Acid Titration Curves

	Hydrochloric acid,1M
	Hydrochloric acid, 0.005M
	Acetic acid,1M (40,000 ppm)
	Acetic acid,0.005M (200 ppm)
	Hypochlorous acid, 1 ppm as Cl
	Carbonic acid,0.01M (500 ppm as CaCO3)

The first two graphs show the behavior of hydrochloric acid, a strong acid, at two concentrations: 1 molar and 0.005 molar. The pH's of the initial solutions are equal to the negative logarithms of the acid concentration-- 0 (zero) and 2.3, respectively. After 90% of the acid has been neutralized ( f = 0.9), the acid concentration is one tenth of the its original value, so the pH has gone up by one unit-- to 1.0 and 3.3, respectively. Similarly, after 99% has been neutralized, the pH is 2 units higher than it was at the start. After this, the pH rapidly approaches neutrality. As soon as more base has been added than is needed to neutralize all of the acid ( f > 1), the solution would rapidly become alkaline, in sort of a mirror image of the previous part of the curve. With a strong acid /strong base system like this the pH will be very low or very high most of the time. It is hard to control the pH at an intermediate value because a small addition of acid or base causes large changes in pH near the equivalence point-- that is, at f values near 1.

The next two curves show the behavior of a weak acid, acetic acid at the same two concentrations. Aside from the pH near the pure acid (f = 0) range, the curves are not very different from one another. The more concentrated acid solution has a lower pH. As strong base is added, the hydroxide ion reacts with acetic acid to form water and acetate ion, which is a weak base. Over most of the range, the pH is determined by the ratio of the concentration of undissociated acetic acid to the concentration of acetate ion. At the f = 0.5 point, where the acetic acid and acetate concentrations are equal, the pH is equal to the pK value of the acid, 4.7.
On the other hand, the partially neutralized 1M acid will have a much greater " buffering" ability than the 0.005M acid. For example, with the 0.005M solution, if you have a liter of an acetic acid/acetate mixture with a pH of 5.0 (f = 0.7) and you add 0.002 moles of a strong acid, you will be moving the system to the point where f = 0.3)-- and the pH will be lowered to about 4.3 or 4.4. (See graph.) But with a 1M acetic acid/acetate buffer at pH7, adding the same amount of strong acid will only lower the f from 0.7 to 0.698, and the pH change will be almost undetectable.

The next curve illustrates hypchlorous acid, HOCl, the weak acid which forms -- along with hydrochloric acid-- when chlorine gas is dissolved in water. The curve is for a concentration equivalent to 1 mg/L chlorine. From the curve you can see that at a pH of 7, only 20% of the chlorine is in the basic form of hypochlorite ion, OCl^- (f = 0.2), meaning that 80% is in the form of HOCl. At pH8, the situation is reversed. This has importance in water and wastewater treatment, as HOCl is a more potent disinfectant than OCl^-.

The last example is of a dibasic acid, carbonic acid, H₂CO₃. This is the acid which forms when carbon dioxide gas, CO₂, dissolves in and reacts with water, H₂O. As strong base is added, it reacts with the H₂CO₃ to form water and HCO₃^-, which is called bicarbonate ion. The pK of carbonic acid is 6.3 (making it a weaker acid than acetic acid). So a pH of 6.3 represents the middle of the first "buffer range" of this acid.
If we continue to add strong base after all of the carbonic acid has been converted to bicarbonate ion, (at f = 1), the HCO₃^- reacts with the hydroxide to form water and carbonate ion, CO₃⁼. So bicarbonate ion is both a weak base and a weak acid. As an acid, it has a pK of 10.3, so this is the middle of the second buffer range. The pH of a pure carbonate solution, such as could be made by dissolving powdered sodium carbonate in water, would depend on the concentration: the higher the concentration, the higher the pH (just as a higher concentration of pure carbonic acid would result in a lower pH.) A pure bicarbonate solution has a pH of about 8.3, and is less affected by the concentration.

Note: These graphs were developed in Lotus 1-2-3 Release 5 for Windows^© using equilibrium equations and constants which can be found in the textbook, "Aquatic Chemistry" by Werner Stumm and James J. Morgan, Wiley-Interscience, 1970. The graphs were copied to the clipboard and into a graphics program, then saved in gif format.

Back to "More about pH and Alkalinity"

This page hosted by Get your own Free Home Page

Last update, Sept. 17, 1998

This page accessed times since Sept. 17, 1998.