Substances that dissociate completely into ions when placed in water are referred to as
strong electrolytes because the high ionic concentration allows an electric current to pass through the solution. Most compounds with ionic bonds behave in this manner; sodium chloride is an example.
By contrast, other substances—like the simple sugar glucose—do not dissociate at all and exist in solution as molecules held together by strong covalent bonds. There also are substances—like sodium carbonate (Na 2CO 3)—that contain both ionic and covalent bonds. (See Figure 1.)
Figure 1. Ionic and covalent bonding in Na2CO3.
The sodium carbonate is a strong electrolyte, and each formula unit dissociates completely to form three ions when placed in water.
The carbonate anion is held intact by its internal covalent bonds.
Substances containing polar bonds of intermediate character commonly undergo only partial dissociation when placed in water; such substances are classed as weak electrolytes. An example is sulfurous acid:
A solution of sulfurous acid is dominated by molecules of H 2SO 3 with relatively scarce H 3O + and 
ions. Make sure that you grasp the difference between this case and the previous example of the strong electrolyte Na 2CO 3, which completely dissociates into ions.
Acids and bases are usefully sorted into strong and weak classes, depending on their degree of ionization in aqueous solution.
The dissociation of any acid may be written as an equilibrium reaction:
where A denotes the anion of the particular acid. The concentrations of the three solute species are related by the equilibrium equation
where K a is the acid ionization constant (or merely acid constant). Different acids have different K a values—the higher the value, the greater the degree of ionization of the acid in solution. Strong acids, therefore, have larger K a than do weak acids.
Table 1 gives acid ionization constants for several familiar acids at 25°C. The values for the strong acids are not well defined; therefore, the values are stated only in orders of magnitude. Examine the “Ions” column and see how every acid yields a hydronium ion and a complementary anion in solution.
Use the equilibrium equation and data from the preceding chart to calculate the concentrations of solutes in a 1 M solution of carbonic acid. The unknown concentrations of the three species may be written
where x represents the amount of H 2CO 3 that has dissociated to the pair of ions. Substituting these algebraic values into the equilibrium equation,
To solve the quadratic equation by approximation, assume that x is so much less than 1 (carbonic acid is weak and only slightly ionized) that the denominator 1 – x may be approximated by 1, yielding the much simpler equation
x 2 = 4.3 × 10 –7
x = 6.56 × 10 –4 = [H 3O +]
This H 3O + concentration is, as conjectured, much less than the nearly 1 molarity of the H 2CO 3, so the approximation is valid. A hydronium ion concentration of 6.56 × 10 –4 corresponds to a pH of 3.18.
You will recall from the review of organic chemistry that carboxylic acids have a single hydrogen bonded to an oxygen in the functional group. (See Figure 2.) To a very small extent, this hydrogen can dissociate in an aqueous solution. Therefore, members of this class of organic compounds are weak acids.
Carboxylic acids.
Summarize the treatment of acids so far. A strong acid is virtually completely dissociated in aqueous solution, so the H 3O + concentration is essentially identical to the concentration of the solution—for a 0.5 M solution of HCl, [H 3O +] = 0.5 M. But because weak acids are only slightly dissociated, the concentrations of the ions in such acids must be calculated using the appropriate acid constant.
- If an aqueous solution of acetic acid is to have a pH of 3, how many moles of acetic acid are needed to prepare 1 liter of the solution?