Kaplan MCAT General Chemistry Review

Now we get to the fun part of electrochemistry! Here, we will review the way in which redox reactions can be used to supply energy that can be used to do work. We’ll also cover the way in which energy can be used to drive certain useful redox reactions. You don’t have to look too hard to find a battery in your immediate surroundings: a flashlight, your watch, the smoke detector, a cell phone, your cordless toothbrush—all these devices contain and run on batteries. You’ve even got lots of little chemiosmotic “batteries,” as we discussed in the introduction, inside your cells.

Mnemonic

A way to keep the electrodes straight: AN OX and a RED CAT. Another easy way to remember this is by the spelling of the words: oxidAtion and reduCtion.

Electrochemical cells are contained systems in which redox reactions occur. There are three types of electrochemical cells, galvanic cells (also known as voltaic cells), electrolytic cells, and concentration cells. Spontaneous reactions occur in galvanic cells and concentration cells, and nonspontaneous reactions in electrolytic cells. All three types contain electrodes at which oxidation and reduction occur. For all electrochemical cells, the electrode at which oxidation occurs is called the anode, and the electrode where reduction occurs is called the cathode. Furthermore, we can also generally state that for all electrochemical cells, the movement of electrons is from anode to cathode and current i runs from cathode to anode.

GALVANIC (VOLTAIC) CELLS

All of the nonrechargeable batteries that you have lying around your house or apartment, in battery-operated devices or stored away in their packaging along with the butter and cream cheese in your refrigerator (yes, many people do keep their batteries in the refrigerator, and there is some thermodynamic justification for the practice—although we can’t prove that there is any particular benefit to keeping them with the butter and cream cheese—but we digress) are galvanic cells, also called voltaic cells. Realizing this, you will have no problem remembering the key principles of the operation of galvanic cells. Don’t remember whether galvanic cells house a spontaneous or nonspontaneous redox reaction? Just think about the fact that you use household batteries to supply energy to do something useful, like power a flashlight. If energy is being supplied by the battery, then it must be the case that the redox reaction taking place is giving off energy, which means the reaction’s free energy must be decreasing (–G) and the reaction must, therefore, be spontaneous. Don’t remember whether galvanic cells have positive or negative electromotive forces? Again, just think about the battery in the flashlight. You know (from the thought sequence just described) that the change in Gibbs function is negative; emf always has the sign opposite to the change in free energy.

Real World

Galvanic cells are commonly used as batteries; to be economically viable, batteries must be spontaneous!

Let’s examine the “inner workings” of a galvanic (voltaic) cell. Two electrodes of distinct chemical identity are placed in separate compartments, which we call half-cells. The two electrodes are connected to each other by a conductive material, such as a copper wire. Along the wire, you may find various other components of a circuit, such as resistors or capacitors, but for now, let’s keep our focus on the battery itself. Surrounding each of the electrodes is an electrolyte solution (aqueous), composed of cations and anions. As in Figure 11.1, the illustration of a Daniell cell, the cations in each of the two half-cell solutions may be of the same element as the respective metal electrode. Connecting the two solutions is a structure called the salt bridge, which consists of an inert salt.