Strengths of Acids and Bases

When you pour concentrated hydrochloric acid on yourself, it hurts. A lot. As a student who frequents the Gen Chem lab, where my peers very rarely think to wipe off the side of an acid container before handing it to the next person in line, I’ve learned this firsthand. When we work with 16 M HCl, people get hurt. When we work with 0.1 M, however, very few people end up in the emergency room.

I believe this is what people think of when they think of the difference between a strong and a weak acid. However, there are some acids that are inherently stronger than others. A 1 molar solution of hydrochloric acid has a pH of 1 (that’s as acidic as it gets!), whereas a solution of the same concentration of acetic acid has a pH of 2.4, which is notably weaker. (For those of you who like cooking, 1 M acetic acid is just vinegar.) Obviously, there would be a difference in making salad dressing with hydrochloric acid instead of vinegar. So, then, what accounts for the innate differences in acid strengths?

Acids and bases both have varying strengths that are independent of concentration. The differences in these strengths have to do with how much of the acid or base ionizes.

When an acid or base ionizes, it splits up into its constituent parts. HCl ionizes into H+ and Cl. Acetic acid, CH3COOH, dissociates into H+ and the acetate anion, CH3COO. Sodium hydroxide, NaOH, dissociates into Na+ and OH-, et cetera.

All acids and bases do this to some extent. However, not all acids or bases ionize completely. Herein lies the key to acid and base strength.

A strong acid or base is an acid or base that ionizes completely, meaning that there is no acid or base left when it is dissolved in water. My favorite strong acid, hydriodic acid, or HI, dissociates like this:

HI (aq) → H+ (aq) + I (aq)

Notice the single-sided arrow. This indicates that this reaction goes to completion, or that essentially all hydriodic acid dissociates into protons and iodide anions in water. You can work through this logically and figure out why the pH is low (or acidity is high) in a solution like this. Since all HI dissociates into protons and iodide, the proton (well, hydronium, but don’t worry about that right now) concentration is higher than it would be in a solution where complete dissociation didn’t occur. For example, let’s look at a slightly anomalous acid and my favorite weak acid, hydrofluoric acid.

Hydrofluoric acid, or HF, is anomalous because, even though it’s a hydrogen halide (meaning it’s got a hydrogen attached to a halogen, or one of the atoms in the column second from the right on the periodic table), it isn’t a strong acid. All of the other acids from this column (HCl, HBr, HI) are strong acids, but fluorine, because it’s tiny and stingy with its attached protons, doesn’t completely ionize. Instead, its dissociation looks like this:

HF (aq) ⇌ H+ (aq) + F(aq)

Notice that, in contrast to the dissociation of hydriotic acid, there is a double-sided arrow, or a symbol that indicates chemical equilibrium, between the acid and its dissociated components. Chemical equilibrium is discussed at length here, but the gist of it is that, at equilibrium, as fast as hydrofluoric acid molecules are breaking apart, protons and fluoride ions are combining. This means that there is an unchanging concentration of each when the reaction “ends,” instead of all of one being completely used up. In contrast to the dissociation of strong acids or bases, where all of the acid or base completely “falls apart,” in weak acids or bases, there is still some acid or base left at the end of the reaction. The proportions of acid or base to constituent parts changes for each reaction, and are described by Ka and Kb, respectively. The pH of a solution of a weak acid has a lower concentration of protons (or hydronium) than a solution of identical acid concentration of a strong acid. Thus, a solution of a weak acid has a higher pH and, therefor, is less acidic (pH < 7 is acidic, pH > 7 is alkaline or basic) than a similar solution of a strong acid.

This talk of acid and base strength is all well and good, but what if we want to know how dissolving something less obvious, like a salt, would affect pH? There is a specific and interesting part of acid-base chemistry devoted just to this, and it works with conjugate acids and bases.

Want to know more about acid/base chemistry? Wish I’d move on? Say so in the comments! Want to correct me? Feel free to do that too! Just, please, don’t cite Wikipedia when you do, because I’ve spent enough time there to know if you’re lying, you rascal, you.


2 thoughts on “Strengths of Acids and Bases

Leave a Reply

Fill in your details below or click an icon to log in: Logo

You are commenting using your account. Log Out /  Change )

Google+ photo

You are commenting using your Google+ account. Log Out /  Change )

Twitter picture

You are commenting using your Twitter account. Log Out /  Change )

Facebook photo

You are commenting using your Facebook account. Log Out /  Change )


Connecting to %s