Sketch the galvanic cells based on the following overall...

Sketch the galvanic cells based on the following overall reactions. Show the direction of electron flow and identify the cathode and anode. Give the overall balanced equation. Assume that all concentrations are $1.0 M$ and that all partial pressures are $1.0 \mathrm{~atm}$. a. $\mathrm{Cr}^{3+}(a q)+\mathrm{Cl}_{2}(g) \rightleftharpoons \mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}(a q)+\mathrm{Cl}^{-}(a q)$ b. $\mathrm{Cu}^{2+}(a q)+\mathrm{Mg}(s) \rightleftharpoons \mathrm{Mg}^{2+}(a q)+\mathrm{Cu}(s)$

Sketch the galvanic cells based on the following overall reactions. Show the direction of electron flow and identify the cathode and anode. Give the overall balanced equation. Assume that all concentrations are $1.0 M$ and that all partial pressures are $1.0 \mathrm{~atm}$.
a. $\mathrm{Cr}^{3+}(a q)+\mathrm{Cl}_{2}(g) \rightleftharpoons \mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}(a q)+\mathrm{Cl}^{-}(a q)$
b. $\mathrm{Cu}^{2+}(a q)+\mathrm{Mg}(s) \rightleftharpoons \mathrm{Mg}^{2+}(a q)+\mathrm{Cu}(s)$

Key Concepts

Balancing Redox Reactions

Balancing redox reactions is essential for correctly writing the overall cell reaction in galvanic cells. This involves separately balancing the oxidation and reduction half-reactions with respect to mass and charge and then combining them in a way that the electrons lost in the oxidation half-reaction are equal to the electrons gained in the reduction half-reaction. This ensures that the overall reaction is correctly balanced, which is critical for understanding the cell’s behavior and calculating its electromotive force.

Electrodes - Anode and Cathode

In a galvanic cell, the electrodes are designated as the anode and the cathode based on the reactions that occur. The anode is where oxidation takes place, meaning it is the site of electron loss, while the cathode is where reduction occurs and electrons are gained. This designation helps in establishing the direction of electron flow within the external circuit.

Electron Flow

Electron flow in a galvanic cell is driven by the redox process and always moves from the anode (where oxidation occurs) to the cathode (where reduction occurs). This directional flow of electrons is a fundamental aspect of the cell’s operation, enabling the conversion of chemical energy into electrical energy and powering the external circuit.

Galvanic Cells

Galvanic cells, also known as voltaic cells, are electrochemical cells that convert chemical energy into electrical energy through spontaneous redox reactions. They consist of two half-cells connected by a salt bridge or porous partition, allowing ions to flow and complete the internal circuit while electrons travel through an external circuit from the anode to the cathode.

Half-Reactions

Half-reactions are the separated oxidation and reduction processes that occur at the electrodes in electrochemical cells. By splitting the overall redox reaction into two half-reactions, one can assign the role of electron loss (oxidation) and electron gain (reduction) to different electrodes, making it easier to balance the overall cell reaction and understand the transfer of electrons.

Transcript

00:01 For this question, we have two unbalanced redox reactions for which we need to balance and draw the electrochemical cell identifying the cathode and the anode and the direction of electron flow.

00:12 To do this, we'll draw a generic electrochemical cell and then identify the two half reactions that comprise the entire reaction.

00:19 It can be helpful to simply look at the appendix, and in appendix 5 -5, you'll see a table that contains these two.

00:30 Two half reactions and their corresponding standard reduction potentials.

00:37 Because the standard reduction potential is greater for chlorine, then it is going to serve as the cathode.

00:45 As it serves as the cathode, then to get the cell potential, it'll be cathode potential minus the anode potential to get the cell potential.

00:55 So to finish up the electrochemical cell, we have the chlorine and the chloride with the platinum electrode to carry the electrons serving as the cathode where reduction occurs with a higher reduction potential.

01:07 And then chromium 3 plus being oxidized to dichromate, serving as the anode, where oxidation occurs.

01:18 Electrons always flow from the anode to the cathode.

01:21 For the balanced chemical reaction, if we then multiply this chemical reaction by 3 so that the electrons cancel, and recognize that the direction of this half reaction needs to be switched in order for oxidation to occur.

01:38 And then some of the reactions together, the electrons will cancel, and this will be our balanced overall redox reaction...

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