The 1970’s was the decade that illuminated the threat of acid rain to the citizens of the US. It had been known to exist several years before, but the sources of the problem did their best to suppress the information. It wasn’t until the environmental damage became significant enough to draw national attention that it would lead to the US enacting regulations to stop acid rain.
Truthfully, most of the public was probably still unaware of what acid rain actually was. The default mental image that comes to the mind of the non-chemist is large drops of battery acid raining down from the heavens and devouring everything. This is not quite the case, however. Pure water has a neutral pH of 7. Normal rain is actually slightly acidic as it picks up CO2 from the air, making carbonic acid. But when this “normal” rain mixes with the byproducts of industrial plants that pump out large amounts of SO2 (sulfuric dioxide) and NO (nitrogen oxide) into the atmosphere, it becomes even more acidic – down to a pH of 3. This “acid” rain has the acidity of citrus juice, so it’s not going to set the world on fire. But it will wreak havoc on local ecosystems.
The 1990’s brought with it tough government regulations on the output of SO2 and NO by large factories, pretty much eliminating acid rain in the US. The rise and fall of acid rain is a great example of why we should educate ourselves on the basic chemistries that define our lives, even though we might not be actual chemists. In this article, we’re going back to your first year of college and hash out just what defines an acid and base. And solidify our understanding of the pH scale. It is essential for the future biohacker to have this knowledge in their toolbox.
The Power of Hydrogen
Our journey begins with plain old water. You probably believe that pure water is nothing but H2O. The problem with this picture is that water has a tendency to auto-ionize. What happens is that a hydrogen atom will break off of one water molecule and join another. This results in two molecules – H3O (hydronium) and OH (hydroxide). As you might have already guessed, the +H3O has a positive charge and the -OH has a negative charge. The process forms an equilibrium in water at 25 C, so that solution remains neutral. This means that pure water at 25 C is composed of H2O and equal parts of +H3O and -OH.
Adding different molecules to the water can disrupt the balance of the hydronium and hydroxide ions.
- When [+H3O] = [-HO], the solution is neutral.
- When [+H3O] > [-HO], the solution is an acid.
- When [+H3O] < [-HO], the solution is a base.
The number of +H30 ions in the solution will determine how acidic it is. Generally, chemists use the mole to count the number of atoms or molecules in a sample. The mole is kind of a “chemist’s dozen”. A dozen is a count of twelve, a couple is the count of two, and the mole is the count of 6.022 x 1023. However, the number of +H3O ions present in an aqueous solution can vary by several orders of magnitude, so another type of scale was created to better quantify this amount. It’s called the pH scale, with pH standing for “power of Hydrogen”.
To avoid getting our brains tied up in not-very-helpful conversions, take it as a given that the concentration of +H30 ions in neutral water is about 1.0 x 107 molar. So we say this neutral water has a pH of 7. As the number of +H3O ions increases, so does the pH. Orange juice has a pH of about 3.5, which means its +H3O concentration is 1.0 x 10-3.5 M. Ammonia has a concentration of 1.0 x 10-11 M, which gives it a pH of 11. You get the idea. Simple, isn’t it!
The pH scale is logarithmic, so acid rain with a pH of 3 has 10 x 10 x 10 or 1,000 more times the amount of +H30 ion concentration that typical rain water, which has a pH of 6.
In our previous article, we defined an acid as a substance that produces hydrogen ions when dissolved in water, which is known as the Arrhenius model. While this definition is sufficient for most of what you’ll run across as a biohacker, we should take the time to mention a more inclusive definition developed by Danish chemist Johannas Bronsted in 1932. His definition can explain acids and bases without the need for them to be dissolved in an aqueous solution. But it’s a little more in-depth.
Consider the above example of how water will autoionize. We explain this by a hydrogen atom breaking free of a water molecule and then joining another to make H3O, and leaving an OH in its wake. While this is helpful for explanation, it’s not entirely accurate. When the H atom leaves its water molecule, it leaves its electron behind. And since a hydrogen atom consists of one proton and one electron (there is no neutron in H), that means it’s basically a free proton. Now, this proton has a positive charge, and the oxygen in water is slightly negative. So the proton will be attracted to the oxygen atom on other H2O molecules. H3O could be written more accurately as [H2O * H+]. With the understanding that free hydrogen ions in an acidic solution are (hydrogen) protons, a better definition can be developed. The Bronsted definition defines an acid as a solution with proton donors, and a base as a solution with proton acceptors.
Consider the example of hydrofluoric acid. When you mix hydrogen fluoride (HF) with water, the hydrogen atoms and fluorine atoms break apart is a process known as dissociation. The hydrogen atom leaves its electron with the fluorine, so you wind up with something like:
HF(aq) + H2O(l) <–> +H3O(aq) + -F(aq) (aq) means aqueous and (l) means liquid state of matter
By the Arrhenous definition, the HF makes free hydrogen ions and is therefore an acid. With the Bronsted definition, the HF donates a proton (the hydrogen atom without its electron), and is therefore a proton donor. Which also makes it an acid.
You should now have a basic idea of what an acid and base are, and why they are defined the way they are. This essential knowledge is needed before we take on the much more exciting amino acids – the building blocks of proteins, and all life on earth.