Pure water always contains a small number of molecules that have dissociated into hydrogen ions (H+) and hydroxyl ions (OH), as illustrated by Equation 3.1.
The water dissociation constant, Kw is defined as the product of the concentrations of H+ and OH- ions, expressed in moles per liter:
where enclosing a species in square brackets is chemical symbolism that represents the species concentration in moles per liter.
Because the degree of dissociation increases with temperature, Kw is temperature dependent. At 25° C,
If, for example, an acid is added to water at 25° C, the H+ concentration increases but the product expressed by Equation 3.3 will always be equal to 1.0 x 10-14 (mol/L)2. This means that if [H+] increases, [OH-] must decrease. Adding a base causes [OH-] to increase and [H+] to decrease correspondingly.
In pure water or in water with no other sources or sinks of H+ or OH-, Equation 3.1 leads to equal numbers of H+ and OH- species. Thus, at 25°C, the values of [H+] and [OH-] must each be equal to 1.0 x 10-7 mol/L, since:
Kw,25C = (1.0 x 10-7 mol/L)(1.0 x 10-7 mol/L) = 1.0 x 10-14 (mol/L)2.
Pure water is neither acidic nor basic. Pure water defines the condition of acid-base neutrality. Therefore, acid-base neutral water always has equal concentrations of H+ and OH-, or [H+] = [OH-].
In neutral water at 25° C, [H+] = [OH] = 1 x 10-7 mol/L.
In neutral water at 50° C, [H+] = [OH] = 4.3 x 10-7 mol/L.
Whatever their separate values, the product of hydrogen ion and hydroxyl ion concentrations must be equal to 1 x 10-14 at 25° C, as in Equation 3.3. If for example [H+] = 10-5 mol/L, then it is necessary that [OH] = 10-9 mol/L, so that their product is 10-14 (mol/L)2.
Many compounds dissociate in water to form ions. Those that form hydrogen ions, H+, are called acids because when added to pure water they cause the condition [H+] > [OH]. Compounds that cause the condition [H+] < [OH] when added to pure water are called bases. An acid water solution gets its acidic properties from the presence of H+. Because H+ is too reactive to exist alone, it is always attached to another molecular species. In water solutions, H+ is often written as H3O+ because of the almost instantaneous reaction that attaches it to a water molecule
H3O+ is called the hydronium ion. It does not make any difference to the meaning of a chemical equation whether the presence of an acid is indicated by H+ or H3O+. For example, the addition of nitric acid, HNO3, to water produces the ionic dissociation reaction
Both equations are read "HNO3 added to water forms H+ (or H3O+) and NO3- ions." Defining pH
The concentration of H+ in water solutions commonly ranges from about 1 mol/L (equivalent to 1 g/L or 1000 ppm) for very acidic water, to about 10-14 mol/L (10-14 g/L or 10-11 ppm) for very basic water. Under special circumstances, the range can be even wider.
Rather than work with such a wide numerical range for a measurement that is so common, chemists have developed a way to use logarithmic units for expressing [H+] as a positive decimal number whose value normally lies between 0 and 14. This number is called the pH, and is defined in Equation 3.6 as the negative of the base10 logarithm of the hydrogen ion concentration in moles per liter:
Note that if [H+] = 10-7, then pH = -log10(10-7) = - (-7) = 7. A higher concentration of H+ such as [H+] = 10-5 yields a lower value for pH, i.e., pH = -log10(10-5) = 5. Thus, if pH is less than 7, the solution contains more H+ than OH- and is acidic; if pH is greater than 7, the solution is basic.
In acid-base reactions, protons (H+ ions) are transferred between chemical species, one of which is an acid and the other is a base. The proton donor is the acid and the proton acceptor is the base. For example, if an acid, such as hydrochloric acid (HCl), is dissolved in water, water acts as a base by accepting the proton donated by HCl. The acid-base reaction is written: HCl + H2O —> Cl- + H3O+. A water molecule that behaved as a base by accepting a proton is turned into an acid, H3O+, a species that has a proton available to donate. The species H3O+, as noted above, is called a hydronium ion and is the chemical species that gives acid water solutions their acidic characteristics. An HCl/water solution contains water molecules, hydronium ions, hydroxyl ions (in smaller concentration than H3O+), and chloride ions. The solution is termed acidic, with pH (at 25° C) < 7. The measurable parameter pH indicates the concentration of protons available for acid-base reactions.
1. In an acid-base reaction, H+ ions are exchanged between chemical species. The species that donates the H+ is the acid. The species that accepts the H+ is the base.
2. The concentration of H+ in water solutions is an indication of how many hydrogen ions are available, at the time of measurement, for exchange between chemical species. The exchange of hydrogen ions changes the chemical properties of the species between which the exchange occurs.
3. pH is a measure of [H+], the hydrogen ion concentration, which determines the acidic or basic quality of water solutions. At 25° C:
The [H+] of water in a stream = 3.5 x 10-6 mol/L. What is the pH? Answer:
pH = -logw[H+] = -logw(3.5 x 10-6) = - (-5.46) = 5.46.
