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Brønsted–Lowry acids and bases

In the Brønsted–Lowry definition of acids and bases, an acid is a proton (H⁺) donor, and a base is a proton acceptor. When a Brønsted–Lowry acid loses a proton, a conjugate base is formed. Similarly, when a Brønsted–Lowry base gains a proton, a conjugate acid is formed. A Brønsted–Lowry acid (or base) and its conjugate base (or acid) are known as a conjugate acid–base pair. Created by Sal Khan.

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  • purple pi purple style avatar for user Sreeramsmallan
    what is the definition of a conjugate base & acid
    (44 votes)
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  • piceratops tree style avatar for user SalomonJohn123
    How come the Chlorine gained 2 extra electrons when hydrogen only has one. Also what made the Chlorine negative.
    (24 votes)
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  • area 52 blue style avatar for user No
    Are there more kinds of acids & bases?
    (10 votes)
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    • starky sapling style avatar for user Dave!
      I believe you might be referring to the proposed or suggested title of the acids/bases in this video as "Bronsted Lowry acids/bases." There is no particular acid or base that is a Bronsted Lowry acid or base; it is simply those researchers' way of describing the behavior of acids and bases in general, and what defines them as each.

      There is another popular description of acids and bases which complements the Bronsted Lowry model called "Lewis Acids and Bases." Lewis focused on the transfer of electrons rather than Hydrogen ions, so essentially it's the same thing, just from a different perspective.

      Of course, there are going to be instances where only one of these behaviors (proton movement or electron movement) is observed and, in those cases, an acid or base would only be EITHER a Bronsted-Lowry acid/base OR a Lewis acid/base.

      I might be way off base here with exactly what you're asking, but I know this hung me up for a bit when I was first learning about Bronsted-Lowry and Lewis. I hope that helps.
      (25 votes)
  • leaf green style avatar for user wannabeDoc
    Are all acids and bases about gaining or loosing Hydrogen (proton) or Hydroxide (electron)?
    (5 votes)
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    • blobby green style avatar for user martha halihan
      There are also Lewis acids and bases, which is a different theory about acid base chemistry. The Bronsted-Lowery definition refers to the loss or gain of an H+ (proton). The acid is a proton donor, and the base is a proton acceptor. The Arrhenius definition of an acid is an H+ producer and the base is an OH- producer. This approach is more limited than the Bronsted-Lowery theory.
      (12 votes)
  • duskpin sapling style avatar for user Maggie
    In my science textbook, the Arrhenius definition is used to define acids, bases and salts. However, it seems like Khan Academy uses the Bronsted-Lowry definition. Which definition is more commonly used and why?
    (7 votes)
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    • leafers ultimate style avatar for user William H
      It really depends on what type of chemistry you're doing or whatever your teacher prefers or what you're talking about. For example is talking about biological buffer solutions (carbocate+bicarbonate in blood) you'd use Arrhenius since it's in water and that's the easiest. If you however were talking about ammonia and boron trifluoride (bear with me, i know it's a bit esoteric), and ammonia donates a pair of electrons and acts as a base without OH- being involved, you would use the bronsted-lowry definition.
      (10 votes)
  • old spice man blue style avatar for user Sean McCrea
    Uh i'm still unsure about Hydronium could someone tell me about in a simple way?
    (6 votes)
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  • winston baby style avatar for user roanaldo
    At around Sal said that the chlorine was an anion with anegative charge, but if the atom has 8 electrons won't the ion become an atom with no net charge or a neutral charge?
    (3 votes)
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  • leaf green style avatar for user Jeffrey Feng
    At , a note at the bottom corrects "Covalent Bond" to Dative Bond. Can anyone elaborate for me?
    (3 votes)
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    • leaf red style avatar for user Richard
      Dative bonds, or coordinate bonds, are similar to covalent bonds in that the atoms which form the bonds share the electrons. But unlike in covalent bonding where each of the two atoms contribute half of the bonding electrons, dative bonding involves one of the two atoms contributing all of the bonding electrons.

