By the end of this section, you will be able to:
The process in which like molecules react to yield ions is called autoionization. Liquid water undergoes autoionization to a very slight extent; at 25 °C, approximately two out of every billion water molecules are ionized. The extent of the water autoionization process is reflected in the value of its equilibrium constant, the ion-product constant for water, Kw:
The slight ionization of pure water is reflected in the small value of the equilibrium constant; at 25 °C, Kw has a value of 1.0×10−14. The process is endothermic, and so the extent of ionization and the resulting concentrations of hydronium ion and hydroxide ion increase with temperature. For example, at 100 °C, the value for Kw is about 5.6×10−13, roughly 50 times larger than the value at 25 °C.
A solution of an acid in water has a hydronium ion concentration of 2.0×10−6 M. What is the concentration of hydroxide ion at 25 °C?
Use the value of the ion-product constant for water at 25 °C
to calculate the missing equilibrium concentration.
Rearrangement of the Kw expression shows that [OH−] is inversely proportional to [H3O+]:
Compared with pure water, a solution of acid exhibits a higher concentration of hydronium ions (due to ionization of the acid) and a proportionally lower concentration of hydroxide ions. This may be explained via Le Châtelier’s principle as a left shift in the water autoionization equilibrium resulting from the stress of increased hydronium ion concentration.
Substituting the ion concentrations into the Kw expression confirms this calculation, resulting in the expected value:
What is the hydronium ion concentration in an aqueous solution with a hydroxide ion concentration of 0.001 M at 25 °C?
[H3O+] = 1 10−11 M
Hydronium and hydroxide ions are present both in pure water and in all aqueous solutions, and their concentrations are inversely proportional as determined by the ion product of water (Kw). The concentrations of these ions in a solution are often critical determinants of the solution’s properties and the chemical behaviors of its other solutes, and specific vocabulary has been developed to describe these concentrations in relative terms. A solution is neutral if it contains equal concentrations of hydronium and hydroxide ions; acidic if it contains a greater concentration of hydronium ions than hydroxide ions; and basic if it contains a lesser concentration of hydronium ions than hydroxide ions.
A common means of expressing quantities that may span many orders of magnitude is to use a logarithmic scale. One such scale that is very popular for chemical concentrations and equilibrium constants is based on the p-function, defined as shown where “X” is the quantity of interest and “log” is the base-10 logarithm:
The pH of a solution is therefore defined as shown here, where [H3O+] is the molar concentration of hydronium ion in the solution:
Rearranging this equation to isolate the hydronium ion molarity yields the equivalent expression:
Likewise, the hydroxide ion molarity may be expressed as a p-function, or pOH:
Finally, the relation between these two ion concentration expressed as p-functions is easily derived from the Kw expression:
At 25 °C, the value of Kw is 1.0×10−14, and so:
The hydronium ion molarity in pure water (or any neutral solution) is 1.0×10−7 M at 25 °C. The pH and pOH of a neutral solution at this temperature are therefore:
And so, at this temperature, acidic solutions are those with hydronium ion molarities greater than 1.0×10−7 M and hydroxide ion molarities less than 1.0×10−7 M (corresponding to pH values less than 7.00 and pOH values greater than 7.00). Basic solutions are those with hydronium ion molarities less than 1.0×10−7 M and hydroxide ion molarities greater than 1.0×10−7 M (corresponding to pH values greater than 7.00 and pOH values less than 7.00).
Since the autoionization constant Kw is temperature dependent, these correlations between pH values and the acidic/neutral/basic adjectives will be different at temperatures other than 25 °C. For example, the hydronium molarity of pure water at 80 °C is 4.9×10−7 M, which corresponds to pH and pOH values of:
At this temperature, then, neutral solutions exhibit pH = pOH = 6.31, acidic solutions exhibit pH less than 6.31 and pOH greater than 6.31, whereas basic solutions exhibit pH greater than 6.31 and pOH less than 6.31. This distinction can be important when studying certain processes that occur at other temperatures, such as enzyme reactions in warm-blooded organisms at a temperature around 36–40 °C. Unless otherwise noted, references to pH values are presumed to be those at 25 °C (Table 22.1).
Summary of Relations for Acidic, Basic and Neutral Solutions
|Relative Ion Concentrations
|pH at 25 °C
|[H3O+] > [OH−]
|pH < 7
|[H3O+] = [OH−]
|pH = 7
|[H3O+] < [OH−]
|pH > 7
Figure 22.1 shows the relationships between [H3O+], [OH−], pH, and pOH for solutions classified as acidic, basic, and neutral.
(The use of logarithms is explained in Appendix B. When taking the log of a value, keep as many decimal places in the result as there are significant figures in the value.)
Air-saturated water has a hydronium ion concentration caused by the dissolved CO2 of 2.0×10−6 M, about 20-times larger than that of pure water. Calculate the pH of the solution at 25 °C.
(On a calculator take the antilog, or the “inverse” log, of −7.3, or calculate 10−7.3.)
Normal rainwater has a pH between 5 and 6 due to the presence of dissolved CO2 which forms carbonic acid:
Acid rain is rainwater that has a pH of less than 5, due to a variety of nonmetal oxides, including CO2, SO2, SO3, NO, and NO2 being dissolved in the water and reacting with it to form not only carbonic acid, but sulfuric acid and nitric acid. The formation and subsequent ionization of sulfuric acid are shown here:
Carbon dioxide is naturally present in the atmosphere because most organisms produce it as a waste product of metabolism. Carbon dioxide is also formed when fires release carbon stored in vegetation or fossil fuels. Sulfur trioxide in the atmosphere is naturally produced by volcanic activity, but it also originates from burning fossil fuels, which have traces of sulfur, and from the process of “roasting” ores of metal sulfides in metal-refining processes. Oxides of nitrogen are formed in internal combustion engines where the high temperatures make it possible for the nitrogen and oxygen in air to chemically combine.
Acid rain is a particular problem in industrial areas where the products of combustion and smelting are released into the air without being stripped of sulfur and nitrogen oxides. In North America and Europe until the 1980s, it was responsible for the destruction of forests and freshwater lakes, when the acidity of the rain actually killed trees, damaged soil, and made lakes uninhabitable for all but the most acid-tolerant species. Acid rain also corrodes statuary and building facades that are made of marble and limestone (Figure 22.2). Regulations limiting the amount of sulfur and nitrogen oxides that can be released into the atmosphere by industry and automobiles have reduced the severity of acid damage to both natural and manmade environments in North America and Europe. It is now a growing problem in industrial areas of China and India.
For further information on acid rain, visit this website hosted by the US Environmental Protection Agency.
The pH can be found from the pOH:
pOH = 11.6, pH = 2.4
The acidity of a solution is typically assessed experimentally by measurement of its pH. The pOH of a solution is not usually measured, as it is easily calculated from an experimentally determined pH value. The pH of a solution can be directly measured using a pH meter (Figure 23.3).
The pH of a solution may also be visually estimated using colored indicators (Figure 22.4). The acid-base equilibria that enable use of these indicator dyes for pH measurements are described in a later section of this chapter.
This content is provided to you freely by BYU Open Learning Network.
Access it online or download it at https://open.byu.edu/general_college_chemistry_2/autoionization_and_p.