Notice that since logarithms are dimensionless, the pH unit has no dimensions or units. Frequently, pH is unnecessarily assigned units called SU, or standard units, even though pH is unitless. This mainly serves to avoid blank spaces in a table that contains a column for units, or to satisfy a database that requires an entry in a units field. An alternate and useful form of Equation 3.6 is:
The pH of water in a stream is 6.65. What is the hydrogen ion concentration? Answer:
Importance of pH
Measurement of pH is one of the most important and frequently used tests in water chemistry. pH is an important factor in determining the chemical and biological properties of water. It affects the chemical forms and environmental impact of many chemical substances in water. For example, many metals dissolve as ions at lower pH values precipitate as hydroxides and oxides at higher pH and redissolve again at very high pH. Figure 3.1 shows the pH scale and typical pH values of some common substances.
pH also influences the degree of ionization, volatility, and toxicity to aquatic life of certain dissolved substances, such as ammonia, hydrogen sulfide, and hydrogen cyanide. The ionized form of ammonia, which predominates at low pH, is the less toxic ammonium ion NH4+. NH4+ transforms to the more toxic form of unionized ammonia NH3, at higher pH. Both hydrogen sulfide (H2S) and hydrogen cyanide (HCN) behave oppositely to ammonia; the less toxic ionized forms, S2- and CN-, are predominant at high pH, and the more toxic unionized forms, H2S and HCN, are predominant at low pH. The pH value is an indicator of the chemical state in which these compounds will be found and must be considered when establishing water quality standards.
The pH of environmental waters is most commonly measured with electronic pH meters or by wetting with sample, special papers impregnated with color-changing dyes. Battery-operated field meters are common. A pH measurement of surface or groundwater is valid only when made in the field or very shortly after sampling. The pH is altered by many processes that occur after the sample is collected, such as loss or gain of dissolved carbon dioxide or the oxidation of dissolved iron. A laboratory determination of pH made hours or days after sampling may be more than a full pH unit (a factor of 10 in H+ concentration) different from the value at the time of sampling.
Loss or gain of dissolved carbon dioxide (CO2) is one of the most common causes for pH changes. When CO2 dissolves into water, by diffusion from the atmosphere or from microbial activity in water or soil, the pH is lowered. Conversely, when CO2 is lost, by diffusion to the atmosphere or consumption during photosynthesis of algae or water plants, the pH is raised.
Rules of Thumb
1. Under low pH conditions (acidic)
a. Metals tend to dissolve.
b. Cyanide and sulfide are more toxic to fish.
c. Ammonia is less toxic to fish.
2. Under high pH conditions (basic)
a. Metals tend to precipitate as hydroxides and oxides. However, if the pH gets too high, some precipitates begin to dissolve again because soluble hydroxide complexes are formed (see Metals).
b. Cyanide and sulfide are less toxic to fish.
c. Ammonia is more toxic to fish.
The pH of pure water at 25° C is 7.0, but the pH of environmental waters is affected by dissolved carbon dioxide and exposure to minerals. Most unpolluted groundwaters and surface waters in the U.S. have pH values between about 6.0 and 8.5, although higher and lower values can occur because of special conditions such as sulfide oxidation which lowers the pH, or low carbon dioxide concentrations which raises the pH. During daylight, photosynthesis in surface waters by aquatic organisms may consume more carbon dioxide than is dissolved from the atmosphere, causing pH to rise. At night, after photosynthesis has ceased, carbon dioxide from the atmosphere continues to dissolve and lowers the pH again. In this manner, photosynthesis can cause diurnal pH fluctuations, the magnitude of which depends on the alkalinity buffering capacity of the water. In poorly buffered lakes or rivers, the daytime pH may reach 9.0 to 12.0.
The permissible pH range for fish depends on factors such as dissolved oxygen, temperature, and concentrations of dissolved anions and cations. A pH range of 6.5 to 9.0, with no short-term change greater than 0.5 units beyond the normal seasonal maximum or minimum, is deemed protective of freshwater aquatic life and considered harmless to fish. In irrigation waters, the pH should not fall outside a range of 4.5 to 9.0 to protect plants.
Domestic water supplies: 5.0-9.0. Freshwater aquatic life: 6.5-9.0.
Rules of Thumb
1. The pH of natural unpolluted river water is generally between 6.5 and 8.5.
2. The pH of natural unpolluted groundwater is generally between 6.0 and 8.5.
3. Clean rainwater has a pH of about 5.7 because of dissolved CO2.
4. After reaching the surface of the earth, rainwater usually acquires alkalinity while moving over and through the earth, which may raise the pH and buffer the water against severe pH changes.
5. The pH of drinking water supplies should be between 5.0 to 9.0.
6. Fish acclimate to ambient pH conditions. For aquatic life, pH should be between 6.5 to 9.0 and should not vary more than 0.5 units beyond the normal seasonal maximum or minimum.
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