To install click the Add extension button. That's it.

The source code for the WIKI 2 extension is being checked by specialists of the Mozilla Foundation, Google, and Apple. You could also do it yourself at any point in time.

Kelly Slayton
Congratulations on this excellent venture… what a great idea!
Alexander Grigorievskiy
I use WIKI 2 every day and almost forgot how the original Wikipedia looks like.
What we do. Every page goes through several hundred of perfecting techniques; in live mode. Quite the same Wikipedia. Just better.

Brønsted–Lowry acid–base theory

From Wikipedia, the free encyclopedia

The Brønsted–Lowry theory is an acid–base reaction theory which was proposed independently by Johannes Nicolaus Brønsted and Thomas Martin Lowry in 1923.[1][2] The fundamental concept of this theory is that when an acid and a base react with each other, the acid forms its conjugate base, and the base forms its conjugate acid by exchange of a proton (the hydrogen cation, or H+). This theory is a generalization of the Arrhenius theory.

YouTube Encyclopedic

  • 1/5
    300 326
    122 383
    431 968
    83 284
    3 368
  • ✪ Bronsted-Lowry definition of acids and bases | Biology | Khan Academy
  • ✪ What is the Bronsted Lowry Theory | Chemistry for All | The Fuse School
  • ✪ Identify Conjugate Acid Base Pairs (Bronsted Lowry)
  • ✪ Conjugate Acid Base Pairs, Arrhenius, Bronsted Lowry and Lewis Definition - Chemistry
  • ✪ Bronsted - Lowry acid , base Concept


- [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.


Definitions of acids and bases

Johannes Nicolaus Brønsted and Thomas Martin Lowry, independently, formulated the idea that acids are proton (H+) donors while bases are proton acceptors.

In the Arrhenius theory acids are defined as substances that dissociate in aqueous solution to give H+ (hydrogen ions), while bases are defined as substances that dissociate in aqueous solution to give OH (hydroxide ions).[3]

In 1923 physical chemists Johannes Nicolaus Brønsted in Denmark and Thomas Martin Lowry in England both independently proposed the theory that carries their names.[4][5][6] In the Brønsted–Lowry theory acids and bases are defined by the way they react with each other, which allows for greater generality. The definition is expressed in terms of an equilibrium expression

acid + baseconjugate base + conjugate acid.

With an acid, HA, the equation can be written symbolically as:

The equilibrium sign, ⇌, is used because the reaction can occur in both forward and backward directions. The acid, HA, can lose a proton to become its conjugate base, A. The base, B, can accept a proton to become its conjugate acid, HB+. Most acid-base reactions are fast so that the components of the reaction are usually in dynamic equilibrium with each other.[7]

Aqueous solutions

Acetic acid, CH3COOH, is composed of a methyl group, CH3, bound chemically to a carboxylate group, COOH. The carboxylate group can lose a proton and donate it to a water molecule, H2O, leaving behind an acetate anion CH3COO− and creating a hydronium cation H3O+. This is an equilibrium reaction, so the reverse process can also take place.
Acetic acid, a weak acid, donates a proton (hydrogen ion, highlighted in green) to water in an equilibrium reaction to give the acetate ion and the hydronium ion. Red: oxygen, black: carbon, white: hydrogen.

Consider the following acid–base reaction:

Acetic acid, CH3COOH, is an acid because it donates a proton to water (H2O) and becomes its conjugate base, the acetate ion (CH3COO). H2O is a base because it accepts a proton from CH3COOH and becomes its conjugate acid, the hydronium ion, (H3O+).[8]

The reverse of an acid-base reaction is also an acid-base reaction, between the conjugate acid of the base in the first reaction and the conjugate base of the acid. In the above example, acetate is the base of the reverse reaction and hydronium ion is the acid.

The power of the Brønsted–Lowry theory is that, in contrast to Arrhenius theory, it does not require an acid to dissociate.

Amphoteric substances

The amphoteric nature of water
The amphoteric nature of water

The essence of Brønsted–Lowry theory is that an acid only exists as such in relation to a base, and vice versa. Water is amphoteric as it can act as an acid or as a base. In the image shown at the right one molecule of H2O acts as a base and gains H+ to become H3O+while the other acts as an acid and loses H+ to become OH.

Another example is furnished by substances like aluminium hydroxide, Al(OH)3.

, acting as an acid
, acting as a base

Non-aqueous solutions

The hydrogen ion, or hydronium ion, is a Brønsted–Lowry acid in aqueous solutions, and the hydroxide ion is a base, by virtue of the self-dissociation reaction

An analogous reaction occurs in liquid ammonia

Thus, the ammonium ion, NH+
, plays the same role in liquid ammonia as does the hydronium ion in water and the amide ion, NH
, is analogous to the hydroxide ion. Ammonium salts behave as acids, and amides behave as bases.[9]

Some non-aqueous solvents can behave as bases, that is, proton acceptors, in relation to Brønsted–Lowry acids.

where S stands for a solvent molecule. The most important such solvents are dimethylsulfoxide, DMSO, and acetonitrile, CH3CN, as these solvents have been widely used to measure the acid dissociation constants of organic molecules. Because DMSO is a stronger proton acceptor than H2O the acid becomes a stronger acid in this solvent than in water.[10] Indeed, many molecules behave as acids in non-aqueous solution that do not do so in aqueous solution. An extreme case occurs with carbon acids, where a proton is extracted from a C-H bond.

