Clackamas Community College

CH 105: INTRODUCTORY CHEMISTRY

 

 

Contact instructor:

Eden Francis

Physical Science
19600 Molalla Avenue
Oregon City, OR 97045
(503) 594-3352
TDD (503) 650-6649

Lesson 7: Acid & Base Strength

In this section we'll review two of the definitions of acids and bases and discuss acid and base strength in more depth.

Definitions of Acids & Bases | Strong vs. Weak Acids | Strong vs. Weak Bases

 

Definitions of Acids & Bases (a review)

There are several useful definitions of acids and bases. The simplest is the Arrhenius definition, which says that an acid is any chemical that produces hydronium ions in water, while a base is any chemical that produces hydroxide ions in water.

The hydronium ion, H3O+ (aq), is sometimes written H+(aq), or just H+, for the sake of brevity.

The two equations on your screen involving HCl are just two ways of describing the same reaction. The H+ ion – also called a proton – is always attached to some other molecule or ion. In aqueous solutions, it attaches itself to a water molecule to form the hydronium ion, H3O+. We will usually just write H+.

 

A more useful definition was devised by Brønsted and Lowry, who said that an acid was anything that could donate a proton (an H+ ion), while a base was anything that could accept a proton.

By this definition, the dissociation of a metal hydroxide is not an acid-base reaction, it is simply the dissociation of an ionic compound in water.

NaOH Na+ + OH-

The resulting solution is basic, however, because the hydroxide ion, OH-, is capable of accepting a proton. That is, by the Brønsted-Lowry definition, it is a base.

 

Here is the general equation for an acid-base reaction. Notice, in particular, the double arrows instead of the equal sign separating the reactants from the products. The double arrows simply imply that the reaction can also go backwards.

In this generalized reaction, the A- can be a single ion, like Cl- or Br-, or it could also represent a polyatomic ion like ClO3- or NO3-. The key is that once the H+ ion has been donated, the species that is left has a charge one less than the original acid. Thus, if HA represents a charged species, like HSO4-, then A- actually represents the sulfate ion, SO42-:

HSO4- + H2O SO42- + H3O+

 

In the forward reaction, HA donates a proton to H2O. HA is the acid (the proton donor) and H2O is the base (the proton acceptor).

In the reverse reaction, H3O+ donates the proton back to the A-. H3O+ is therefore the conjugate acid and A- is the conjugate base.

A- is called the conjugate base of HA because it is the base that can accept the proton back once HA (acting as an acid) has donated it. Similarly, H3O+ is called the conjugate acid of H2O because once H2O, acting as a base, has accepted the proton, H3O+ can donate it back, thus playing the part of an acid.

 

Top of page

 

Strong vs. Weak Acids

You can look at this reaction as a competition between A- and H2O for the proton. If H2O has a stronger attraction for the proton than A-, then most of the protons will stick to the water, forming H3O+ ions. If this is the case, we say that HA is a strong acid – it has a strong tendency to donate its proton to water.

When a stong acid dissolves in water, most of the acid molecules donate their protons to water molecules so that the concentration of hydronium ions is high and the concentration of HA is low.

The stronger the acid, the higher the hydronium ion concentration and the lower the HA concentration. But even for the strongest acids, there will always be a few HA molecules in solution.

 

If, on the other hand, A- has a stronger attraction for the proton than the H2O does, then most of the protons will remain stuck to the A-. In this case, we say that HA is a weak acid, it has a weak tendency to donate its proton to water.

 

When a weak acid dissolves in water, only a few of the HA molecules donate their protons to water molecules, so the concentration of hydronium ions is low and the concentration of HA is high.

The weaker the acid, the lower the hydronium ion concentration. In fact, one way to measure the strength of an acid is to measure the concentration of hydronium ions when it dissolves in water.

 

When an acid dissolves in water, protons transfer back and forth between the A- and the H2O until a balance is achieved. This balance is called dynamic equilibrium.

 

When HA is first added to water, there is virtually no H3O+ present.

As protons are transferred to water, the concentration of H3O+ slowly increases while the concentration of HA falls.

As a result, the rate at which hydronium ions donate their protons back to the A- slowly increases while the rate at which HA transfers its protons to water slowly decreases.

Eventually, the rate at which HA molecules are donating protons to water is the same as the rate at which hydronium ions are donating their protons back to the A-.

At this point, the reaction is said to have reached equilibrium and the concentration of the various species remain constant.

