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

In 1923, Johannes Brønsted and Thomas Lowry redefined acids and bases, where an acid is a proton (H⁺) donor and a base is a proton acceptor. This Brønsted-Lowry definition expands on the Arrhenius definition, which is more restrictive, only considering substances that produce H⁺ or OH⁻ ions in aqueous solutions as acids and bases, respectively. Brønsted-Lowry acids and bases form conjugate pairs; when a base accepts a proton, it becomes a conjugate acid, and when an acid donates a proton, it becomes a conjugate base. There is an inverse relationship between the strength of an acid or base and its conjugate: strong acids have weak conjugate bases, and weak acids have relatively strong conjugate bases. This relationship is pivotal because it influences the extent of ionization and the direction of chemical equilibrium in acid-base reactions. Understanding this concept is crucial for predicting the behavior of acids and bases in various chemical contexts.

Bronsted-Lowry Acids vs Bases

Here we're going to say that in 1923, Johannes Bronsted and Thomas Lowry developed a new definition for acids and bases. Here we're going to say that the Bronsted-Lowry definition is that it is a proton donor, so acids donate H⁺ to water, producing H₃O⁺ ion, which is coined the hydronium ion.

Here we take a look. We have hydrobromic acid reacting with water. Here, hydrobromic acid is serving as our Bronsted-Lowry acid, so it's going to donate H⁺ visually. If you want to think about it, you can think of it like this, like it's giving away its H⁺, and in doing that, H₂O accepts it. So H₂O is acting as our base. What's the result of these actions? Well, we're going to say here that HBr gave away an H, so all that's left of it is Br^-. H₂O accepted an H and as a result, it creates H₃O⁺.

Now a Bronsted-Lowry base is a proton acceptor. If we take a look here, we have ammonia reacting with water. In this case, ammonia acts as the base, so it's going to accept H⁺, and it is water that will be donating the H. So visually, think of it. H is going to go here. What's going to happen as a result? Well, NH₃ picks up an H⁺ and becomes NH₄⁺ as a product. So here water is acting as the acid and then water gave away an H⁺. So what's left of it? Well, what's left is OH^-. So that's what we have left of our water molecule.

So just remember, this goes a little bit further into our understanding of what constitutes an acid, what constitutes a base. An acid will donate H⁺ and a base will accept H⁺.

Bronsted-Lowry Acids and Bases Example

Identify each compound as either a Bronsted-lowry acid or base. So for the first one we have ammonium ion. This represents a positively charged amine. Remember, positively charged amines are weak acids. So here this will be a Bronsted-lowry acid.

For the second one, we have lithium hydride. We have a group 1A metal connected to the hydride ion. This represents an ionic base in it. It's the hydride ion which is negatively charged that would accept an H⁺. So this would be a base.

Next, we have CH₃NH₂. This is methylamine. It's a neutral amine. Neutral amines are weak bases. So this is a Bronsted-lowry base. It could accept an H⁺, and if it did it would create CH₃NH₃⁺.

Finally, we have H₂Te. This represents a binary acid, so it could donate one of its H⁺. If it did that, it would become HTe^-. So this is how we classify each of the compounds given to us within this example question.

Comparing Bronsted-Lowry and Arrhenius

Now let's take a look at some questions that try to relate Arrhenius acids and bases to Bronsted-Lowry acids and bases. The first question says are all Arrhenius acids and bases considered Bronsted-Lowry acids and bases. The answer to this question would be a yes. So for example we have hydrobromic acid. Under the Arrhenius definition it would be an acid since it produces H+1 ion, but it would also be a Bronsted-Lowry acid because it is donating an H+1 ion. So under that definition it would be both an Arrhenius acid and a Bronsted-Lowry acid.

Now sodium hydroxide. Here sodium hydroxide under the Arrhenius acid and base definition is a base because it produces OH-1. Now this OH-1 could accept an H from the water molecule surrounding it. It would also be a Bronsted-Lowry base. Now are all Bronsted-Lowry bases considered Arrhenius bases? The answer here would be a no. We're going to say bases that do not contain OH hydroxide ion are not Arrhenius bases. Here we have ammonia ion. When you put it into water, it doesn't immediately produce OH-1 in the way that we are accustomed to seeing under the Arrhenius definition. Yes, it could accept an H+1 ion from water to create OH-1, but it itself is not releasing OH-1 into the solution.

