Bronsted-Lowry Acids and Bases - Video Tutorials & Practice Problems
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concept
Bronsted-Lowry Acids vs Bases
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Here we're going to say that in 1923, Johannes Braunstead and Thomas Lawrie developed a new definition for acids and bases. Here we're going to say that the Bronsted Lowry acid definition is that it is a proton donor. So acids donate H+ to water producing H3O+ ion, which is coined the hydronium ion. Here if we take a look we have hydrobromic acid reacting with water. Here hydrobromic acid is serving as our Bronstonorium 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 in H+, and as a result it creates H3O+. Now, a bracidylori 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 plus is gonna go here. What's gonna happen as a result? Well, NH 3 picks up an H+ and becomes NH 4 positive 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 minus. 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+.
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example
Bronsted-Lowry Acids and Bases Example
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Identify each compound as either a Bronston Lauri 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 would be a Bronsted Lowry acid. For B, 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, CH3 NH2. This is Methylamine, it's a neutral amine. Neutral amines are weak bases, so this is a bromosinlauric base. It could accept NH+ and if it did, it would create CH3 NH3 plus 1. And then finally, here we have H2Te. This represents a binary acid, so we could donate one of its H pluses. If it did that it would become HTE minus. So, this is how we classify each of the compounds given to us within this example question.
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concept
Comparing Bronsted-Lowry and Arrhenius
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Now let's take a look at some questions that try to relate Arrhenius acids and bases to Brostin Lawrie 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 Aedes definition, it would be an acid since it produces h plus ion. But it would also be a Bronsted Lowry acid because it is donating an H+ion. So under that definition it would be both an Arrhenius acid and a Prostate Lauri acid. Now, sodium hydroxide. Here, sodium hydroxide under the Arrhenius acid and base definition is a base because it produces Oh minus. Now, this Oh minus could accept an H plus from the water molecule surrounding it, so it would also be a bronston lawrie base. Now, are all bronston lawrie bases considered Aurignac bases? The answer here would be a no. We're gonna say bases that do not contain Oh minus hydroxide ion are not Arrhenius bases. Here we have ammonia ion. When you put it into water, it doesn't immediately produce Oh minus in the way that we are accustomed to seeing under the Arrhenius definition. Yes, it could accept an H+ ion from water to create Oh-, but it itself is not releasing Oh minus into the solution. Now finally, are all Bronstonoriasis considered Arrhenius acid? This would be a yes. Because what is a Bronston Lauri acid? It is an H+ donor, so it has H+ present, meaning that if I put it within an aqueous solution it'll release H+ ions, so it would be an Arrhenius acid too. From this information, we can see that Bronson Lawrie takes a much more broader view of acids and bases than Arrhenius does. Arrhenius is much more limited in its scope. Brocellulari is much more broad, so we can fit even more compounds under the definition of either an acid or a base.
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example
Bronsted-Lowry Acids and Bases Example
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Identify each compound as either a Bronston Lawrie 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, since it can dissolve into H+ ions when placed into an aqueous solution, it could technically donate donate that h+ as well. So it also represent a Bronsted Lowry acid. So here we're gonna say it's both in Arrhenius acid and a Bronsted Lowry acid. So we'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 minus from itself so it would not be an Arrhenius base. So here it is only a Bronsted Lowry base which we're gonna 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 brancellore base. Again, it is not composed of a metal hydroxide, therefore it doesn't release o h minus from its, dissolution. So we couldn't say that it is a, Arrhenius base. So this is just gonna be a Broston Lohrey base. Next, we have strontium hydroxide. It is the metal hydroxide, so it is an Arrhenius base. It also has a presence of Oh- which could accept an H+ and therefore it represents a Bronsted and Lore base. So we're gonna say here this is an Arrhenius base, plus a Bronsted Lowry base, so let's 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+1 ion it could donate that H+1 ion. So it is also a Bronsted Lowry acid. So we're gonna 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 minus it could accept an H+ ion so it is a Bronci 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 base. And that's how we'd classify each one of these compounds.
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concept
Conjugate Acid-Base Pairs
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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 plus, it transforms into a conjugate acid. So basically to create a conjugate acid we have to add an H plus to a base. Here we have Methylamine. Here I've drawn lone pairs on the nitrogen, it possesses those lone pairs because when Methylamine gained an H plus it goes to the nitrogen. Carbons already making all the bonds it can make. So by accepting this H plus we get now CH3 NH3 positive. This is Methylammonium ion, this represents the conjugate acid. Now, when an acid donates a proton it transforms into a conjugate base. So basically when you remove an H+ from an acid you create a conjugate base. Here we have nitric acid. We're going to remove an H+ from it, 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 broxler and laurie 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.
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example
Bronsted-Lowry Acids and Bases Example
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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, n h two n h two, 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 NH2, NH3 positive, or you could just put the H+ on the first nitrogen. Either one would be correct. Next we have looks like an oxyacid, 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 minuteus. I get the Chloride ion. Right? So these will represent each of the conjugates for the following compounds.
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Problem
Problem
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)
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Problem
Problem
In the following reaction, label Bronsted-Lowry acid and base, along with conjugate acid and base.
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
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concept
Strength of Conjugate Acids and Bases
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We're gonna 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 plus. So here we have H c l reacting with water. H c l is the acid, so it's gonna donate an H+ to the water. Water becomes H 3 O plus as a result. H c l lost an h plus, so what's left is c l minus, it's conjugate base 4. Remember, if you're a strong acid, that means your conjugate base will be weak. So here we're gonna have a weak conjugate base as a product.
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concept
Conjugate Acid-Base Relationships
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1m
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Alright. 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 to more readily accept an H plus. Here we have HCM which is a weak acid, hydrocyanic acid, it reacts with water. Water accepts an H+ and becomes H three 0 plus as a result. HCM gave away an H+, now it's c n minus. Here it's a weak acid, so that means it's gonna be a stronger conjugate base. It's still weak, but it's stronger because it came from some place 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 a 100% of these ions. A vast majority of it is still in the reactant form, which is why the arrow is pointing towards the reactant. That side is more favored. There's more of it. So we're gonna say here, stronger the base, the weaker the conjugate acid. Weak 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. So 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.
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example
Acid and Base Strength Example
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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 perbromic acid, which is one of the strong 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 bases. So here, the only answer that works is option b.
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Problem
Problem
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
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Problem
Problem
Which of the following bases will have the weakest conjugate acid?