Bronsted Lowry Acid and Base: Study with Video Lessons, Practice Problems & Examples

In 1923, Brønsted and Lowry refined the definitions of acids and bases. They defined acids as proton donors (H⁺) and bases as proton acceptors. This contrasts with Arrhenius, who defined acids as substances that increase H⁺ concentration in water. Brønsted-Lowry acids and bases form conjugate acid-base pairs, differing by one hydrogen atom. For example, H₃O⁺ and H₂O are conjugate pairs. Understanding these concepts is crucial for grasping acid-base reactions and their applications in various chemical contexts.

Bronsted Lowry Acids & Bases Concept 1

Guys, in this new video, we're going to take a look at another type of acids and bases and probably the most important of the 3, Bronsted-Lowry. So here we're going to say it was in 1923 when Bronsted and Lowry developed a new definition for acids and bases. Arrhenius devised Arrhenius acids and bases near the end of the 1800s. And it was in the new century, 1923, where these two guys helped to make that definition a little bit better. We're going to say according to their definition, an acid was considered to be a proton donor. And when we say proton, we mean H⁺. So, this first definition still goes hand in hand with the Arrhenius definition because the Arrhenius acid is something that increases H⁺ concentration when dissolved in water. Bronsted-Lowry still agreed with that. They said that, "We agree an acid should have an H⁺." Where they disagreed though, Bronsted-Lowry didn't believe that a base needed an OH^- to be a base. What they said instead was, "if the acid donates H⁺, then the base must accept the H⁺." Their new definition for a base was, bases are proton, which means H⁺, acceptors. This is where they departed from Arrhenius. What we're going to say here is that unlike acids and bases, for Arrhenius, these could be used in solutions that were not just made up of water. And we're going to say here, Arrhenius acids say they're H⁺, they increase H⁺, Bronsted-Lowry says they give H⁺, so both agree. So you can say that Arrhenius acids are Bronsted-Lowry acids. Now, we're also going to say here that according to Bronsted-Lowry, a base accepts H⁺. Why would a base accept H⁺? Because it has lone pairs or it has a negative charge. Something positive, something negative are naturally attracted to one another. This kind of goes in line with Arrhenius a little bit because according to Arrhenius, we have to produce OH^-. OH^- is negative so it could accept H⁺. So there's a little bit of disagreement but also a little bit of agreement in terms of bases. Now we're going to say that Bronsted-Lowry acids and bases are always occurring in pairs which we call conjugate acid-base pairs. What you need to remember is that conjugate acid-base pairs differ by only one hydrogen. A good example is we have H₂O, we could have OH^-, we could have H₃O⁺. Here, these two are conjugate acid-base pairs. They're only different by 1 H. Water has 2 Hs, OH^- only has 1. And you're going to say that these two could be conjugate acid-base pairs as well. H₃O⁺ has 3 Hs. H₂O has 2. They're only different by 1 H⁺. So that's what we mean by conjugate acid-base pairs.

Bronsted Lowry Acids & Bases Example 1

Now, let's take a look at example 1. It says, write the formula of the conjugate base for the following compound. Now, the following compound is HSO₄^-. And realize when I say conjugate base, I mean remove H⁺. And since we're going to be using some of the spaces here guys to write graphs to help us follow along, I'm going to remove myself from the image so we can better focus. So remember, conjugate base means we'll remove H⁺. So just do a number line. This number line will help us figure out what the new charge is going to be. Because remember, you're not only moving an H, you're moving a positive H. So you're causing the charge to change once that H positive comes off. So it's important to track and see what your new charge will be. We start off as negative one. And if you're removing an H⁺, you're removing a positive, you're going to become more negative. So let's remove that H⁺. So now we're SO₄. Because we're going to become a little bit more negative, we fall down to minus 2. So, SO₄^2- would be the conjugate base of HSO₄^-.