      So in the example at the oxygen atom donates both electrons in the form of a lone pair to the hydrogen to form a new single bond. Since the hydrogen didn't contribute any electrons for the new bond, it would be classified as a dative bond.

      Dative bonding is more common when dealing with Lewis acids/bases and coordination compounds involving transition metals.

      Hope that helps.
      (4 votes)
  • aqualine sapling style avatar for user Ladysaint3
    Are acids and/or bases particularly a negative or positive ion or molecule?
    (3 votes)
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    • female robot grace style avatar for user tyersome
      Not exactly.

      These are all acids:
      NH₄⁺, H₃O⁺, HCl, H₂SO₄, HSO₄¯, H₂C₂O₄ (oxalic acid), HC₂O₄¯

      These are all bases:
      NH₃, OH¯, SO₄²¯, C₂O₄²¯
      I can't think of any net positively charged compounds that act as bases, but I'm sure they exist§.

      So you could say that positively charged ions are less likely to act as bases, and negatively charged ions are less likely to act as acids.

      Note that it probably better to think of compounds acting as acids or bases in a particular reaction, since many compounds can do both depending on the circumstances. Also, many compounds (e.g. amino acids) have both positive and negative charges ...

      §ADDENDUM: At pH < 2.18 lysine will have a net charge greater than +1 so lysine⁺ (where both amino groups are protonated, but the carboxyl is deprotonated) can act as a base.
      (4 votes)
  • piceratops seedling style avatar for user Ng Chaw Chuen
    so can the reaction move backwards?
    (4 votes)
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Video transcript