Now finally, are all Bronsted-Lowry acids considered Arrhenius acids? This would be a yes, because what is a Bronsted-Lowry acid? It is an H+1 donor, so it has H+1 present, meaning that if I put it within an aqueous solution it will release H+1 ions, so it would be an Arrhenius acid too. From this information we can see that Bronsted-Lowry takes a much more broader view of acids and bases than Arrhenius does. Arrhenius is much more limited in its scope. Bronsted-Lowry is much more broad, so we can fit even more compounds under the definition of either an acid or a base.

Bronsted-Lowry Acids and Bases Example

Identify each compound as either a Bronsted-Lowry acid or base, Arrhenius acid or base or both. So here in the first one we have carbonic acid. Carbonic acid can dissociate into H⁺ ions when we place it within an aqueous solution. This would make it an Arrhenius acid. At the same time synthetic can dissolve into H⁺ ions. When placed into an aqueous solution, it could technically donate that H⁺ as well, so it also represent a Bronsted-Lowry acid. So here we're going to say it's both an Arrhenius acid and a Bronsted-Lowry acid. So I'll say it's both.

Next we have methylamine. This is a neutral amine. Remember, neutral amines represent weak bases. Because it is a neutral amine it could accept an H⁺ and therefore represents a Bronsted-Lowry base. However, it is not a metal hydroxide. Dissolving it in an aqueous solution does not allow it to release an OH^- from itself. So it would not be an Arrhenius base. So here it is only a Bronsted-Lowry base, which one abbreviate as BLB.

Next we have potassium amide. This represents a strong base. Because it is a base, it could accept an H⁺ ion from the water molecules around it. Once you place it within an aqueous solution so it can serve as a Bronsted-Lowry base. Again, it is not composed of a metal hydroxide, therefore it doesn't release OH^- from its dissolution. So we couldn't say that it is Arrhenius base. So this is just going to be a Bronsted-Lowry base.

Next we have strontium hydroxide. It is the metal hydroxide, so it is an Arrhenius base. It also has the presence of OH^-, which could accept an H and therefore it represents a Bronsted-Lowry base. So we're going to say here this is an Arrhenius base plus a Bronsted-Lowry base. So we'll say it's both.

Next we have hydrofluoric acid HF. It could release H⁺ when dissolved in a solution, so it represents an Arrhenius acid. And because it can release an H ion, it could donate that H ion. So it is also a Bronsted-Lowry acid. So we're going to say it's an Arrhenius acid plus a Bronsted-Lowry acid. So it's both.

And then finally we have calcium hydride. Calcium hydride is a strong base here. The presence of the hydride ion H^-, it could accept an H⁺ ion. So it is a Bronsted-Lowry base, but it is not a metal hydroxide. So we won't say that it represents an Arrhenius base. So here we're just going to say it's a Bronsted-Lowry base and that's how we classify each one of these compounds.

Conjugate Acid-Base Pairs

Now, Bronsted-Lowry acids and bases occur in what we call conjugate pairs. We're going to say when a base accepts a proton, remember proton is H+, it transforms into a conjugate acid. So basically to create a conjugate acid we have to add an H+ to a base.

Here we have methyl amine. Here I've drawn lone pairs on the nitrogen. It possesses those lone pairs because when methylamine gained an H+, it goes to the nitrogen carbons already making all the bonds it can make. So by accepting this H+, we get now CH₃NH₃⁺. This is methyl ammonium ion. This represents the conjugate acid.

Now, when an acid donates a proton, it transforms into a conjugate base. So basically, when you're removing H+ from an acid, you create a conjugate base. Here we have nitric acid. We're going to remove an H+ from it, and when we do that, we're going to create the nitrate ion. The nitrate ion is the conjugate base of my nitric acid, right.

So just keep this in mind when we're talking about Bronsted-Lowry acid definitions of acids and bases. We're going to say here if you are adding an H+ to a base, you create a conjugate acid. If you're removing an H+ from an acid, you're creating a conjugate base.

Bronsted-Lowry Acids and Bases Example

Here it says to provide the formulas of the conjugates for each of the following compounds. So remember, if we're given an acid, that means we have to show the conjugate base. If we're given a base, then we have to show the conjugate acid.

If we take a look here at the first one and H₂ and H₂, that is a neutral amine. Its name is hydrazine, but you don't need to know that it's a neutral amine. Therefore it is a weak base. Since it is a base, we have to show the conjugate acid. Remember to give the conjugate acid you add an H to it. So here we would get NH₂NH₃⁺. Or you could just put the H on the first nitrogen. Either one would be correct.

