So far we've seen a few definitions of an acid, with the most popular being the Brønsted-Lowry one that tells us an acid transfers proton. We also looked at examples of strong acids and weak acids and how their equilibriums in water play out. Up until this section, we could figure out the relative strength of an acid by knowing its acid-dissociation constant: the bigger the acid-dissociation constant, the stronger the acid.
In the spirit of keepin' it real and honest, it's probably best to just memorize a bunch of the acids (and bases) we went over already and whether they are strong or weak. If you forget a few, we'll talk about the properties that determine acid strength in this section—it just might help you out if you find yourself in a pickle.
The Periodic Table is a basically a chemistry treasure map. There might not be pot of gold hidden in it, but it might help you ace your next test. We'll be referring to it in this guide to help predict acid strength.
Two big predictors of acid strength are the H-A bond strength and the H-A bond polarity. Both properties can be mapped on to the Periodic Table and follow certain trends.
Let's take a look at the strength of the halogen (Group 7) containing acids HF, HCl, HBr, and HI.
The bond strength of an acid generally depends on the size of the 'A' atom: the smaller the 'A' atom, the stronger the H-A bond. When going down a row in the Periodic Table (see figure below), the atoms get larger so the strength of the bonds get weaker, which means the acids get stronger. For the halogen-containing acids above, HF has the strongest bond and is the weakest acid. The strong bond between the more similarly-sized 'H' and 'F' atoms doesn't want to break and allow the 'H' to transfer.
HI, on the other hand, is a very strong acid. The big 'I' atom overpowers the helpless little 'H' and the H-I bond is very weak. Therefore, HI is a great proton transferer and one heck of an acid (Ka ~ 109 M).
When comparing acids that have 'A' atoms in the same row, bond polarity differences are more important in determining acid strength. That's because bond strength differences are much smaller between atoms nearby each other in the same row.
Bond polarity is largely determined by the electronegativity difference between the two atoms involved in the bond. Electronegativity is basically how much an atom wants electrons. Think of electronegativity as a measure of an atom's electron crush. Yes, we're talking about the boy band kind of crush. For bonds that involve an atom that has a huge electron crush and an atom that just isn't feeling the electron love, the bond tends to be really polar. The electrons reciprocate the crush and go for the atom that likes them the most. Isn't that nice?
The hydrofluoric acid (HF) bond is polar because F really loves electrons. Let's compare this to CH4. Carbon is in the same row as fluorine (check out the figure above), but HF is a much stronger acid then CH4. The C-H bonds are not polar compared to the H-F bond. CH4 is a really weak acid.
Here's another class of acids that also have predictable strengths based on the Periodic Table: the oxoacids. They have the general formula, HnYOm.
Some real-world examples are H2CO3, H2PO4, and HNO3. These acids contain an O-H bond that dissociates to form a hydronium ion and a conjugate base:
(Note: In this section, when we say "Y atom," we don't mean yttrium. We're using the Y as a placeholder for an element.)
The stronger an acid is the more the right side of the equilibrium is favored. The more the Y atom is able to stabilize the Y-O- negatively charged product on the right side of the equilibrium, the stronger the acid will be.
If Y has a big electron crush (is highly electronegative) it will be happy to be on the right side of the equilibrium. That's because there are more electrons on the molecule on the right side of the equilibrium. Even though Y doesn't get the electrons all to itself, it still gets some satisfaction from knowing the neighboring oxygen atom is enjoying them. Isn't that nice?
Overall, the more electronegative the Y atom is, the better it can stabilize the Y-O- product and the stronger the acid will be. The hypohalous acid series does a good job showing increasing acid strength with increasing electronegativity of the Y atom (in this case, a halogen atom).
Anytime the O-H bond is weakened the stronger the acid will be. In the example above, the O-H bond is weakened by increasing the electronegativity of the Y atom. Think of the Y atom as an electron vacuum cleaner that sucks the electrons out of the O-H bond so that they no longer get shared with the proton and end up on the conjugate base molecule. The stronger the vacuum pulling the electrons out the bond, the stronger the acid will be.
The same principle holds for acids that contain the same Y atom but different numbers of oxygen atoms. Oxygen atoms are also like electron vacuum cleaners. They weaken the O-H bond through the central Y atom and stabilize the negatively charged product. As a result, the more oxygen atoms are attached to the central Y atom, the stronger the HnYOm acid.
The series of oxoacids of chlorine illustrates what we mean:
Increasing the number of oxygen atoms that are attached to the central atom also increase the oxidation number of the central atom. High oxidation numbers of the central atom represent positive charge on that atom.
Since opposite charges attract, a very positive central atom would be more attracted to the negative charge from the electron lone pair on the neighboring oxygen atom. In order for the neighboring oxygen to get the coveted lone pair electrons, it transfers a proton and acts as an acid.
While the above trends and examples will be useful, you will inevitably have close encounters with alien molecules. Reminds us of an old Spielberg movie. When we need to compare the acidity of molecules we haven't seen before, try to determine how many electron vacuum cleaners (if any) there are on the molecule. These "vacuum cleaner" atoms are typically those with high electronegativities, like oxygen. These atoms stabilize the lone pair of electrons that usually result when an acid transfers a proton. The better able the molecule can stabilize the extra electrons, the stronger it will be as an acid.