ChemBytes - Chapter 8
February 23, 1998
Acids and Bases
CHEMByte 15: Acids, Bases, and the Hydrogen Ion
The first theory of acids, or acidic behavior, that is at all comprehensible in modern usage is attributable to Lavoisier (1787). Late in the eighteenth century, he erroneously identified oxygen as the "acidifying principle" which converted elements such as carbon, nitrogen and sulfur into acids such as carbonic, nitric and sulfuric acids. Unfortunately, his assumption that all acids contained oxygen led to the view that hydrochloric acid... and therefore chlorine... contained oxygen. The British chemist, Sir Humphrey Davy (1811), questioned that view, and the eventual discoveries of hydrobromic, hydriodic, and hydrocyanic acids, all of which are notable for their lack of oxygen, caused Lavoisier's definition to be set aside. What was observed, however, was the presence of hydrogen in all known acids. Still, established theories often die slowly and this one was no exception; some notable chemists of the day such as Berzelius and Gay-Lussac backed the earlier ideas on acids well into the nineteenth century. And although the notion that oxygen must be present in order to observe acidic behavior is now long dead, we are easily reminded of Lavoisier's views in the very word "oxygen" which is derived through the French oxygéne from the Greek for sour, a characteristic property of acids.
In 1834, Michael Faraday, a towering genius of 19th century British science, showed acids and bases to be substances he called "electrolytes," setting the stage for modern understanding of acidic substances. With the theory of dissociation of electrolytes into ion, as presented by Wilhelm Ostwald and Arrhenius at the end of the nineteenth century, the notion of acidity could be linked to hydrogen ions in aqueous solutions and chemical equilibrium. Now acidity could be treated quantitatively for the first time. Similarly, the constitution of bases could be understood to be associated with hydroxide ions in solution and the reaction of acids and bases together could be explained in terms of their neutralization, namely,
H+ + OH- 
H2O
But it was Brönsted who generalized these late nineteenth century ideas, in ways that removed most of the difficulties and ambiguities that still clouded acid-base theory, when he presented a definition of an acid in terms of the reaction scheme,
A B + H+
where A is an acid, B is its conjugate base and the symbol H+ represents the proton, not the hydronium ion of variable composition found in water.
Born in the same year as Einstein (1879), Johannes Brönsted, the son of a Danish engineer, was Professor of Chemistry at the University of Copenhagen and an outspoken anti-Nazi who managed to survive the war in Europe.
CHEMByte 16: Carboxylic Acids and Vinegar … and "Mother of Vinegar"
Carboxylic acids are an especially important class of acidic substances that are distinguished by the presence of the -COOH (carboxyl) group, a cluster of atoms that imparts a general set of acidic properties to molecules in which it is found. Formic, acetic and benzoic acids are monoprotic carboxylic acids; oxalic acid is a diprotic carboxylic acid; citric acid is triprotic:
| HCOOH | formic acid | monoprotic | 1 carboxyl group |
| CH3COOH | acetic acid | monoprotic | 1 carboxyl group |
| C6H5COOH | benzoic acid | monoprotic | 1 carboxyl group |
| HOOC-COOH | oxalic acid | diprotic | 2 carboxyl groups |
| HOOCC(OH)(CH2COOH)2 | citric acid | triprotic | 3 carboxyl groups |
In each case, even though there are other hydrogen atoms in the molecule, only one carboxylic acid hydrogen atom from each -COOH group can be transferred to a base in an acid-base reaction. As a class of compounds, they are generally weakly acidic. Acetic acid is by far the most important in terms of quantity used and manufactured... on the order of 10 billion pounds annually selling for about $0.30/lb, exclusive of the amount produced by fermentation and sold as a dilute aqueous solution known as vinegar. In commercial markets and as a commodity chemical it appears as glacial acetic acid of about 99.5 percent purity, so-called because on cold days it freezes into an ice-like solid. The melting point of pure acetic acid is 16.6°C.
Vinegar, an approximately 5% aqueous solution of acetic acid, is one of the oldest known fermentation products, predated perhaps only by wine. First derived from the spoilage (air oxidation) of alcohol in wine, vinegar is largely produced by the oxidation of ethanol (ethyl alcohol) by bacteria of the Acetobacter genus and is used as a food preservative, medicinal agent, topical antiseptic, primitive antibiotic, and even today as a household cleaning agent. Since the end of World War II, however, several synthetic acetic acid processes have replaced the fermentation process, except vinegar for food uses, which is produced by a double fermentation process: first, an alcoholic fermentation of a sugary mash by a suitable yeast which is then followed by a second fermentation to oxidize the alcohol.
You can make your own vinegar if you can get hold of a starter culture of microorganisms commonly known as "mother of vinegar." The bacteria will convert alcohol in leftover red wine, for example, into acetic acid, the characteristic ingredient. Where do you get mother of vinegar? From a few tablespoons of unpasteurized red wine left in air for an hour or so. Then add the few tablespoons to a liter placed in a ceramic crock (or preferably a wooden bucket), cover with cheesecloth, and allow to sit undisturbed in a dark, cool place for a couple of weeks. And now the vinegar. Carefully (gently) float the mother of vinegar -- which should be plainly evident if the right strain of bacterium has done its job -- onto the top of the wine you have chosen to be fermented into vinegar. After being left to ferment for a couple of peaceful, undisturbed weeks, the vinegar should be ready for enjoyment. Note: If you are unsuccessful, it may be due to competition among many other airborne bacteria which can produce less palatable flavors. Should that be the case, try purchasing the mother of vinegar starter from a vinegar-making friend, a hobby shop that supplies home beer brewers, a local vinegar manufacturer.... or revert to a gourmet food shop and just buy some good stuff.
