Lesson Explainer: Lewis Acids and Bases Chemistry • First Year of Secondary School

In this explainer, we will learn how to explain what Lewis acids and bases are, along with their characteristic properties, and identify them in chemical reactions.

Chemists look for patterns in chemical reactions. They put chemicals in groups and compare them with one another. Two of these groups are acids and bases.

The descriptions of “acid” and “base” have changed over time, and multiple descriptions have existed at once. Today, a few of these descriptions are still in use, although some substances are considered acidic or basic in one system, but not in others.

The simplest description of an acid could be a substance that tastes sour, while the simplest description of a base could be a substance that reacts with an acid.

In 1887, Svante Arrhenius described acids and bases on the basis of whether they produce hydrogen ions () or hydroxide ions () when added to water:

Hydrogen ions and hydroxide ions react to make water. This accounts for the fact that, generally, Arrhenius bases will react with Arrhenius acids in a consistent way.

Another way Arrhenius bases are described is as “substances that when added to water raise the hydroxide ion concentration.” This would apply to substances that do not contain hydroxide ions but raise the hydroxide ion concentration of water by reacting with it, such as ammonia (). The distinctions at this level show that descriptions and definitions can vary a little.

In 1923, Johannes Brønsted and Thomas Lowry extended the description of bases, including substances that do not yield hydroxide ions in solution.

A Brønsted–Lowry base is any substance that can accept a proton. For instance, ammonia () can react with hydrogen ions () in water, forming the ammonium ion (). Generally speaking, Arrhenius acids are also Brønsted–Lowry acids, since they have a hydrogen ion they can donate. However, it is important to realize that Arrhenius acids are those chemicals that can donate hydrogen ions to water specifically.

There are also some substances that will act as acids in some circumstances and as bases in others; these are known as amphoteric substances.

Another description of acids and bases was invented by Gilbert Lewis in 1923. He described a Lewis acid as any substance that can accept a lone pair of electrons, forming a bond. He also described a Lewis base as any substance that has a lone pair of electrons that can be donated to form a bond.

In general, Brønsted–Lowry bases are Lewis bases and vice versa. However, while all Brønsted–Lowry acids are Lewis acids, not all Lewis acids are Brønsted–Lowry acids.

To see how the definitions of acids and bases can be applied, consider hydrogen chloride ().

Hydrogen chloride can be described by all three descriptions of acids.

However, now consider borane (). This is the Lewis structure of borane:

The boron atom in the middle has two vacancies in its valence shell. It will readily accept a lone pair from another substance.

This is the full table for .

The table above shows that borane can be described as a Lewis acid, but not an Arrhenius nor Brønsted–Lowry acid or base.

For sodium hydroxide (), the following table can be constructed.

We can see how behaves as both a Brønsted–Lowry base and a Lewis base in a single diagram. The oxygen of the hydroxide ion of sodium hydroxide can donate one of its lone pairs to form a bond with a hydrogen ion:

A fourth, and final, example is ammonia ().

Here, ammonia can definitely be described as a Brønsted–Lowry and Lewis base. However, its description as an Arrhenius base is not as definite and shows why other descriptions were needed.

The advantage of Lewis’s description is that the set of Lewis acids and bases is bigger than the set of substances included by Arrhenius, or Brønsted and Lowry, and Lewis acids and Lewis bases react in predictable ways.

Lewis’s description allows us to easily compare a greater number of chemical reactions. If we see an area of one molecule that is electron deficient and another area that is electron rich with a reactive lone pair, we can guess how they might interact, meaning that we can predict where bonds are likely to form. In practice, a little more information than this is needed, but it is a great starting point.

When a Lewis acid reacts with a Lewis base, they form what is called a Lewis acid–base adduct. The chemical equation below shows a Lewis base (ammonia) reacting with a Lewis acid (borane):

The lone pair from the nitrogen of ammonia can form a bond with the boron of borane. The result, , is an example of a Lewis acid–base adduct.

The adduct can be written showing the coordinate covalent bond arrow:

At the center of the molecule is a boron atom that can accept electrons, and the ammonia molecule has a reactive lone pair on the nitrogen. We can predict that the nitrogen will form a bond with the boron.

One of the useful features about Lewis acid and base theory is the strong relationship it has with Lewis structures. Lewis structures allow us to highlight lone pairs as well as electron-deficient areas, for instance, when an atom or ion does not have a full octet.

However, not all lone pairs are reactive. The reasons for this are complex, but some common Lewis bases are

• halides (, , , ),
• oxide (),
• hydroxide (),
• water (),
• cyanide (),
• cyano groups (),
• amines (),
• ammonia ().

We tend to see more reactive lone pairs on more electronegative elements.

Examples of some electron-deficient species that will act as Lewis acids are

• boranes (, ),
• aluminum trihalides (e.g., ),
• various transition metal ions.

In many cases, knowing how to draw Lewis structures and identify lone pairs and incomplete octets should make it clear where bonds would form.

It is important that we can identify Lewis acids and bases in reactions. For example, this is the reaction of trifluoroborane () and a fluoride ion ():

If we draw the Lewis structures, we can see where the lone pair that forms the bond comes from:

The fluoride ion with its reactive lone pair forms a bond with the electron-deficient boron. This can also be drawn with a coordinate covalent bond, although in reality the negative charge is distributed evenly in :

The fluoride ion is donating a lone pair to form a bond—it is acting as a Lewis base.

The trifluoroborane is accepting a lone pair, forming a bond—it is acting as a Lewis acid.

There are some substances that are Lewis acids under some circumstances and Lewis bases under others. These are known as Lewis amphoteric substances.

When hydrogen chloride dissolves in water, a water molecule reacts with the hydrogen ion in , forming the hydronium ion and the chloride ion. We often treat the hydronium ion () like . However, whenever we see , we are really dealing with a Lewis acid–base adduct. One of the lone pairs on the oxygen in water is donated to form a bond with . In this example, water is acting as a Lewis base, donating a lone pair:

We see a different behavior from water in the formation of ammonium hydroxide when ammonia is added to water. In this case, it is the lone pair on the nitrogen that forms a bond with the ion. What we form is the ammonium ion () and the hydroxide ion ():

In this example, the water is the Lewis acid as one of the hydrogens accepts an electron pair from ammonia. Water is therefore amphoteric; it can be a Lewis acid or, in other circumstances, a Lewis base.