Lesson 5: Lewis Concept & Comparing the Concepts

The Lewis Concept | Comparing the Concepts

The Lewis Concept

The third conceptual definition of acids and bases is that of Lewis. You've heard of him before. Electron dot diagrams are sometimes called Lewis diagrams. Lewis focused his thoughts about acids and bases on what happens when a base accepts a proton and what structural features of the base allow it to accept a proton. Let's look at what happens. (This is also shown in Ex. 17 in your workbook.) {Note: The charges shown below apply to the entire polyatomic ion.}

When a hydroxide ion accepts a proton, you can see that the hydrogen ion, the proton, attaches itself to a pair of electrons from the oxygen.	··  _ H : O :     ·· + H+ ·· H : O : H ··
When a hydrogen ion reacts with ammonia, it attaches itself to a pair of electrons on the nitrogen.	H ·· H : N :    ·· H + H+ H       ··   + H : N : H ·· H
When a hydrogen ion reacts with carbonate ion it attaches itself to a pair of electrons on one of the oxygen atoms that surround the carbon in the carbonate ion.	··          ··  2- : O : C : O : ··   : :   ·· : O : + H+   ··          ··    - : O : C : O : H ··   : :   ··   : O :

 

What happens in each case is that the base provides or donates a pair of electrons to the hydrogen ion in order to allow it to form a covalent bond. So, Lewis' definition of a base is that a base is an electron pair donor.

Now, think about what the hydrogen ion is doing in each case. The hydrogen ion is attaching to or accepting an electron pair. So, just as a base is an electron pair donor, an acid is an electron pair acceptor.

Once you say that an acid is an electron pair acceptor, then you're not limited to the idea that an acid has to be a hydrogen ion or some chemical that contains or releases hydrogen ions. It could be something else, as long as it can accept (or bond to) a pair of electrons.

One commonly cited example of a Lewis acid that does not contain hydrogen is boron trifluoride, shown here accepting a pair of electrons from ammonia. There are many other examples of Lewis acids. I won't go into them because we won't be dealing with them very much.

··   : F : ··   ··    : F : B     ··   ··      : F :   ·· + H   ··     : N : H ··   H  ··      : F :   H ··   ··     ··    : F : B  : N : H ··   ··    ··    : F :   H ··      acid base

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Comparing the Concepts

The three conceptual definitions can be compared with one another by looking at how each one views what happens when an acid dissociates in water. Let's use hydrochloric acid (HCl) as an example. (These equations are also shown in example 18 in your workbook.)

The first equation shows that hydrochloric acid dissociates to form hydrogen ion and chloride ion. This is the Arrhenius view.	HCl  H+ + Cl¯
The second equation shows a closer approximation to what is really going on. The HCl doesn't just dissociate, but reacts with water to form hydronium ion, H3O+. The idea here is that a proton is transferred from one chemical to another in an acid-base reaction. This is the Brønsted-Lowry point of view.	HCl + H2O  H3O+ + Cl¯
The third equation shows the Lewis way to look at the formation of the hydronium ion. A hydrogen ion from HCl combines with a water molecule to form H3O+. The reaction occurs because H2O provides an electron pair for the hydrogen ion to accept. The electron dots emphasize the Lewis point of view.	H+     + ·· H : O : H ··     ··  + H : O : H ·· H

 

We will deal primarily with acids and bases from the Brønsted and Lowry point of view. At times we will also use the Arrhenius definition of acids and bases. Actually, what we will do is use whichever definition is most appropriate to the particular aspect of acids and bases that we are trying to study at the time.

 

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