Next, we have looks like an oxy acid, so this is formic acid. Since it's an acid, we have to show the conjugate base, so we have to remove an H. So when we do that, we're going to get CHO₂^-.

Next we have hydrogen sulfate. This compound is amphoteric or amphiprotic. It has the presence of a hydrogen in the front and a negative charge. Remember, amphiprotic or amphoteric species can act as an acid or base depending on what they're reacting with. And because of this I have to specify, does this amphoteric species represent a base or an acid? Here I'm saying it represents a base, which means you need to show the conjugate acid. To show the conjugate acid you add an H to it. So here the conjugate acid would be H₂SO₄ sulfuric acid.

And then finally we have here chlorous acid. It is an acid so I need to show the conjugate base so remove an H⁺. When I do that I get ClO₂^-. I get the chlorite ion right? So these will represent each of the conjugates for the following compounds.

Identify conjugate acid and conjugate base in the following reaction.

                                                    H₂PO₄⁻ (aq) + H₂O (l) ⇌ HPO₄²⁻ (aq) + H₃O⁺ (aq)

a) HPO₄²⁻ (conjugate acid), H₃O⁺ (conjugate base) 

b) HPO₄²⁻ (conjugate base), H₃O⁺ (conjugate acid)

c) H₂PO₄⁻ (conjugate acid), H₂O (conjugate base)

d) H₂PO₄⁻ (conjugate base), H₂O (conjugate acid)

• A. HPO₄²⁻ (conjugate acid), H₃O⁺ (conjugate base)
• B. HPO₄²⁻ (conjugate base), H₃O⁺ (conjugate acid)
• C. H₂PO₄⁻ (conjugate acid), H₂O (conjugate base)
• D. H₂PO₄⁻ (conjugate base), H₂O (conjugate acid)

In the following reaction, label Bronsted-Lowry acid and base, along with conjugate acid and base.

                               

H₂C₆H₆O₆ (aq) + H₂O (l) ⇌ HC₆H₆O₆⁻ (aq) + H₃O (aq)

           

• A. H₂C₆H₆O₆ conjugate acid
  H₂O conjugate base
  HC₆H₆O₆⁻ base
  H₃O⁺ acid
• B. H₂C₆H₆O₆ acid
  H₂O base
  HC₆H₆O₆⁻ conjugate base
  H₃O⁺ conjugate acid
• C. H₂C₆H₆O₆ conjugate base
  H₂O conjugate acid
  HC₆H₆O₆⁻ acid
  H₃O⁺ base
• D. H₂C₆H₆O₆ base
  H₂O acid
  HC₆H₆O₆⁻ conjugate acid
  H₃O⁺ conjugate base

Strength of Conjugate Acids and Bases

We're going to say there's an inverse relationship between strengths of acids and bases and their conjugates. A rule of thumb is a strong acid will have a relatively weak conjugate base, and in fact the stronger the acid the weaker the conjugate base. And a weak conjugate base has a low affinity for proton. That means it doesn't want to accept an H⁺.

So here we have HCl reacting with water. HCl is the acid so it's going to donate an H⁺ to the water. Water becomes H₃O⁺. As a result, HCl lost an H to what's left is Cl^- its conjugate base.

Remember, if you're strong acid, that means your conjugate base will be weak. So here we're going to have a weak conjugate base as a product.

Conjugate Acid-Base Relationships

All right, so remember there's an inverse relationship. So that means a weak acid will have a relatively strong conjugate base. And in fact, the weaker the acid, the stronger the conjugate base. A stronger conjugate base has a high affinity for protons, so it has it wants some more readily accepted H.

Here we have HCM, which is a weak acid hydrocyanic acid. It reacts with water. Water accepts an H⁺ and becomes H₃O⁺. As a result, HCM gave away an H⁺ analog C and minus. Here it's a weak acid, so that means it's going to be a stronger conjugate base. It's still weak, but it's stronger because it came from someplace else that's weak.

Now notice also that we have reversible arrows here. The arrow that points to the reactant is longer. That means reactants are more highly favored. This makes sense because remember, weak acids and weak bases don't completely ionize. We don't make 100% of these ions. A vast majority of it is still in the reactant form, which is why the arrows pointing towards the reactant, that side is more favored. There's more of it.

So we're going to say here, stronger the base, the weaker the conjugate acid. We conjugate acids less readily donate protons. The weaker the base then the stronger the conjugate acid. Stronger acids more readily donate protons or more easily donate protons to. Just remember this inverse relationship. If you're strong in one particular way, you're weaker in the opposite way. A strong acid would equate to a weaker conjugate base.

