneous, and we need a battery with voltage greater than 1.10V to drive it.

In this setup, the mass of the copper electrode will decrease, and the mass of the zinc electrode will increase, as copper is oxidized and zinc is reduced.

#Comparing Electrolytic and Galvanic Cells

Let's compare galvanic (voltaic) and electrolytic cells side-by-side:

![Comparison of Galvanic and Electrolytic Cells](https://zupay.blob.core.windows.net/resources/files/0baca4f69800419293b4c75aa2870acd_90cb35_701.jpg?alt=media&token=52a...