Bronsted-Lowry Acids and Bases - Online Tutor, Practice Problems & Exam Prep

The Arrhenius model defines acids as substances that increase H⁺ ions in solution and bases as those that increase OH^- ions. The Brønsted–Lowry theory expands this by defining acids as proton donors and bases as proton acceptors, applicable beyond aqueous solutions. For example, HCl donates H⁺ to water, forming H₃O⁺. This model introduces conjugate acid-base pairs, essential for understanding acid-base reactions and titrations, highlighting the importance of proton transfer in various chemical contexts.

Bronsted-Lowry Acid-Base

So we started with the simplest definitions of acids and bases on the Arrhenius model where the acid was just increasing the amount of H⁺ ion in solution and the base was increasing the amount of OH^- concentration in solution. After that, we move on to the Bronsted-Lowry definitions of acids and bases. Now in 1923, Johannes Brønsted and Thomas Lowry developed their new set of definitions for what constitutes an acid and a base. Now according to them, acids are considered proton, which is the same thing as your hydrogen ion or hydronium ion donors. And bases are considered to be proton acceptors.

So unlike Arrhenius acids and bases, they're not limited to aqueous environments. So we're expanding out from a solvent of water and moving into other solvents, other situations in which this definition still holds true. Now here, every Arrhenius acid is a Bronsted-Lowry acid. So think about it. Our example of HCl for Arrhenius acid.

Here, under the Arrhenius acid, we said it broke up into 2 ions to give us H⁺. Now under the Bronsted-Lowry model, what's really going on is HCl is donating its H⁺ to water to give us H₃O⁺ which is the same thing as H⁺. That's why every Arrhenius acid is a Bronsted-Lowry acid. Now every Arrhenius base is also a Bronsted-Lowry base. If we had NaOH, that OH^- which is the base since it's negative, it could easily accept an H⁺ ion from HCl if it had to. And because it can accept an H⁺ from HCl, it constitutes a Bronsted-Lowry base.

Now our Bronsted-Lowry acids and bases are just called our conjugate acid-base pairs when we're talking about a reaction. Hopefully, this is just a quick review for all of you. This was a predominant portion of general chemistry when we talked about acid-base reactions. We're just getting the definitions out of the way before we head on to more advanced calculations. Here, remember, Arrhenius was the simplest definition of an acid and base.

Bronsted-Lowry is just taking a few steps further into talking about more complex situations that don't involve water as a solvent. We still have another form of acids and bases that still remain after this one. But for right now, we're focused on our acids being proton donors and our bases being proton acceptors. Now that we've got the fundamentals down, try to do example 1. Hopefully, you guys remember these types of questions from Gen Chem in order to get the correct answer for example 1.

If you don't quite remember, don't worry. Just come back and we'll quickly go over this example.

Bronsted Acid-Base

So here, we have to write the formula of the conjugate base for the following compound. Since the compound is acting as an acid, to find its conjugate base, we need to remove an H⁺ from the compound. Remember, when asked for the conjugate base, just remove an H⁺ from the compound, and you'll have your conjugate base. So, by removing an H⁺ from it, we get SO₃.

Remember, we're not only losing a hydrogen but also losing a positive charge. Here, it's already negative one. Losing a positive charge means it becomes more negative, so it becomes SO₃²⁻. In actuality, what's happening is that we have HSO₃⁻ reacting with water. It is acting as the acid.

Water is acting as the base. It's going to donate an H⁺ away, and that's how it becomes SO₃²⁻. In the process, water accepts an H⁺ to become H₃O⁺. Here, the sulfate ion represents the conjugate base of bisulfite, which is HSO₃⁻. Now that you've seen that example, think of it the opposite way in which we have to find the conjugate acid of the example given below.

Once you do that, click on the next video and see if your answer matches up with mine.

Bronsted Acid-Base

Here, we have to write the conjugate acid for the compound given. If we want its conjugate acid, that must mean that it acts as a base and accepts an H⁺. Here, by accepting an H⁺, it becomes H₂PO₃. Here, you're not only accepting a hydrogen but you're accepting a +1 charge as well. Its charge is already 2 minus.

By accepting a plus one, it becomes more positive. Now it's H₂PO₃. What's happening is that it's reacting with water. It is acting as the base now, so water is acting as the acid. Water is going to donate an H⁺ to it.

That's how it becomes H₂PO₃⁻. In the process, by losing an H⁺, water becomes OH⁻. We'd say here that this is called dihydrogen phosphite, H₂PO₃⁻, which represents the conjugate acid of our starting material of HPO₃²⁻. Now that we've seen those two examples, put it all together and classify each one of the following compounds under the words provided. Once you do that, come back and see how your answer matches up with mine.

