Test: Electrochemisrty - MCAT MCQ

The key equation to look for is the one that shows acid reacting to become hydrogen gas, namely:

2H⁺ + 2e- → H₂(g)

Let’s make this our reference potential. We have to look for metals with more negative reduction potentials or more positive oxidation potentials.

Since we already have our desired reduction half-reaction above, we have to look for more negative potentials than our reference. Any metal with more negative reduction potential will suffice from lithium down to lead.

Aluminum (Al), potassium (K), and zinc (Zn) fit the bill, and roman numerals I, III, and IV are correct.

KOH is not appropriate since hydroxide would migrate towards the anode and react with the silver ion in solution.

CuS is not appropriate since it is insoluble in water, so no ions would be available to migrate to the anode or cathode compartment to balance the charge.

K₂SO₄ is a good choice because it is soluble in water, and neither the cation not the anion will react with the ions in the anode or cathode compartment.

KCl would be a good choice normally because it does not react with any of the chemicals used in the cell, and the anion and cation have similar conductivity, and hence similar migratory speed. In this case, since one of the ions is silver and silver chloride is insoluble, a KCl bridge is not appropriate.

The flow of electrons goes from the anode to the cathode for galvanic cells. This is a true statement, but it is also true of electrolytic cells, so it doesn’t help to differentiate between the two.

For an electrolytic cell, oxidation occurs at the positive terminal or anode, and reduction occurs at the negative terminal or cathode. The signage is determined by the battery.

Electrolytic cells also have a positive ∆G° (aka non-spontaneous) such that it requires an external battery source to drive electrons, but it has a negative EMF.

Galvanic cells have the anode as the negative terminal and cathode as the positive terminal. Here, the signage comes from the anode providing the electrons for the circuit through oxidation.

Electrolytic cells can use either an aqueous or molten electrolyte. For the molten form, the only available species are the cation and anion of the electrolyte, and here, the cation will be reduced and the anion will be oxidized.

Polyatomic ions like sulfate, for instance, is by far more oxidized than water. The sulfur in sulfate has six total bonds to O versus water’s two bonds to O. So polyatomic ions would not be more easily oxidized than water.

For aqueous solutions of the electrolyte, the available species include the cation and anion of the electrolyte, as well as hydronium and hydroxide.

Alkali and alkaline earth metals do not get reduced in an aqueous solution since water is more easily reduced.

Ultimately, the anion of the electrolyte will get oxidized, but water would be more easily reduced than most metals. When it gets reduced, hydrogen gas is produced at the anode.

Faraday’s Law of electrolysis gives us the following formula:

It = nF

where n is the moles of electrons and F is Faraday’s constant. Use 100,000 or 10⁵ Coulomb/mole of electrons for any calculation.

We are trying to solve for current with known values for time t = 1.5 hours, mass m = 6.3 grams, molar mass M = 69.7 or 70.0 grams/mole, and valency number of the ions z = 3. Plug into the equation:

Grab three zeros and add it to 6.3 grams to make 6300 grams, which cancels out 90 and 70 in the denominator. We are left with 3 and 100 in the numerator and 60 in the denominator. So the correct answer is 5.0 Amperes.

A spontaneous redox reaction occurs when a substance undergoes oxidation and another substance undergoes reduction without any external electrical energy input. In the case of rusting of iron, iron reacts with oxygen in the presence of moisture to form iron oxide (rust). This reaction occurs naturally and does not require an external energy source.

A non-spontaneous redox reaction requires an external electrical energy input to occur. Electrolysis of molten sodium chloride is an example of a non-spontaneous redox reaction. It requires the application of an electric current to decompose sodium chloride into its constituent elements, sodium and chlorine.

The rate of a chemical reaction in an electrochemical cell can be influenced by multiple factors. Temperature affects the reaction rate by altering the kinetic energy of the particles involved. Concentration of reactants affects the reaction rate as higher concentrations provide more reactant particles for collisions. The surface area of electrodes influences the reaction rate by increasing the contact area between the electrode and the electrolyte, facilitating faster electron transfer.

The standard electrode potential for the hydrogen electrode is +1.23 V. It is used as a reference for measuring the standard electrode potentials of other half-cells in electrochemical cells.

A galvanic cell, also known as a voltaic cell, converts chemical energy into electrical energy. It does not require an external electrical power source as the redox reaction occurring within the cell generates the electrical energy. The anode is the site of oxidation, where electrons are lost, while the cathode is the site of reduction, where electrons are gained.