Bronsted Lowry Acids & Bases Example 2

For example 2, we're doing the complete opposite. Now, we're looking for a conjugate acid, so that means adding NH⁺. Remember, adding a positive to you makes you more positive, so you become more positive. Number line again. Let's add an H to it. It becomes HVO₂. Just add the H to the front of the compound. That's all you have to do. Not to the center, to the side. Here it was negative 2. When I add an H⁺ it becomes more positive, so it jumps up to negative one. So, our compound now is HVO₂^-1. Now that you guys have seen how to do conjugate acids and bases, I want you guys to attempt to do practice question 1 and practice question 2. Good luck.

Write the formula of the conjugate base for the following compound:

H₂Se

• A.
• B.
• C.
• D.

Write the formula of the conjugate for the following compound:

NH₂NH₂

• A.
• B.
• C.
• D.

Bronsted Lowry Acids & Bases Example 3

Hey, guys. In this new video, we're going to put to practice some of the concepts we learned about Bronsted-Lowry acids and bases. So, let's take a look at the first example. Here it says identify the acid, the base, the conjugate acid, and the conjugate base based on the following reactions. Here we have HF aqueous plus H₂O aqueous which gives us F^- aqueous plus H₃O⁺. So, remember, what does a Bronsted-Lowry acid do? It gives away H⁺. A Bronsted-Lowry base, on the other hand, will accept that H⁺. So if we take a look, here we have HF. But then look, what happens to the HF? It becomes F^-. What must have happened? That HF gave away an H⁺ to the water. Because it gave its H⁺ away, it's the acid. The water on the other hand accepted that H⁺. That's how it became H₃O⁺ over here. So, here, this must be the base. Then we're going to say, HF gives away an H⁺ to become F^-; since we're taking away an H⁺, this must be the conjugate base. H₂O accepts an H⁺ to give us H₃O⁺, so H₃O⁺ must be the conjugate acid. We're going to say that these two are connected together as our conjugate acid-base pair, and then these two are connected together as our other set of conjugate acid-base pairs. Remember, we talked about this earlier. They only differ by 1 hydrogen. If they're different by 1 H, they're conjugates of each other.

Bronsted Lowry Acids & Bases Example 4

Now that we've done that one, let's take a look at example 2. In example 2, we have to do the same thing once again. So here, we have CN^- but all of a sudden it becomes HCN. How did that happen? The CN^- must have accepted an H⁺ because by accepting the H⁺ it's a base. Who's giving it that H⁺? It must have been the water. Water gives away an H⁺, making it an acid. When water gives away the H⁺, what happens to the water? The water becomes OH^-. Here, we would say that this is the conjugate base. Whatever you are, your conjugate is the opposite. If this is a base, this is a conjugate acid. Then based on that, we'd say that these two are conjugates of one another and then these two also. Just realize, water acts as a base in the first example but as an acid in the second. Something that can act as both an acid and a base, we said, was called amphoteric. Water is the best example of an amphoteric species. Depending on what it's next to, it could act as either an acid or a base. So always be careful. And we know that this has to be the base because it's negative. We're still using the rules we've learned before.

And if we go back up to water, actually, how do we know HF is the acid and H₂O is the base? Because remember, they're both going to have H connected to an electronegative element. F and O are both in the same period. Remember, we said when you're in the same period, what do we look at? We look at electronegativity. So HF is definitely a stronger acid than H₂O. As a result, HF must be the acid, H₂O must be the base. The rules we learned earlier play a huge role in what we're doing right now. Here, I gave us the product so it's easy to say who was the acid and who was the base. But on your exam, you may not get that luxury. Your professor might just give you HF plus water and ask you what are your products. So you still have to remember who would be the stronger acid. That person will be the acid. The other one would have to be the base.

