In 1923, within several months of each other, Johannes Nicolaus Brønsted (Denmark) and Thomas Martin Lowry (England) published essentially the same theory about how acids and bases behave. Since they came to their conclusions independently of each other, both names have been used for the theory name.
II. The Acid Base Theory
Using the words of Brønsted:
". . . acids and bases are substances that are capable of splitting off or taking up hydrogen ions, respectively."
Or an acid-base reaction consists of the transfer of a proton from an acid to a base. KEEP THIS THOUGHT IN MIND!!
Here is a more recent way to say the same thing:
Remember: proton, hydrogen ion and H+ all mean the same thing.
Very common in the chemistry world is this definition set:
In fact, your teacher may define acids and bases this way and insist that you give those definitions back on the test. OK, go ahead and do it, but please recognize that the truth is slightly different than "donor" and 'acceptor" imply.
In an acid, the hydrogen ion is bonded to the rest of the molecule. It takes energy (sometimes a little, sometimes a lot) to break that bond. So the acid molecule does not "give" or "donate" the proton, it has it taken away. In the same sense, you do not donate your wallet to the pickpocket, you have it removed from you.
The base is a molecule with a built-in "drive" to collect protons. As soon as the base approaches the acid, it will (if it is strong enough) rip the proton off the acid molecule and add it to itself.
Now this is where all the fun stuff comes in that you get to learn. You see, some bases are stronger than others, meaning some have a large "desire" for protons, while other bases have a weaker drive. It's the same way with acids, some have very weak bonds and the proton is easy to pick off, while other acids have stronger bonds, making it harder to "get the proton."
One important contribution coming from Lowry has to do with the state of the hydrogen ion in solution. In Bronsted's announcement of the theory, he used H+. Lowry, in his paper (actually a long letter to the editor) used the H3O+ that is commonly used today. Here is what Lowry had to say:
"It is a remarkable fact that strong acidity is apparently developed only in mixtures and never in pure compounds. Even hydrogen chloride only becomes an acid when mixed with water. This can be explained by the extreme reluctance of a hydrogen nucleus to lead an isolated existence.... The effect of mixing hydrogen chloride with water is probably to provide an acceptor for the hydrogen nucleus so that the ionisation of the acid only involves the transfer of a proton from one octet to another.
HCl + H2O [an equilibrium sign] Cl¯ + OH3+
(Lowry also draws this equilibrium with all the electron "dots" to show the full octets)
The ionised acid is then really an ionised oxonium salt."
T. M. Lowry, "The Uniqueness of Hydrogen" Chemistry and Industry 42 (19 January 1923) pp43-47.
III. Sample Equations written in the Brønsted-Lowry Style
A. Reactions that proceed to a large extent:
HCl + H2O <===> H3O+ + Cl¯
HCl - this is an acid, because it has a proton available to be transfered.
H2O - this is a base, since it gets the proton that the acid lost.
Now, here comes an interesting idea:
H3O+ - this is an acid, because it can give a proton.
Cl¯ - this is a base, since it has the capacity to receive a proton.
HNO3 + H2O <===> H3O+ + NO3¯
The acids are HNO3 and H3O+ and the bases are H2O and NO3¯.
Remember that an acid-base reaction is a competition between two bases (think about it!) for a proton. If the stronger of the two acids and the stronger of the two bases are reactants (appear on the left side of the equation), the reaction is said to proceed to a large extent.
B. Reactions that proceed to a small extent:
If the weaker of the two acids and the weaker of the two bases are reactants (appear on the left side of the equation), the reaction is said to proceed to only a small extent:
HC2H3O2 + H2O <===> H3O+ + C2H3O2¯
NH3 + H2O <===> NH4+ + OH¯
Identify the conjugage acid base pairs in each reaction.
IV. Problems with Arrhenius' Theory
This theory works very nicely in all protic solvents (water, ammonia, acetic acid, etc.), but fails to explain acid base behavior in aprotic solvents such as benzene and dioxane. That job will be left for a more general theory, such as the Lewis Theory of Acids and Bases.