Factors affecting acid strength | Acids and bases | AP Chemistry | Khan Academy

Factors that affect acid strength include bond polarity, bond strength, and conjugate base stability.

Let's think about a generic acid, HA, that donates a proton to water to form the hydronium ion, H3O⁺, and the conjugate base, A⁻. First, let's consider the polarity of the bond between H and A. If A is more electronegative than H, A withdraws electron density, so we could draw an arrow pointing towards A, right, as the electrons in the bond between them are pulled closer to A. As the electronegativity of A increases, there's an increase in the polarity of the bond. As the bond polarity increases, more electron density is drawn away from the H, which makes it easier for HA to donate a proton. Therefore, in general, an increase in the polarity of the H-A bond means an increase in the strength of the acid.

Next, let's think about the factor of bond strength. Let's consider the strength of the bond between H and A. The weaker the bond, the more easily the proton is donated. Therefore, in general, a decrease in the bond strength means an increase in the strength of the acid. The stability of the conjugate base can also affect the strength of the acid. The more stable the conjugate base, the more likely the acid is to donate a proton.

If we think about that for HA, the conjugate base is A⁻, and the more stable A⁻ is, the more likely HA will donate a proton in solution. Therefore, in general, the more stable the conjugate base, the stronger the acid. So, let's go ahead and write here: an increase in the stability of the conjugate base means an increase in the strength of the acid. Even though acid strength is usually due to all three of these factors—bond polarity, bond strength, and conjugate base stability—when we look at examples, we're only going to consider one or two factors that are the main contributors to acid strength.

Let's look at the binary acids from group 7A on the periodic table, so that's hydrofluoric acid, hydrochloric acid, hydrobromic acid, and hydroiodic acid. As we go down the group from fluorine to chlorine to bromine to iodine, there's an increase in the strength of the acid. So, out of these four, hydroiodic acid is the strongest. The main factor determining the strength of the binary acids in group 7A is bond strength.

Looking up values for bond enthalpies allows us to figure out the strengths of these bonds. For example, the HF bond has a bond enthalpy of X kilojoules per mole, while the HI bond has a bond enthalpy of 299 kilojoules per mole. The lower the value for the bond enthalpy, the easier it is to break the bond, and because bond enthalpies decrease as we go down the group, that means there's a decrease in bond strength. A decrease in bond strength means it's easier for the acid to donate a proton; therefore, we see an increase in the strength of the acid as we go down the group.

Next, let's compare the strengths of some oxyacids. These oxyacids all have the general formula XOH, where X is a halogen. The acidic proton is the proton that's directly bonded to the oxygen. If an oxyacid donates its proton to water, that forms the hydronium ion, H3O⁺, and the conjugate base to the oxyacid. For these three oxyacids, the halogens are iodine, bromine, and chlorine. As we go from iodine to bromine to chlorine in group 7A on the periodic table, that's an increase in the electronegativity. So, chlorine is the most electronegative out of these three halogens.

As we go up the group in our halogens, there's an increase in the strength of the acid. So, hypochlorous acid is the strongest of the three. The main factor determining the strength of these oxyacids is the bond polarity, which is affected by the electronegativity of the halogen. The polarity of this oxygen-hydrogen bond is affected by the presence of the halogen. As the electronegativity of the halogen increases, the halogen is able to withdraw more electron density away from the right side of the molecule. That increases the polarity of the O-H bond and makes it easier to donate this proton. Therefore, as the electronegativity of the halogen increases, the acidity of the oxyacid increases. This effect of the electronegative atom increasing the acidity is often referred to as the inductive effect.

Let's compare hypochlorous acid to two other oxyacids and notice how I've left the formal charges off of these acids just so we can focus on general structure. Notice how the acidic proton is directly bonded to the oxygen in all three of these oxyacids, and in all three of these oxyacids, the oxygen is directly bonded to a chlorine. Notice what happens to the structure as we move to the right, comparing chlorous acid to hypochlorous acid. Chlorous acid has an additional oxygen bonded to the chlorine, and looking at perchloric acid, instead of only one additional oxygen, there are three additional oxygens.

So, as we move to the right, we're increasing in the number of oxygens bonded to the chlorine. Oxygen is a very electronegative element, so as we move to the right, we're increasing in the number of electronegative atoms in the acid. As the number of electronegative atoms increases, more electron density is pulled away from the acidic proton, which increases the polarity of the oxygen-hydrogen bond. So, bond polarity increases as we move to the right, which predicts an increase in the strength of the acid as we move to the right, and that's what we observe experimentally.

Perchloric acid is the strongest of the three. In reality, all three factors affect the strength of the acid; however, for simplicity's sake, we could just say that increasing bond polarity is the main factor for the increasing acid strength in these oxyacids. Carboxylic acids are a group of acids that all contain a carboxyl group. A carboxyl group consists of carbon, oxygen, oxygen, and hydrogen.

So, if we look at acetic acid, I'll circle the carboxyl group on acetic acid, and I can also circle the carboxyl group on formic acid. The acidic proton in a carboxylic acid is the one that's directly bonded to the oxygen in the carboxyl group. One reason why this proton is acidic is because of the presence of this oxygen in the carboxyl group. This oxygen-hydrogen bond is already polarized, but the presence of another electronegative atom further increases the polarity of the oxygen-hydrogen bond. Increasing the bond polarity makes it more likely to donate the proton, which increases the acidity.

For carboxylic acids, it's also important to consider the stability of the conjugate base. When acetic acid donates its proton, it turns into its conjugate base, which is the acetate anion. Notice that the oxygen that used to be bonded to the proton now has a negative formal charge on it. There are two possible resonance structures that you can draw for the acetate anion. The first is with the negative charge on this oxygen, and then we could draw another resonance structure with a negative charge on the other oxygen.

In reality, neither resonance structure is a perfect representation of the acetate anion, and we need to think about a hybrid of these two resonance structures. In the hybrid, the negative charge isn't on one of the oxygens; that one negative charge is spread out or delocalized over the two oxygens. So, it's like one oxygen has a charge of negative one-half and the other oxygen has a charge of negative one-half. This delocalization of the negative charge stabilizes the conjugate base, and the more stable the conjugate base, the stronger the acid. Therefore, the stability of the conjugate base also affects the acidity of the carboxylic acid. So because carboxylic acids have conjugate bases that are resonance stabilized, carboxylic acids like acetic acid and formic acid are more acidic.