7.2.2: Brønsted–Lowry Acids and Bases

Brønsted–Lowry¹²⁸ Acids and Bases

The Arrhenius acid–base model is fairly easy to understand but its application is limited to certain kinds of reactions. Rather than continue down this road, chemists found that they needed to expand their model of acids and bases and how they react. The first of these expansions was the Brønsted–Lowry model. In the Brønsted–Lowry model, an acid is characterized as a proton (H⁺) donor and a base as a proton acceptor. If we revisit the reactions we looked at earlier in the context of the Brønsted–Lowry acid-base model, we see that HCl is the proton donor; it gives away H⁺ and water is the proton acceptor. In this scheme, HCl is the acid and water is the base:

HCl(g) + H2O(l) ⇄ H3O+ (aq) + Cl–(aq)

acid base acid base

conjugate conjugate

The resulting species are called the conjugate acid (so H₃O⁺ is the conjugate acid of H₂O) and the conjugate base (Cl^– is the conjugate base of HCl). This is because H₃O⁺ can and generally does donate its H⁺ to another molecule (most often another water molecule) and Cl^– can accept an H⁺.

A major (and important difference) between the Brønsted–Lowry and Arrhenius acid–base models is that a Brønsted–Lowry acid must always have an accompanying base to react with— the two are inseparable. A proton donor must have something to donate the protons to (a base)— in this case, water. Remember that bond breaking requires energy, whereas bond formation releases energy. Some energy input is always required for a reaction in which the only thing that happens is the breaking of a bond (for example the Cl–H bond in HCl). Acid–base reactions are typically exothermic; they release energy to the surroundings and the released energy is associated with the interaction between the H+ and the base. In other words, the proton does not drop off the acid and then bond with the base. Instead, the acid–H bond starts to break as the base–H bond starts to form. One way that we can visualize this process is to draw out the Lewis structures of the molecules involved and see how the proton is transferred.

As shown in the figure, we use a dotted line to show the growing attraction between the partial positive charge on the H of the H—Cl molecule and the partial negative charge on the oxygen. This interaction results in the destabilization of the H—Cl bond. Because the Cl is more electronegative than the H, the electrons of the original H—Cl bond remain with the Cl (which becomes Cl^-) and the H⁺ forms a new bond with a water molecule. Essentially, a Brønsted–Lowry acid–base reaction involves the transfer of a proton from an acid to a base, leaving behind the original bonding electrons.

Another example of an acid–base reaction is the reaction of ammonia with water:

NH_3(aq) + H₂O_(l) ⇄ NH₄⁺_(aq) + ^–OH_(aq)

base acid conjugate acid conjugate base

In this case, oxygen is more electronegative than nitrogen. The proton is transferred from the oxygen to the nitrogen. Again, the dotted line in the figure represents the developing bond between the hydrogen and the nitrogen. As the H—O bond breaks, a new H—N bond forms, making the resulting NH₄⁺ molecule positively-charged. The electrons associated with the original H—O bond are retained by the O, making it negatively-charged. So, water is the acid and ammonia is the base! An important difference between this and the preceding HCl–H₂O reaction is that H2O is a much weaker acid than is HCl. In aqueous solution, not all of the NH₃ reacts with H₂O to form NH₄⁺. Moreover, the reaction between NH₃ and water is reversible, as indicated by the ⇄ symbol. The next chapter will consider the extent to which a reaction proceeds to completion. You may be wondering why the water does not act as a base in the reaction with NH₃, like it does with HCl. If you draw out the products resulting from a proton transfer from nitrogen to oxygen, you will see that this process results in a mixture of products where the more electronegative atom (O) now has a positive charge, and the less electronegative atom (N) has a negative charge. It does not make sense that the most electronegative atom would end up with a positive charge, and indeed this process does not happen (to any measurable extent).

We will soon return to a discussion of what makes a compound acidic and/or basic. At the moment, we have two acid–base reactions: one in which water is the acid and the other in which water is the base. How can this be? How can one molecule of water be both an acid and a base, apparently at the same time? It is possible because of the water molecule’s unique structure. In fact, water reacts with itself, with one molecule acting as an acid and one as a base:

H₂O_(l) + H₂O_(l) ⇄ H₃O⁺ _(aq) + ^–OH_(aq)

acid base conjugate acid conjugate base

As shown in the figure, we can again visualize this process by drawing out the Lewis structures of the water molecules to see how the proton is able to move from one water molecule to another, so that it is never “alone” and always interacting with the lone pairs on the oxygens.

Questions to Ponder

• Between the Arrhenius model and the Brønsted–Lowry model of acids and base, which is more useful? Why?

Questions to Answer

• Which do you think is more likely to happen? The reaction H₂O + H₂O → H₃O⁺ + ^–OH? Or the reverse process H₃O⁺ + ^–OH → H₂O + H₂O? Could they both happen at once?

• What do you think the relative amounts of H2O, H3O+ + –OH might be in a pure sample of liquid water? How would you measure the relative amounts?

• Now that you know HCl is an acid and ammonia is a base, can you predict the reaction that occurs between them?

• Is water a necessary component of a Brønsted–Lowry acid–base reaction? How about for an Arrhenius acid–base reaction?

References

¹²⁸ This theory was postulated simultaneously by both Brønsted and Lowry in 1923.