In this next series of posts we are going to discuss the reactions of alcohols.
As a functional group, alcohols are introduced fairly early in organic chemistry. Their reactions, however, are usually not covered until near the end of Org 1 – at least after subjects like substitution and elimination reactions have been explored. That’s because, as we will see, the reactions of alcohols often fall into these categories.
We’ll cover the reactions of alcohols in due course. But today, let’s just get our feet wet by familiarizing ourselves with their structure, nomenclature and physical properties.
Structure and Nomenclature
Alcohols are organic molecules containing the “hydroxyl” functional group, “OH” directly bonded to carbon. The carbon directly attached to OH is technically called the “carbinol” carbon, although this nomenclature is often not introduced in introductory classes. The carbinol carbon (carbon attached to OH), however, is the key to understanding the most common classifications we use for alcohols, that being “primary”, “secondary”, and “tertiary” alcohols.
To determine if an alcohol is primary, secondary, or tertiary, examine the carbon attached to OH. If that carbon is attached to one carbon, the alcohol is primary; two, secondary; three, tertiary. If zero carbons and three hydrogens (a unique situation) it is methanol. Hydroxyl groups attached to aromatic rings are called, “phenols”.
The deeper nomenclature of alcohols is not something I’m going to discuss on this blog at the present. There are all kinds of fantastic videos and resources on the internet that describe the naming of alcohols in detail, and I will happily refer you to those for now.
Here’s an example by Leah Fisch, for example.
The key to understanding the physical properties of alcohols – and as we shall see later, their reactivity – is to appreciate just how polarized the hydroxyl group is. Oxygen has an electronegativity of 3.5 and hydrogen 2.2. Oxygen, being “greedier” for electrons than hydrogen, will have more than its “share” of the two electrons in the O-H bond. This means that oxygen will be more “electron rich” (more negative) and hydrogen more “electron poor” (more positive) than they would be in a bond where electrons were shared perfectly equally. We say that electron density in the O-H bond is strongly “polarized” toward oxygen. Another way of saying the same thing is that the O-H bond has a strong “dipole”.
Since opposite charges attract, these partial charges will line up in solution in a way where the partial negative oxygens on one molecule interact with the partially positive hydrogens on another. These interactions between molecules, which we call, “hydrogen bonds” are about 1/10 as strong as normal bonds, and are much more transient, lasting only a fraction of a second on average. Still, the effect of hydrogen bonding is to lead molecules to “stick” to one another.
A non-scientific analogy would be to compare the dipoles that lead to hydrogen bonding to the way Velcro “hooks” stick to Velcro “fuzz”. They’re not strong bonds, necessarily, but are sticky enough:
This has two important implications.
First, hydroxyl groups greatly increase boiling points. Look what happens to the boiling point of propane (an alkane) when a CH3 group is replaced with an OH group – the molecular weights are the same, but there’s over a 100° C difference in boiling point. [Why is it important to compare molecules of roughly similar molecular weight? Because boiling points also with increased Van Der Waals interactions, which are roughly proportional to increased surface area. Keeping the molecular weights constant ensures we’re comparing apples to apples in this respect].
If one is good, then surely more is better? Yes! Replacing a CH3 on propanol with an OH group gives us “ethylene glycol”, a “di-ol” (actually called a “vicinal diol” since the alcohols are on adjacent carbons). Having twice as many hydroxyl groups, we would expect it to be even more polar than propanol, and thus have a higher boiling point. This is true – the boiling point of ethylene glycol is 197°C.
The differences in boiling point between primary, secondary, and tertiary alcohols can be subtle. Generally, the more “exposed” the hydroxyl group, the more other OH groups it will be able to interact with, and the higher the boiling point. The isomers of butanol are a perfect example. See how the primary alcohols (1-butanol and 2-methyl-1-propanol) have higher boiling points than the secondary alcohol (2-butanol) which has a higher boiling point than the tertiary alcohol (t-butanol).
The second key implication of hydrogen bonding is that the hydroxyl group imparts much greater water solubility upon organic molecules. That’s because water itself is a hydrogen-bonding solvent, and therefore the dipole of the hydroxyl group can interact favourably with the hydroxyl group of H2O. Ethanol and water, for example, are mixable (“miscible” is the word we actually use) in all proportions. [It isn’t until butanol that the greasy alkyl chain starts interfering with water solubility.
We often say that hydroxyl groups are a “polar” functional group. Solvents containing OH are generally considered to be “polar” solvents because of the large dipole. You may have also heard of alcohols being referred to as “polar protic” solvents. What’s that about? It has to do with one of the key properties of alcohols – their ability to participate in acid-base reactions. That’s the subject of the next post!