Lec XIX

 Chem 1200

Angel C. de Dios

Reactions XI

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For this section, we will focus on chemical reactions that involve a transfer of electrons, reduction-oxidation (redox) reactions. Some of you may have seen this in high school chemistry and have worked on balancing these reactions. Our focus, however, is on how these reactions can be facilitated by an electric wire. The reactants are not in the same compartment but with a piece of wire, electrons can be transferred from one compartment to the other. Reduction occurs in one compartment, and oxidation happens in the other compartment. 

As you know, most metals are electropositive, meaning they tend to lose electrons. However, as in any process, some metals are more eager than others when it comes to losing electrons. One can therefore envision a metal reacting with a cation of another metal, as illustrated in the following example:

Zn(s)  + Cu²⁺(aq)  à  Cu(s)  +  Zn²⁺(aq)

Such a reaction happens. If one dips a strip of Zn metal into an aqueous solution containing Cu(II) ions, copper gets plated on the Zn strip. In this focus, we are going to look deeper into this type of reactions, but the electrons are not transferred directly from a zinc atom to copper(II) ion. The electrons will be travelling through a wire. This is electrochemistry.

But, first, let us review some basic concepts in reduction-oxidation chemistry.

First, whenever there is reduction, there is oxidation.

Thus, if electrons are being transferred, we need to keep track of electrons. We accomplish this by assigning oxidation numbers to each element.

Of course, there are rules we must follow in assigning oxidation numbers. For most of the reactions we will study this semester, assigning oxidation numbers will be straightforward since we will normally encounter metals, which are elements, and are therefore assigned an oxidation number of zero. In addition, we will likewise see cations. And in these cases, the oxidation number of a cation is equal to its charge. Nonetheless, it is useful to review how oxidation numbers are assigned in polyatomic species that involve more than one type of element.

These rules are given in hierarchical order. Rules stated first are more important to follow than later rules.

Here are instances where we follow the order of the rules.

Since the rule on F comes first before the rule on O, in the compound OF2, F is assigned an oxidation number of -1, and O must be assigned an oxidation number of +2. 

Since the rule on Group I elements comes first before the rule on H, in the compound NaH, Na is assigned an oxidation number of +1, and H must be assigned an oxidation number of -1.

Since the rule on H comes first before the rule on O, in the compound H2O2, H must be assigned an oxidation number of +1 first, so O must be assigned an oxidation number of -1.

Again, as we shall see in this section, we will be working mostly with metals and their cations. We may also see halogens and halides, and perhaps, protons as well as hydrogen gas. For these cases, assigning oxidation numbers should be easy. The oxidation number of a metal is zero. The oxidation number of a metal cation is equal to its charge. For the elements that are gaseous, their oxidation number is zero. Examples are fluorine (F2), chlorine (Cl2), hydrogen (H2), and oxygen (O2).  For the halides, their oxidation number is equal to their charge, -1. And in water, H is assigned an oxidation number of +1 and O is assigned an oxidation number of -2.

Not all elements have rules because combining the special rules for Group I and II, F, H, O, and halides with the first two rules, we can assign oxidation number for the other elements not mentioned in any rule. The following are examples.

An oxidizing agent gains electrons, it gets reduced. A reducing agent gives electrons, it gets oxidized.

The following is a bit more complicated example.

Most of the discussion here apply only to inorganic or general chemistry. In these areas, compounds often have equivalent atoms for each element. In organic chemistry, compounds may contain several carbon atoms, but not all of these carbon atoms are equivalent. Assigning one oxidation number to all the carbon atoms is therefore invalid. One example that you see in the previous topics is acetic acid:  CH3COOH. The two C's in this compound are clearly not identical and should not be assigned the same oxidation number.

So, now let us look at the thermodynamics of redox reactions.

In electrochemistry, if the redox reaction is spontaneous, we call the reaction cell, a galvanic or voltaic cell, or in common terms, a battery.

The following is a diagram describing a galvanic cell.

We reviewed the parts of an electrochemical (galvanic, battery) cell:

Finally, electrons travel through the wire and ions migrate inside the salt bridge. Anions (negatively charged ions) migrate toward the anode and cations (positively charged ions) migrate toward the cathode.

The electrode at the anode is the first one to be listed in the notation.

The electrode at the cathode is the last one to be listed in the notation. Both electrodes are connected through a wire. The electrodes need not be participating in the reaction, (In this example, Zn is a reactant) an electrode can be an inert metal whose sole purpose is to be connected to a wire.