Some non-aqueous solvents can behave as acids. An acidic solvent will increase basicity of substances dissolved in it. For example, the compound CH3COOH is known as acetic acid because of its acidic behaviour in water. However it behaves as a base in liquid hydrogen chloride, a much more acidic solvent.[11]

Comparison with Lewis acid–base theory

In the same year that Brønsted and Lowry published their theory, G. N. Lewis proposed an alternative theory of acid–base reactions. The Lewis theory is based on electronic structure. A Lewis base is defined as a compound that can donate an electron pair to a Lewis acid, a compound that can accept an electron pair.[12][13] Lewis's proposal gives an explanation to the Brønsted–Lowry classification in terms of electronic structure.

In this representation both the base, B, and the conjugate base, A, are shown carrying a lone pair of electrons and the proton, which is a Lewis acid, is transferred between them.

Adduct of ammonia and boron trifluoride
Adduct of ammonia and boron trifluoride

Lewis later wrote in "To restrict the group of acids to those substances that contain hydrogen interferes as seriously with the systematic understanding of chemistry as would the restriction of the term oxidizing agent to substances containing oxygen."[13] In Lewis theory an acid, A, and a base, B:, form an adduct, AB, in which the electron pair is used to form a dative covalent bond between A and B. This is illustrated with the formation of the adduct H3N−BF3 from ammonia and boron trifluoride, a reaction that cannot occur in aqueous solution because boron trifluoride reacts violently with water in a hydrolysis reaction.

These reactions illustrate that BF3 is an acid in both Lewis and Brønsted–Lowry classifications and emphasizes the consistency between both theories.[citation needed]

Boric acid is recognized as a Lewis acid by virtue of the reaction

In this case the acid does not dissociate, it is the base, H2O that dissociates. A solution of B(OH)3 is acidic because hydrogen ions are liberated in this reaction.

There is strong evidence that dilute aqueous solutions of ammonia contain negligible amounts of the ammonium ion

and that, when dissolved in water, ammonia functions as a Lewis base.[14]

Comparison with the Lux–Flood theory

The reactions between oxides in the solid or liquid state are not included in Brønsted–Lowry theory. For example, the reaction

does not fall within the scope of the Brønsted–Lowry definition of acids and bases. On the other hand, magnesium oxide acts as a base when it reacts with an aqueous solution of an acid.

Dissolved SiO2 has been predicted to be a weak acid in the Brønsted–Lowry sense.[15]

According to the Lux–Flood theory compounds such as MgO and SiO2 in the solid state may be classified as acids or bases. For example, the mineral olivine may be regarded as a compound of a basic oxide, MgO, with an acidic oxide, silica, SiO2. This classification is important in geochemistry.


  1. ^ Brönsted, J. N. (1923). "Einige Bemerkungen über den Begriff der Säuren und Basen" [Some observations about the concept of acids and bases]. Recueil des Travaux Chimiques des Pays-Bas. 42 (8): 718–728. doi:10.1002/recl.19230420815.
  2. ^ Lowry, T. M. (1923). "The uniqueness of hydrogen". Journal of the Society of Chemical Industry. 42 (3): 43–47. doi:10.1002/jctb.5000420302.
  3. ^ Myers, Richard (2003). The Basics of Chemistry. Greenwood Publishing Group. pp. 157–161. ISBN 978-0-313-31664-7.
  4. ^ Masterton, William; Hurley, Cecile; Neth, Edward (2011). Chemistry: Principles and Reactions. Cengage Learning. p. 433. ISBN 1-133-38694-6.
  5. ^ Ebbing, Darrell; Gammon, Steven D. (2010). General Chemistry, Enhanced Edition. Cengage Learning. pp. 644–645. ISBN 0-538-49752-1.
  6. ^ Whitten, Kenneth; Davis, Raymond; Peck, Larry; Stanley, George (2013). Chemistry. Cengage Learning. p. 350. ISBN 1-133-61066-8.
  7. ^ Lew, Kristi (2009). Acids and Bases. Infobase Publishing. ISBN 9780791097830.
  8. ^ Patrick, Graham (2012). Instant Notes in Organic Chemistry. Taylor & Francis. p. 76. ISBN 978-1-135-32125-3.
  9. ^ Holliday, A.K.; Massy, A.G. (1965). Inorganic Chemistry in Non-Aqueous Solvents. Pergamon Press.
  10. ^ Reich, Hans J. "Bordwell pKa Table (Acidity in DMSO)". Department of Chemistry, University of Wisconsin, U.S. Archived from the original on 9 October 2008. Retrieved 2008-11-02.
  11. ^ Waddington, T.C. (1965). Non-Aqueous Solvent Systems. New York: Academic Press.
  12. ^ Miessler, G. L., Tarr, D. A., (1991) "Inorganic Chemistry" 2nd ed. Pearson Prentice-Hall pp. 170–172
  13. ^ a b Hall, Norris F. (March 1940). "Systems of Acids and Bases". Journal of Chemical Education. 17 (3): 124–128. Bibcode:1940JChEd..17..124H. doi:10.1021/ed017p124.
  14. ^ Housecroft, C. E.; Sharpe, A. G. (2004). Inorganic Chemistry (2nd ed.). Prentice Hall. p. 187. ISBN 978-0-13-039913-7.
  15. ^ Pauling, Linus (1960). The Nature of the Chemical Bond (3rd ed.). Ithaka: Cornell University Press. p. 557.
This page was last edited on 11 November 2019, at 04:50
Basis of this page is in Wikipedia. Text is available under the CC BY-SA 3.0 Unported License. Non-text media are available under their specified licenses. Wikipedia® is a registered trademark of the Wikimedia Foundation, Inc. WIKI 2 is an independent company and has no affiliation with Wikimedia Foundation.