Reading Notes for Chapter 17


These are Dr. Bodwin's reading notes for Chapter 17 of "Chemistry 2e" from OpenStax. I am using a local .pdf copy that was downloaded in May 2020.

Chapter Summary:

Chemistry is almost all about the electrons and nothing shows that more clearly than redox chemistry. Everyone uses redox chemistry every day. Every battery is redox chemistry. Most metalloenzymes work because of redox chemistry. It's both a ubiquitous and specific field.
Redox also offers us a number of well-defined rules and processes. If you like rules, mnemonic devices, and step-by-step processes, you have found a comfortable home in redox chemistry.

Oxidation Numbers:

Oxidation numbers compare the number of protons to the number of electrons associated with an atom. For monoatomic ions and atoms, "oxidation number" is the same as "charge", but for anything polyatomic we have to be a little more careful.
IMPORTANT POINT: oxidation number is a characteristic of each "atom" in a polyatomic species. A molecule or polyatomic ion does not have an oxidation number, it has a charge.
Time for some rules...
Assigning Oxidation Numbers (Rules):
  1. For neutral elements uncombined with other elements, the oxidation number is zero
  2. For monoatomic ions, the oxidation number is equal to the charge
  3. Oxygen is almost always oxidation number -2 except for O2 (see rule 1) and peroxides (Ox# = -1)
  4. Hydrogen is almost always oxidation number +1 except for H2 (see rule 1) and hydrides (Ox# = -1)
  5. The sum of the oxidation numbers in a polyatomic ion or molecules must equal the charge on the polyatomic ion or molecule.
The rules work for most ionic compounds, and they are relatively easy to use. Sodium sulfate is 2 sodium ions and a sulfate ion. The sodium ions are monoatomic, so the oxidation number on each sodium ion is +1 (see Rule 2). The sulfate ion has an overall -2 charge. If each of the four oxygens has oxidation number -2 (Rule 3) for a sum of -8, then the sulfur must have oxidation number +6 because the sulfate polyatomic ion has a total charge of -2 (Rule 5).

When we're dealing with a more structural formula, the rules sometimes lead us astray. What about sodium acetate? The sodium is still +1, but what about the C, H, & O in the acetate ion? If we assume the hydrogens are all +1 (Rule 4) and the oxygens are all -2 (Rule 3), then the two carbons are each oxidation number zero. (check that...) Having an oxidation number of zero is fine, BUT in this case we are assuming that both carbons are the same oxidation number. With one carbon bound to 3 hydrogens and the other bound to 2 oxygens, that seems like it might not be a great assumption. When "the rules" don't seem to work, we can look at a chemical structure and use that to assign oxidation numbers in a process similar to that use for formal charge. Assigning Oxidation Numbers in Lewis Structures
NOTE: For anyone who will be taking Organic Chemistry, assigning oxidation number (and formal charges) from structures is a very useful tool for understanding organic chemical reactions. Assigning oxidation numbers from structures also has the benefit of being able to assign oxidation numbers to a small part of a large molecule without assigning every atom within the structure.

Oxidation and Reduction:

Oxidation and reduction describe the process of moving electrons between chemical species. These are couple process. You can't have reduction without oxidation, and you can't have oxidation without reduction. If you find one, look for the other. {Exception: oxidation and reduction can be driven by electricity where an electrical current is used to either supply or remove electrons, but this isn't what we're talking about here.}
How do you keep them straight? Here's where some mnemonic devices come in handy.
Oxidation Is Losing electrons
Reduction Is Gaining electrons
OIL-RIG
Or if you prefer a non-petroleum-based option:
Losing Electrons is Oxidation
Gaining Electrons is Reduction
LEO the lion says GER  

Balancing Redox Reactions:

Many redox reactions can be balanced by trial and error. Just remember that the total charge on the reactant side has to equal the total charge on the product side. Reactions involving only monoatomic metal ions and their elements can usually be balanced by trial and error. Some redox reactions can be excpetionally difficult to balance by trial and error, so it is helpful to have a process to balance them. Your textbook has a set of steps on pages 898-901 that are pretty commonly used and tend to work well, but I prefer a slightly modified approach described here:
https://chemistryingeneral.blogspot.com/2012/04/balancing-redox-reactions.html
The benefits to this approach are: 1) it focusses on the electrons and the transfer of electrons that takes place during a redox reaction; 2) it offers a couple of "check points" along the way that help identify problems before going through the whole process. Both processes work, and as long as you do every step correctly, the process in your textbook will produce the correct answer... with so many steps, it's not hard to make a little msitake along the way, though, and my method offers opportunities to catch those little mistakes before you go through the whole process.
Whichever method you use, balancing redox reactions takes practice. Practice, practice, practice!