Both processes are still occurring, however, albeit at equal rates, so the equilibrium is said to be dynamic.

 

For a strong acid, we say the equilibrium lies to the right, that is, it favors formation of H3O+.

One definition of a strong acid is that it is one whose tendency to donate protons to water is greater than the tendency of the H3O+ ion to donate protons back to the A- ion. In other words, at equilibrium, the concentration of H3O+ is greater than the concentration of HA.

The only strong acids we will encounter are HCl, HBr, HI, HNO3, H2SO4, and HClO4. All other acids are weak acids. You should memorize this list.

For a weak acid, we say the equilibrium lies to the left, that is, it favors formation of HA.

There is, unfortunately, some ambiguity in the use of the terms “strong” and “weak”. A solution is weakly acidic if it has only a few H3O+ ions. This might be due to the fact that it is a solution of a weak acid and only a small percentage of the protons are transferred to water, or it might also be due to the fact that, although the solution contains a strong acid, there is very little of the acid there – in other words, the concentration of the acid is low. Take care to note whether what is being described is a weak acid or a weakly acidic solution.

Reactions like the acid-base reactions we have been looking at that can go in both directions are said to be reversible.

 

There are some reactions that have no tendency whatsoever to proceed in the reverse direction. An example would be the combustion of butane

2 C4H10 + 13 O2 8 CO2 + 10 H2O

Such reactions are written with an arrow that points only one way. They are referred to as irreversible reactions.

If HA is a strong acid, then the A- ion has a very weak attraction for the proton, which is therefore easily transferred to a water molecule.

 

Also recall that the A-, since it can accept the proton back, can function as a base and is called the conjugate base of the acid HA. The Brønsted-Lowry definition of a base is that it is a proton acceptor.

It may seem paradoxical that what characteristics a “strong” acid is its weak hold on the proton. If you don’t mind anthropomorphizing a bit, think of it as a “strong generosity” on the part of the acid to donate its proton and the weakness of the conjugate base as its “reluctance” to take back its donation.

Since the A- ion has a very weak tendency to accept a proton, it is, by the Brønsted-Lowry definition, a very weak base. The same property that makes HA a strong acid makes A- a very weak base and therefore, in general, stong acids have very weak conjugate bases.

 

By the same token, if A- has a strong attraction fro the proton, it is by definition a strong base, and HA will have a very weak tendency to donate the proton, making it a very weak acid. Very weak acids have strong conjugate bases.

 

Between these two extremes, moderately weak acids have moderately weak conjugate bases and vice-versa. Usually we leave out the “moderately” and just say that weak acids have weak conjugate bases and weak bases have weak conjugate acids.

 

Top of page

 

Strong vs. Weak Bases

Metal hydroxides such as NaOH and Mg(OH)2 are considered bases because they contain hydroxide ions, OH-. In the solid, however, the hydroxide ion is bound to a positive metal ion and has no tendency to accept a proton. Therefore, the solids themselves are not basic according to the Brønstead-Lowry definition.

When metal hydroxides dissolve in water, however, the hydroxide ion separates from the metal ion. It is the free hydroxide ion that is strongly basic - it has a strong tendency to accept a proton.

If OH- is a strong base, its conjugate acid must be a very weak acid. This is fortunate, as the conjugate acid of OH- is H2O water.

The neutral ammonia molecule, NH3, on the other hand, does have a significant ability to attract protons, and therefore is a base according to the Brønsted-Lowry definition.

 

Metal hydroxides that dissolve easily in water release many hydroxide ions and so are strong bases, while insoluble or slightly soluble metal hydroxides do not release many OH- ions into solution and are considered weak bases.

NaOH and KOH are the only two strong metal hydroxide bases we will encounter. All other metal hydroxides are insoluble or only slightly soluble in water, and so are considered weak bases. Note that this definition of base strength has nothing do do with the tendency to accept a proton – the Brønsted-Lowry concept. Rather, it is based on Arrhenius’s idea that bases from hydroxide ions in water. Thus, strong bases are those that can dissolve readily and release many hydroxide ions into the solution, while weak bases are those that release only a few hydroxide ions into solution. This way of categorizing metal hydroxide bases focuses on the basicity of the solution they form, not on the intrinsic basicity of their hydroxide ions.

 

In the next page of this lesson, we'll look at how the strengths of acids and bases can be expressed using acid (or base) equilibrium expressions.

 

 

Top of page