Theories of Acids and Bases

Theories of Acids and Bases

Theory Arrhenius, Bronsted-Lowry, and Lewis theories of acids and bases, and explains the relationships between them. It also explains the concept of a conjugate pair - an acid and its conjugate base, or a base and its conjugate acid

• The Arrhenius Theory of Acid and Bases 

The theory, introduced in 1887 by the Swedish scientist Svante Arrhenius, that acids are substances that dissociate in water to yield electrically charged atoms or molecules, called ions, one of which is a hydrogen ion (H+), and that bases ionize in water to yield hydroxide ions (OH−). It is now known that the hydrogen ion cannot exist alone in water solution; rather, it exists in a combined state with a water molecule, as the hydronium ion (H3O+). In practice the hydronium ion is still customarily referred to as the hydrogen ion.

Acids and bases may be classified as strong or weak acids and bases depending on the hydrogen ion or hydroxide ion concentration produced in solution. The reaction between an acidand a base leads to the formation of a salt and water; the latter is the result of the combination of a hydrogen ion and a hydroxide ion.

Limitations of the theory

Hydrochloric acid is neutralised by both sodium hydroxide solution and ammonia solution. In both cases, you get a colourless solution which you can crystallise to get a white salt - either sodium chloride or ammonium chloride.

These are clearly very similar reactions. The full equations are:

In the sodium hydroxide case, hydrogen ions from the acid are reacting with hydroxide ions from the sodium hydroxide - in line with the Arrhenius theory.

However, in the ammonia case, there don't appear to be any hydroxide ions!

You can get around this by saying that the ammonia reacts with the water it is dissolved in to produce ammonium ions and hydroxide ions:

This is a reversible reaction, and in a typical dilute ammonia solution, about 99% of the ammonia remains as ammonia molecules. Nevertheless, there are hydroxide ions there, and we can squeeze this into the Arrhenius theory.

However, this same reaction also happens between ammonia gas and hydrogen chloride gas.

In this case, there aren't any hydrogen ions or hydroxide ions in solution - because there isn't any solution. The Arrhenius theory wouldn't count this as an acid-base reaction, despite the fact that it is producing the same product as when the two substances were in solution.

• The Bronsted-Lowry Theory of Acids and Bases

The Brønsted–Lowry theory is an acid–base reaction theory which was proposed independently by Johannes Nicolaus Brønstedand Thomas Martin Lowry in 1923. 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.

• Bronsted theory

acid + base  conjugate base + conjugate acid.

Conjugate pairs

When hydrogen chloride dissolves in water, almost 100% of it reacts with the water to produce hydronium ions and chloride ions. Hydrogen chloride is a strong acid, and we tend to write this as a one-way reaction:

H2O+HCl→H3O++Cl−$$H_{2}O+HCl\to H_3{O}^{+}+C{l}^{-}$$

In fact, the reaction between HCl and water is reversible, but only to a very minor extent. To generalize, consider an acid HA$HA$, and think of the reaction as being reversible.

HA+H2O⇌H3O++A−$$HA+H_{2}O\rightleftharpoons H_3{O}^{+}+A^{-}$$

Thinking about the forward reaction:

• The HA$HA$ is an acid because it is donating a proton (hydrogen ion) to the water.
• The water is a base because it is accepting a proton from the HA$HA$.

However, there is also a back reaction between the hydronium ion and the A−$A^{-}$ ion:

• The H3O+$H_3{O}^{+}$ is an acid because it is donating a proton (hydrogen ion) to the A−$A^{-}$ ion.
• The A−$A^{-}$ ion is a base because it is accepting a proton from the H3O+$H_3{O}^{+}$.

The reversible reaction contains two acids and two bases. We think of them in pairs, called conjugate pairs.

When the acid, HA$HA$, loses a proton it forms a base, A−$A^{-}$, which can accept a proton back again to refom the acid, HA$HA$. These two are aconjugate pair. Members of a conjugate pair differ from each other by the presence or absence of the transferable hydrogen ion.

• If you are thinking about HA$HA$ as the acid, then A−$A^{-}$ is its conjugate base.
• If you are thinking about A−$A^{-}$ as the base, then HA$HA$ is its conjugate acid.

The water and the hydronium ion are also a conjugate pair. Thinking of the water as a base, the hydronium ion is its conjugate acid because it has the extra hydrogen ion which it can give away again. Thinking about the hydronium ion as an acid, then water is its conjugate base. The water can accept a hydrogen ion back again to reform the hydronium ion.

• The Lewis Theory of Acids and Bases

• An acid is an electron pair acceptor.
• A base is an electron pair donor.

It is easiest to see the relationship by looking at exactly what Brønsted-Lowry bases do when they accept hydrogen ions. Three Brønsted-Lowry bases we've looked at are hydroxide ions, ammonia and water, and they are typical of all the rest.

The Brønsted-Lowry theory says that they are acting as bases because they are combining with hydrogen ions. The reason they are combining with hydrogen ions is that they have lone pairs of electrons - which is what the Lewis theory says. The two are entirely consistent. So how does this extend the concept of a base? At the moment it doesn't - it just looks at it from a different angle.

But what about other similar reactions of ammonia or water, for example? On the Lewis theory, any reaction in which the ammonia or water used their lone pairs of electrons to form a co-ordinate bond would be counted as them acting as a base. Here is a reaction which you will find talked about on the page dealing with co-ordinate bonding. Ammonia reacts with BF3 by using its lone pair to form a co-ordinate bond with the empty orbital on the boron.

As far as the ammonia is concerned, it is behaving exactly the same as when it reacts with a hydrogen ion - it is using its lone pair to form a co-ordinate bond. If you are going to describe it as a base in one case, it makes sense to describe it as one in the other case as well.

Lewis acids Lewis acids are electron pair acceptors. In the above example, the BF3 is acting as the Lewis acid by accepting the nitrogen's lone pair. On the Bronsted-Lowry theory, the BF3 has nothing remotely acidic about it. This is an extension of the term acid well beyond any common use. What about more obviously acid-base reactions - like, for example, the reaction between ammonia and hydrogen chloride gas? What exactly is accepting the lone pair of electrons on the nitrogen. Textbooks often write this as if the ammonia is donating its lone pair to a hydrogen ion - a simple proton with no electrons around it. That is misleading! You don't usually get free hydrogen ions in chemical systems. They are so reactive that they are always attached to something else. There aren't any uncombined hydrogen ions in HCl. There isn't an empty orbital anywhere on the HCl which can accept a pair of electrons. Why, then, is the HCl a Lewis acid? Chlorine is more electronegative than hydrogen, and that means that the hydrogen chloride will be a polar molecule. The electrons in the hydrogen-chlorine bond will be attracted towards the chlorine end, leaving the hydrogen slightly positive and the chlorine slightly negative.
The lone pair on the nitrogen of an ammonia molecule is attracted to the slightly positive hydrogen atom in the HCl. As it approaches it, the electrons in the hydrogen-chlorine bond are repelled still further towards the chlorine. Eventually, a co-ordinate bond is formed between the nitrogen and the hydrogen, and the chlorine breaks away as a chloride ion. This is best shown using the "curly arrow" notation commonly used in organic reaction mechanisms.
The whole HCl molecule is acting as a Lewis acid. It is accepting a pair of electrons from the ammonia, and in the process it breaks up. Lewis acids don't necessarily have to have an existing empty orbital.