In this tutorial, you will compare and contrast different acid base theories. You will be introduced to the Arrhenius and Bronsted-Lowry theories, and the idea of an Arrhenius acid and a Bronsted-Lowry acid (Bronsted acid).
Make sure to check out the Lewis Acids & Bases tutorial to learn about the Lewis theory!
Topics Covered in Other Articles
- Acid-Base Chemistry
- Definition of pH
- Strong Acids & Bases
- Properties of Acids & Bases
- Net Ionic Equations
- Common Ion Effect
What are Acid Base Theories?
There are 3 different acid base theories in chemistry; each of these theories have their own definition of what qualifies as an acid- base reaction. Let’s learn the different qualifications of each theory below!
Arrhenius Acids and Bases
This theory defines both the acid and base as a compound that is able to dissociate in water, meaning it breaks up into its ion counterparts (i.e. HBr → H+ + Br–).
An Arrhenius acid is a compound that yields H+ ions in solution. Nitric acid, HNO3, is an example of an Arrhenius acid because it will dissociate and create H+ and NO3– ions. The hydrogen ions then combine with water to form H3O+; this is because the oxygen’s lone pair attracts the hydrogen, which is too weak to exist alone in solution.
H-A(aq) → H+(aq) + A–(aq)
H-A (aq) + H2O (l) → H3O+(aq) + A–(aq)
On the other hand, an Arrhenius base is a compound that yields OH– ions in solution. Potassium hydroxide, KOH, is an Arrhenius base because it will dissociate and create K+ and OH– ions.
B-OH (aq) → B+(aq) + OH–(aq)
Furthermore, neutralization occurs within the solution due to the H+ and OH– ions also reacting together to form water.
H+(aq) + OH–(aq) → H2O(l)
The Arrhenius theory has the most specificity out of the three theories because it states that either the H+ or OH– ion must be present to be labeled an acid or base; however, this is not true in other theories. Let’s move on to a more general theory that includes more compounds – the concept of a Bronsted-Lowry acid.
Bronsted-Lowry Acids and Bases
This definition of acids and bases can be used to analyze solutions both containing and not containing water; it has to do with how well a compound can accept or donate protons.
A Bronsted-Lowry acid, aka Bronsted Acid, is a proton donor, meaning it can release a proton; when an acid dissociates in solution, it increases the H+ yield. Again, nitric acid, HNO3, can be labeled as an acid because it will dissociate and create H+ and NO3– ions.
H-A(aq) → H+(aq) + A–(aq)
H-A(aq) + H2O(l) → H3O+(aq) + A–(aq)
A Bronsted-Lowry base, aka Bronsted base, is a proton acceptor; when a base dissociates, it takes a proton from water to generate OH– ions in solution. So is NH3 an acid or base? Ammonia, NH3, is an example of a base because it takes a proton from water to form the products of NH4+ and OH–.
B(aq) + H2O(l) → B-H+(aq) + OH–(aq)
In this theory, water can be either an acid or a base; this is because it can accept a proton to form H3O+ or donate a proton to form OH–; when a compound can act as both a Bronsted-Lowry acid or base, it is said to be amphoteric.
The Bronsted-Lowry theory adds onto the Arrhenius theory. The general idea of yielding H+ and OH– ions is the same, but this theory is more general, which allows more compounds to be labeled as acids or bases; for example, bases do not have to contain OH– (they have to in the Arrhenius theory) because taking a proton from water forms the hydroxide ion.
Lewis Acids and Base Theory
This theory is has the least specificity, meaning it allows for a broad selection of acids and bases. A Lewis acid is an electron pair acceptor, while a Lewis base is an electron pair donor. There is no mention of protons, or hydrogens, which creates a chance for more compounds to fit this category. The Lewis theory states that a lewis base donates a pair of electrons that is either shared with or used by a lewis acid. Read our Lewis Acids & Bases tutorial to go into more detail!