Now that we've seen that, let's take a look at practice questions 1 and 2. For this one, we're asking, which of the following is an acid? I'll give you guys a huge help here. Remember, a Brønsted-Lowry acid has to have H⁺ in there. It has to have H. If it doesn't have H, it's out. Second, for something to give away an H, H needs to be connected to something that is electronegative. If H is connected to an electronegative element, it's going to make a polar bond. If that bond is polar, that means it's reactive. That means I can break it and take that H⁺ off. What bonds are nonpolar? Well, if H was connected to itself, this would be a nonpolar bond. So as a result, H₂ does not represent an acid. Also, who else? We can also say that when H is connected to C, their electronegativities are not that far off from one another. So their electronegativities are very similar. So we would expect that bond not to be polar. So as a result, if you have H connected to carbon, we would expect that to not be an acidic bond. So, I gave you guys two huge helps to help you figure out which one of these could be a Brønsted-Lowry acid. Use the rules that we've learned earlier to identify something as either an acid or a base. That'll help guide you to what's a Brønsted-Lowry acid for this particular question.

Which of the following is a Bronsted-Lowry acid?

• A. CH₄
• B. HCN
• C. NH₃
• D. Br₂

Determine the chemical equation that would result when carbonate, CO3^2-, reacts with water.

• A.
• B.
• C.
• D.

Here’s what students ask on this topic:

What is the Brønsted-Lowry definition of acids and bases?

The Brønsted-Lowry definition, introduced in 1923, describes acids as proton donors and bases as proton acceptors. In this context, a proton refers to a hydrogen ion (H⁺). This definition expands on the Arrhenius concept, which defines acids as substances that increase the concentration of H⁺ in water. The Brønsted-Lowry model is more versatile because it applies to reactions in non-aqueous solutions as well. For example, in the reaction between hydrochloric acid (HCl) and ammonia (NH₃), HCl donates a proton to NH₃, making HCl the acid and NH₃ the base.

How do Brønsted-Lowry acids and bases differ from Arrhenius acids and bases?

Brønsted-Lowry acids and bases differ from Arrhenius acids and bases primarily in their definitions and applicability. Arrhenius acids are substances that increase the concentration of H⁺ ions in water, while Arrhenius bases increase the concentration of OH⁻ ions. In contrast, Brønsted-Lowry acids are proton donors, and Brønsted-Lowry bases are proton acceptors. This broader definition allows Brønsted-Lowry theory to apply to reactions in both aqueous and non-aqueous solutions. For example, NH₃ (ammonia) is a Brønsted-Lowry base because it accepts a proton, but it is not an Arrhenius base because it does not produce OH⁻ in water.

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

In the Brønsted-Lowry theory, conjugate acid-base pairs are two species that differ by one 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 water (H₂O) and hydrochloric acid (HCl), H₂O acts as a base and accepts a proton to form H₃O⁺ (hydronium ion), while HCl donates a proton to become Cl⁻ (chloride ion). Here, H₃O⁺ and H₂O are a conjugate acid-base pair, as are HCl and Cl⁻.

Why is the Brønsted-Lowry theory important in chemistry?

The Brønsted-Lowry theory is crucial in chemistry because it provides a more comprehensive understanding of acid-base reactions. Unlike the Arrhenius definition, which is limited to aqueous solutions, the Brønsted-Lowry theory applies to a broader range of chemical environments, including non-aqueous solutions. This theory helps explain the behavior of acids and bases in various chemical reactions, including those in biological systems and industrial processes. It also introduces the concept of conjugate acid-base pairs, which is essential for understanding reaction mechanisms and predicting the outcomes of acid-base reactions.

Can you give an example of a Brønsted-Lowry acid-base reaction?

Sure! A classic example of a Brønsted-Lowry acid-base reaction is the interaction between ammonia (NH₃) and water (H₂O). In this reaction, water acts as a Brønsted-Lowry acid by donating a proton (H⁺) to ammonia, which acts as a Brønsted-Lowry base by accepting the proton. The reaction can be represented as follows:

$H_{2}O+{\mathrm{NH}}_3\to {\mathrm{OH}}_{\mathrm{-}}+{\mathrm{NH}}_4+$

In this reaction, H₂O donates a proton to NH₃, forming OH⁻ (hydroxide ion) and NH₄⁺ (ammonium ion). Here, H₂O and OH⁻ are a conjugate acid-base pair, as are NH₃ and NH₄⁺.