- [Voiceover] You've probably heard the term acid used in your everyday life. But what we want to do in this video is get a more formal definition of an acid. And particular, we'll focus on the one that is most typically used. Although we'll see future videos that there's other fairly common definitions of acids used as well beyond the one that we're going to see here. But the one that we're going to focus on is the Bronsted-Lowry definition. The Bronsted-Lowry definition of acids and bases. And this is a picture of Bronsted. This is a picture of Lowry. And they came up with this acid-base definition in the 1920s. So, we're going to do the Bronsted-Lowry, Bronsted-Lowry definition, definition of acids and bases. So, according to them, according to them, an acid, an acid is a proton, proton, or instead of writing proton we could actually write hydrogen ion donor. So why is a proton and a hydrogen ion the same thing? Well, in the most common isotope of hydrogen, we would, in it's nucleus, we would find just a proton and no neutron. And if it's neutral, you would have an electron buzzing around, jumping around in its orbital. So, you would have it's electron jumping around in its orbital. But if you were to ionize it, you're getting rid of its electron. So, if you're getting rid of it's electron, so, if you're getting rid of this, all you're going to be left with is a proton. So that's why a proton, an H plus, is usually referring to the exact same, is referring to the exact same thing. So, that's what an acid is. So what would a base be? Well, you could imagine by this definition A base, a base would be a proton, would be a proton, or you could say a hydrogen ion acceptor, acceptor. So let's make this a little bit more tangible with some examples. So one of the stronger acids we know is hydrochloric acid. Let me, let me draw. So, it's a hydrogen having a, having a covalent bond. Having a covalent bond with chlorine. With chlorine, with chlorine right over there. And if we want to, let's draw actually chlorine's lone pairs. So outside of the electron that is contributing to this pair in the covalent bond. It also has, it also has three other lone pairs. It also has three other lone pairs, just like that. So, if you were to take hydrochloric acid, place it in an aqueous solution, so it's in an aqueous solution right over here. And actually an aqueous solution, you'll see this written like that. That just means it's in a solution of water. So you could write like this, you could write hey, hydrochloric acid in an aqueous solution if you want to make it a little bit more explicit. You could say hey, look, this is going to be around some water molecules in its liquid form. Aqueous solution just means it's dissolved in liquid water. So, some water molecules in their liquid form. So, this is a water molecule. Whoops, water molecule. Right over here. So, an oxygen bonded to two hydrogens. And sometimes you'll see it written like this, that it's in its liquid, it's in its liquid form. Well, what do you think is going to happen? Well, I already said that this is a strong acid right over here. So this is going to really want to donate protons. It's really going to want to donate this hydrogen, but not let the hydrogen keep its electrons. So what's likely to happen here? Well, the both of these electrons in this pair are going to be grabbed by this chlorine. And then this hydrogen ion, because its electron was grabbed, well this could be nabbed by some water molecule passing by. Remember, in a real solution, it's not like they know what to do. They're just all bumping past each other. And based on how badly they want to do things, these reactions happen. And so you can imagine this lone pair right over here, well maybe it's able to form a covalent bond with this hydrogen. And so what's going to happen? What's going to happen? And I'll draw it with just an arrow because this reaction favorably goes, very strongly goes to the right, because this is such a strong acid. Well, then you're going to be left with, you're gonna be left with, the chlorine is now going to have its three lone pairs that it had before. And then it also grabbed these two electrons right over here. It also grabbed those two electrons right over there, so it gained an extra electron. It now has a negative charge. It is now the chloride anion. So it has a negative charge. And what about this water molecule? Well this water molecule, you have your oxygen, you have your hydrogens, you have your hydrogens, but now you don't just have two hydrogens, you grabbed this hydrogen right over here. And maybe I'll do this hydrogen in a slightly different color so that you could keep track of it. You have this hydrogen right over there. And this lone pair, this lone pair you can view it as now forming this covalent bond. You had your other two covalent bonds to the other two hydrogens. And then you still have this lone pair right over here. You still have that lone pair sitting right over there. And what just happened? Well, this water molecule just gained a proton. This hydrogen did not come with an electron. So if you just gain a proton, you are now, if you were neutral before, you are now going to have a positive charge. So what just happened? You put hydrochloric acid in a water solution, in an aqueous solution, this thing has donated a proton to a water molecule. And so, what is the acid and what is the base here? Well, when we look at the reaction this way, we see that this is the acid, the hydrochloric acid, it's literally called hydrochloric acid. And here, water is acting as a base. Water is acting as a base. And as you could see, water can actually act as an acid or a base. So, water is acting as a base. Now you might be saying, okay, this reaction goes strongly to the right, hey, but like you know, I could imagine in certain circumstances where chloride might accept a proton because it has this negative charge. And you would be right. This reaction goes strongly to the right, but once an acid has donated its proton, the thing that is left over, this is called a conjugate base. And I'll do the same color. So, this is the conjugate base of hydrochloric acid. The chloride anion. Conjugate, conjugate base of hydrochloric acid. And this right over here is the conjugate acid because you could imagine this hydronium ion, this could, under the right circumstances, donate protons to other things. Donate a hydrogen without donating electron to other things. And so this is actually the conjugate acid of H2O. Conjugate acid of water, of a water molecule. And as we'll see, water can act as an acid or a base. But this this gives you a kind of a baseline of at least the Bronsted-Lowry definition of acids and bases. And actually, one other thing I want to add. In some books here, so over here I said, hey, put this in an aqueous solution you're gonna form some hydronium, sometimes you'll see it written like this. And I'll just write it a little bit, a little bit, sometimes you'll see it like this. So you have your hydrochloric acid, and I won't draw the details this time, in an aqueous solution. So it's in a solution of water. And they'll just draw the reaction going like this, where they say hey, you're gonna be left with, you're gonna be left with some hydrogen ions, these protons. And you're going to be left with, and actually we could say it's gonna be in a aqueous solution, aqueous solution. And you're gonna be left with some chloride anions. Some chloride anions and it's in an aqueous solution. Now this isn't incorrect, but it's important to realize what they're talking about when they're talking about these hydrogen ions right over here. We know that if you have the hydrogen ions in an aqueous solution they don't just hang out by themselves. They get grabbed by a water molecule and they form hydronium. So, it's much more, I guess, it's much more close to the actual of what's happening, is if you actually talk about hydronium forming. As opposed to just the protons. 'Cause these protons in an aqueous solution, in a water solution, they're gonna be grabbed by a water molecule to form hydronium. And that's why I did it the way, this way up here.