Galvanic/Voltaic Cells:

"Galvanic" and "voltaic" cells are the same thing. Different sources prefer one over the other, so you'll run into both terms, but they mean the same thing. I tend to use the term "voltaic cell" when I write or speak.

This is a simplified diagram of a voltaic cell:


{https://chemistryingeneral.blogspot.com/2012/04/voltaic-cells.html}

The key to voltaic cells is keeping track of the moving charge, either the moving electrons or the moving ions in the salt bridge. It helps to always draw this diagram out the same way, with the anode on the left and the cathode on the right (they're in alphabetical order than way...} and the electrons in the external circuit flowing left to right.

Oxidation takes place at the anode, reduction takes place at the cathode. If you go to a farm, you might see An Ox or a Red Cat to help you remember this. Or you might remember that anode and oxidation both start with vowels, cathode and reduction both start with consonants.

Cell Notation:

This is one of the places where you get to use the "pipe" character on your keyboard!! It's usually a shifted backslash and appears as a tall vertical line, "|".
Standard cells are usually written as reductions... if you're pairing two half-cells, one of them will have to be reversed to give an oxidation. How do you decide which one to reverse? Well...

Cell Potentials and Half-Cell Potentials:

A cell potential is a measure of the "load" in the above picture. If the electrons are moving as shown in the picture, the cell potential is positive. This is a spontaneous voltaic cell. Rather than measuring every possible combination of half-cells, we measure all of them relative to an accepted reference, usually the Standard Hydrogen Electrode ("SHE").
Oxidation and reduction are opposite processes. If a reduction potential (as measured relative to SHE) is know, changing the sign yields the oxidation potential for the opposite process.
This is another place where your textbook uses a very common pocess that I don't always like. To figure out cell potential, your book uses the equation:
E°cell = E°cathode - E°anode 
That equation works. The reason I don't always like it is that you have to remember which half is the cathode and which is the anode, and you have to put them in the right order because subtraction is directional.
When I'm looking at cell potentials, I use the equation:
E°cell = E°reduction + E°oxidation 
For the oxidation half reaction, change the sign of the standard reduction potential to get a standard oxidation potential, then add them up. Mathematically, "change the sign" is the same as "subtract", so these really are the same process, but explicitly changing the potential of the oxidation reaction to an oxidation potential makes my method more connected to understanding the processes that are taking place and less about memorizing an equation. That's why I prefer it. Either one will give you the same correct answer.

Non-standard Conditions and Equilibrium:

These are closely related to the discussions in the previous chapter and are described well in your textbook. Using these relationships, you should be able to interconvert between K°, ∆G°rxn, and E°cell.

Batteries and Fuel Cells:

These are some of the most obvious applications of redox chemistry. Read them over and make sure you can identify half-reactions and concepts.

Corrosion:

One of the most important reasons to study something like redox chemistry is so that we can use it to our advantage. There are many important investments that are made of metal, from cars to grain bins. Corrosion destroys these investments, so it's worth considerable effort to protect them from corrosion.
The two main ways to protect something from corrosion are:
  1. Anodic Inhibition - When a metal corrodes ("rusts"), it is oxidized from elemental metal with an oxidation number of zero to a metal cation with a postive oxidiation number. Oxidation takes place at the anode, so what we are trying to accomplish by anodic inhibition to to prevent the metal from being an anode. This is most often done with some sort of coating (a paint or oil or grease) that forms a barrier and prevents transfer of eletrons and ions. In the picture above, if the anode is painted, the electron and cation transfer that is required for the cell to work will not happen. Anodic inhibition has the same effect as snipping the wire that connects the two electrodes via the external curcuit... if we break the circuit, the reaction cannot happen.
  2. Cathodic Protection - This is a little more "clever" approach. Rather than just preventing the metal from being the anode, in cathodic protection we force the metal to be a cathode by connecting it to a more active metal. The more active metal serves as the anode and is sacrificed to protect the metal of interest. This is often used with things like grain bins; a large rod of zinc (or another more active metal) is connected to the steel of the grain bin and then pounded into the ground. The zinc rod corrodes while the steel bin does not. These sacrificial anodes must be replaced from time to time as they corrode. Another example of this is galvanized steel. This is steel that has been coated with a thin layer of zinc metal to act as both a coating (anodic inhibition) and a sacrificial anode (cathodic protection) for the steel.

Electrolysis:

As mentioned above, a non-spontaneous redox reaction can be forced by applying an electrical current. In the picture above, the "load" is passively observing the flow of electrons in the extrenal circuit, but if a source rather than a load was connected to the external circuit, we could force the reaction backwards, in the non-spontaneous direction.
Electrolysis is an important process for production of a variety of pure chemicals, but it can be expensive.
From a calculation perspective, let the units guide your problem setup... Faraday's constant and electrical units such as amperes have units that can help guide problem setup.


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