Bronsted Acid-Base

This question is pretty straightforward in terms of recognizing a Bronsted Lowry acid and a Bronsted Lowry base and its conjugates. Here, we simply have to identify the acid, base, conjugate acid, and conjugate bases in the following reaction. Here we have HF. Here we have water. In the process of donating and accepting an H⁺, we produce fluoride ion and the hydronium ion.

We can see that HF transitions into F^-. It lost an H⁺, so it has to represent the acid. Water accepts an H⁺ to go from water to H₃O⁺ so it must be the base. Remember, whatever you are as a reactant, you're the opposite as a conjugate. If HF is the acid, F^- would be its conjugate base. If water is the base, H₃O⁺ is its conjugate acid.

Now realize that HF and F^-, they're different by one hydrogen so they would represent the conjugate acid-base pair to one another. Water and H₃O⁺ are different by one hydrogen, 1 H⁺, so they would represent another conjugate acid-base pair. The idea of Bronsted and Lowry is pretty simple and straightforward. This is important when it comes to calculating the pH of weak acid and weak base solutions because you have to think of the reactions in this format. You have your acid and bases as reactants and your conjugates as products.

This will help later on with the formation of ice charts in order to calculate the concentration of H⁺ or OH^- to, at the end, determine the pH or pOH of solutions. Now that you've got that down, move on to the last classification of acids and bases.

Here’s what students ask on this topic:

What is the difference between Arrhenius and Brønsted-Lowry definitions of acids and bases?

The Arrhenius model defines acids as substances that increase the concentration of H⁺ ions in aqueous solutions and bases as substances that increase the concentration of OH^- ions. In contrast, the Brønsted-Lowry theory defines acids as proton (H⁺) donors and bases as proton acceptors. This definition is broader and not limited to aqueous solutions. For example, HCl is an Arrhenius acid because it increases H⁺ in water, and it is also a Brønsted-Lowry acid because it donates a proton to water, forming H₃O⁺. Similarly, NaOH is an Arrhenius base because it increases OH^- in water, and it is a Brønsted-Lowry base because OH^- can accept a proton.

What are conjugate acid-base pairs in the Brønsted-Lowry theory?

In the Brønsted-Lowry theory, a conjugate acid-base pair consists of two species that transform into each other by the gain or loss of a proton (H⁺). When an acid donates a proton, it becomes its conjugate base, and when a base accepts a proton, it becomes its conjugate acid. For example, in the reaction between HCl and water, HCl donates a proton to form Cl^- (its conjugate base), and water accepts a proton to form H₃O⁺ (its conjugate acid). This concept is crucial for understanding acid-base reactions and equilibrium in various chemical contexts.

Can Brønsted-Lowry acids and bases exist in non-aqueous solutions?

Yes, Brønsted-Lowry acids and bases can exist in non-aqueous solutions. Unlike the Arrhenius model, which is limited to aqueous environments, the Brønsted-Lowry theory applies to any solvent system. This broader definition allows for the study of acid-base reactions in various solvents, such as ammonia, acetic acid, and others. For instance, in liquid ammonia, NH₃ can act as a base by accepting a proton to form NH₄⁺, demonstrating the flexibility of the Brønsted-Lowry model beyond water-based systems.

How does the Brønsted-Lowry theory explain the behavior of HCl in water?

According to the Brønsted-Lowry theory, HCl behaves as an acid in water by donating a proton (H⁺) to a water molecule. This proton transfer results in the formation of hydronium ions (H₃O⁺) and chloride ions (Cl^-). The reaction can be represented as:

${\mathrm{HCl}}_{}+{H2O}_{}\to {H3O}_{}+{\mathrm{Cl}}_{}-$

In this reaction, HCl is the proton donor (acid), and water is the proton acceptor (base). The resulting H₃O⁺ and Cl^- are the conjugate acid and conjugate base, respectively.

Why is every Arrhenius acid also a Brønsted-Lowry acid?

Every Arrhenius acid is also a Brønsted-Lowry acid because both definitions involve the release of H⁺ ions. In the Arrhenius model, an acid increases the concentration of H⁺ ions in aqueous solution. In the Brønsted-Lowry model, an acid is defined as a proton donor. Since releasing H⁺ ions (protons) into solution is essentially donating protons, any substance that qualifies as an Arrhenius acid will also meet the criteria for a Brønsted-Lowry acid. For example, HCl increases H⁺ in water (Arrhenius) and donates a proton to water (Brønsted-Lowry).