Acid Base 2

Introduction

As you are aware, all acids and bases are not created equal. Some are "more acidic" or "more basic" than others. But what do we mean when we say that? What factors determine an acid or base's "strength". We will discuss these questions in this note pack.

Strength: what is it?

To define and discuss the strength of acids or bases, it's useful to use the Arrhenius or Bronsted-Lowry definitions and to remember what strong electrolytes are. Strong electrolytes (from chapter 4 in CHM 110) are substances that dissociate completely in water.

Acids:

Strong acids, then, are substances that dissociate completely in water to form hydronium ions (or, substances that readily donate all of their acidic protons to another substance).
Weak acids are substances that partially dissociate in water to form hydronium ions (or, substances that do not readily donate all of their acidic protons to another substance).

Bases:

Strong bases are substances that dissociate completely in water to form hydroxide ions (or, substances that readily accept all available protons from another substance).
Weak bases are substances that partially dissociate in water to form hydroxide ions (or, substances that readily accept all available protons from another substance).

You should memorize common strong acids and bases if you haven't already.

Common Strong Acids	Common Strong Bases
HCl, hydrochloric acid	Alkali metal hydroxides like NaOH
HNO₃, nitric acid	Ca(OH)₂
H₂SO₄, sulfuric acid	Ba(OH)₂
HI	Sr(OH)₂
HBr

What makes strong acids strong?

For Bronsted-Lowry acids, it's relatively easy to explain acid strengths. Recall that a Bronsted-Lowry acid loses a proton. The acid's strength depends on how easy it is to lose that proton.

The strength of the acid depends on the strength of the bond holding the acidic proton onto the molecule. For binary acids (simple acids that have hydrogen attached to one other element like HF, HCl, etc.), the acid strength is easy to determine - look at the bond energy of the only bond in the molecule. The higher the bond energy, the weaker the acid. Comparing HF and HI, for example: HF has the highest bond energy, and is a weak acid. HI has the lowest bond energy and is a strong acid.

You can make another interesting observation with polyprotic acids like sulfuric acid, H₂SO₄. (Polyprotic acids can lose more than one hydrogen ion from each acid molecule.)

Let's look at what happens when you put sulfuric acid in water:

H₂SO₄(aq) + H₂O(l) <--> H₃O⁺(aq) + HSO₄^-(aq)

HSO₄^-(aq) + H₂O(l) <--> H₃O⁺(aq) + SO₄^2-(aq)

As you can see, sulfuric acid loses one proton, then another. This reaction has been thoroughly studied and it's been found that the loss of the first proton is complete (in other words, sulfuric acid is a strong acid). However, the loss of the second proton doesn't quite proceed to completion. In other words, HSO₄^- is a weak acid! Why would this be the case?

Look at the hydrogen sulfate (HSO₄^-) ion and remember that we're trying to pull a proton (H⁺) off of it. It should be more difficult to pull a positively charged proton off of a substance that is already negatively charged. Remember that opposite charges attract. This agrees with the experimental data - it is harder to pull off the second proton from sulfuric acid. This is true for other polyprotic acids like H₃PO₄ (with three acidic protons). It's much easier to pull off the first proton than the second. It's also much easier to pull off the second proton than the third.

Summary

In this note pack, we've discussed strong and weak acids and bases. You should understand what is meant by "strong" and "weak". If you haven't memorized them already, make sure you know the common strong acids and bases (making a set of flash cards might help here). You should also understand what makes a strong acid strong and a weak acid weak.