Acid and Base Strength Example

Which of the following acids will have relatively strong conjugate bases? So remember a strong conjugate base comes from a weak acid. So basically we have to look and see which one here is a weak acid.

So here we have per bromic acid which is one of the strong acids. So this wouldn't work. Hydrocyanic acid is a weak acid so this would be an answer. Nitric acid is a strong acid so that wouldn't work. And then we have perchloric acid which is a strong acid so that wouldn't work.

Remember a weak acid would equate to stronger or strong conjugate basis. So here the only answer that works is option B.

Which of the following is the strongest base?

a. NO₃⁻                         b. F⁻                              c. Cl⁻                             d. ClO₄⁻                         e. H₂O

• A. NO₃⁻
• B. F⁻
• C. Cl⁻
• D. ClO₄⁻
• E. H₂O

Which of the following bases will have the weakest conjugate acid?

a. CH₃COOH         b. HCl            c. CH₃NH₂           d. (CH₃)₂NH              e. LiOH

• A. CH₃COOH
• B. HCl
• C. CH₃NH₂
• D. (CH₃)₂NH
• E. LiOH

Here’s what students ask on this topic:

How can one identify Bronsted-Lowry acids and bases?

To identify Bronsted-Lowry acids and bases, you need to look at how the substances behave in a chemical reaction. According to the Bronsted-Lowry theory, an acid is any substance that can donate a proton (which is a hydrogen ion, ^H+), and a base is any substance that can accept a proton.

When examining a chemical reaction, identify the species that loses a hydrogen ion; this species is the Bronsted-Lowry acid. Conversely, look for the species that gains a hydrogen ion; this one is the Bronsted-Lowry base. It's important to note that the reaction can be reversible, and in the reverse reaction, the roles of acid and base can switch. This is because the product of the acid, after it donates a proton, becomes a base capable of accepting a proton back (called the conjugate base), and the product of the base, after accepting a proton, becomes an acid capable of donating a proton (called the conjugate acid). Keep in mind that water can act as both a Bronsted-Lowry acid and a base, depending on the reaction it is involved in.

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

The Bronsted-Lowry definition of acids and bases is a conceptual framework used in chemistry to explain the behavior of acids and bases. According to this theory, an acid is a substance that can donate a proton (which is a hydrogen ion, ^H+) to another substance, while a base is a substance that can accept a proton. This definition expands the understanding of acids and bases beyond the earlier Arrhenius definition, which limited acids and bases to substances that produce ^H+ and ^OH- ions in water, respectively.

In the Bronsted-Lowry context, a proton transfer reaction occurs between an acid and a base. When an acid donates a proton, it becomes its conjugate base, which is the species that remains after the proton is lost. Conversely, when a base accepts a proton, it becomes its conjugate acid. This concept introduces the idea of conjugate acid-base pairs, where acids and bases are linked by the transfer of a proton. This theory is widely used because it applies to reactions in both aqueous and non-aqueous solutions.

What are Bronsted-Lowry acids?

Bronsted-Lowry acids are defined as substances that donate protons (hydrogen ions, H⁺) in a chemical reaction. According to the Bronsted-Lowry theory, an acid-base reaction involves the transfer of protons from the acid to another substance, which is called the base. The base, in this context, is defined as a proton acceptor. When an acid donates a proton, it forms its conjugate base, which is the species that remains after the proton is lost. Conversely, when a base accepts a proton, it forms its conjugate acid. This concept broadens the definition of acids and bases beyond those substances that produce H⁺ or OH⁻ (hydroxide ions) in water, allowing for the analysis of acid-base reactions in a wider range of chemical environments. The Bronsted-Lowry theory is fundamental to understanding chemical reactivity, pH, and the behavior of many solutions in chemistry.

What is the Bronsted-Lowry definition of a base?

The Bronsted-Lowry definition of a base is a substance that can accept a proton, which is a hydrogen ion (^H+). This definition expands the concept of bases beyond just those substances that can produce hydroxide ions (^OH-) in solution, which is how bases are defined according to the Arrhenius theory. In the Bronsted-Lowry framework, a base is therefore any substance that can bond to, or accept, a proton from another substance, which is correspondingly called a Bronsted-Lowry acid. This interaction results in the formation of a conjugate acid-base pair. For example, when ammonia (_NH3) acts as a Bronsted-Lowry base, it accepts a proton to form its conjugate acid, ammonium (_NH4⁺). This definition is particularly useful because it allows for the description of acid-base reactions in solvents other than water and includes reactions that do not produce ions.