CHEMByte 17: The Color of Red Cabbage and Indicating Acidity at Home.
A vegetable that is readily available in the supermarket can easily and simply be used to indicate the relative acidity of common household chemical compounds and solutions. Red cabbage contains a dye known as anthocyanin which turns red in strongly acidic solutions and yellow in strongly alkaline solutions. The juice of the red cabbage is a universal indicator:
| red | pH 1 |
| pink | pH 4 |
| lavender | pH 7 |
| green | pH 10 |
| yellow | pH 13 |
The preparation is simple if you have a food blender and an automatic drip coffee pot equipped with a chemical-style filter. Cut half a cabbage into slices and place in the blender; add water to cover, and blend for 15-20 seconds at high speed. A frothy sludge will result which, when passed through the filter paper in your coffee maker, will leave you with a couple of hundred mL of indicator solution that should be a lavender color. Add about 20 mL of the indicator solution to solutions or slurries of the following and observe the results:
Rolaids, Tums, or Milk of Magnesia
Vinegar (Preferably white)
Bicarbonate of soda (Arm & Hammer)
Milk
Orange juice
Table salt
Tide (detergent)
Bar soap
Household ammonia
Club soda
Fresca (carbonated soft drink)
CHEMByte 18: Raining Acid
The pH of typical rainwater samples is a slightly acidic 5.5, lower than the neutral value of 7.0 that we might expect. The acid pH of normal rain results from CO2 in the atmosphere, which dissolves to form carbonic acid:
H2O (liq) + CO2(g) H2CO3 (aq)
Carbonic acid in turn ionizes to form H3O+ ions:
H2CO3 (aq) + H2O (liq) H3O+(aq) + HCO3- (aq)
However, the low Ka(1) for H2CO3 (4.3 x 10-7) and its low concentration combine to give only a very weakly acidic solution.
Burning sulfur-containing fuels leads to contamination of the atmosphere by large amounts of the sulfur oxides, SO2 and SO3. Sulfur trioxide is particularly harmful because of its reaction with water to form sulfuric acid:
SO3 (liq) + H2O (liq) H2SO4 (aq)
The product is rainwater that contains a strong acid and has a low pH, resulting in a devastating effect on pH-sensitive plant and animal life in freshwater lakes. Some lakes in northern New York State and as far south as Georgia are essentially lifeless as a result of their high acidity. Acid rain is especially prevalent at higher elevations along the crest of the Appalachians, where clouds literally dip down into mountain lakes.
In addition to polluting bodies of water, acid rain appears to dissolve clay minerals, releasing soluble aluminum salts that are harmful to trees. Furthermore, statues and buildings constructed of limestone are very susceptible to damage from acid rain. Limestone is composed of the salt CaCO3, which dissolves in acidic solutions:
CaCO3 (s) + H3O+ (aq) Ca2+ (aq) + HCO3- (aq) + H2O (liq)
Such acid rain reactions have inflicted irreversible damage on priceless art treasures and greatly increased the costs of maintaining many municipal buildings and public works.
Because scientists around the world have been studying acid rain and its environmental impact for decades, it is not surprising to find long lists of governmental regulations for emission of contributing air pollutants which industry -- especially steel, paper and pulp, and automotive -- agriculture, and the private citizen driving her car or heating his home must abide by. The Clean Air Act imposed specific controls on sulfur dioxide, which forms sulfurous acid, and nitrogen oxides, which also combine with water to form acids. Countries in The European Union have established similar regulations. What is surprising and disappointing is that the problem has not gone away. At best, it is a standoff; and in some cases, particularly forests, which show higher levels of damage than expected, things are getting worse. The question is why?
Once on the ground, acidic rainwater can be neutralized, diminishing environmental damage, by the action of compounds in the soil but only providing the "acid load" is not too great. As these waters trickle through, the buffering action (Section 8.5) takes over. However, the supply of the basic cations that do the buffering is limited.
So the more we learn, the clearer it becomes. Acid rain is still a problem. Simple solutions anticipated twenty-five years ago by the first generation to address these issues failed because of the unexpected complexity of the problem. The environmental dilemma continues!
CHEMByte 19: Buffers and Body Fluids
Body fluids such as blood and saliva function properly only over a very narrow pH range. Nature achieves this range by the use of buffers. For example, blood has two buffering systems. One is a phosphate buffer containing the ions H2PO4- and HPO42-. The dihydrogen phosphate ion (H2PO4- ) is a weak acid, ionizing in aqueous solution as follows:
H2PO4- + H2O(liq) H3O+(aq)+ HPO42- (aq)
The equilibrium constant for this reaction is Ka(2), the second proton dissociation for H3PO4, which is 6.2 x 10-8 and pKa(2) = 7.21. Since the average pH for blood is 7.40, then,
pH = pKa + log( [base]/[acid] )
7.40 = 7.21 + log( [HPO42-]/[H2PO4-] )
7.40 - 7.21 = 0.19 = log( [HPO42-]/[H2PO4-] )
log( [HPO42-]/[H2PO4-] ) = 100.19 = 1.55
which is the ratio of [HPO42-] to [H2PO4-] in the blood.
However, the [HPO42-]/[H2PO4-] salts are not the principal buffers in the blood. The total concentration of these two ions in the blood is only about one-twelfth of the total concentration of the components of the [HCO3-]/[H2CO3] buffer. Most of the buffer action in blood is performed by species involved in the carbon dioxide and hydrogen carbonate ion equilibrium:
CO2(g) + H2O (liq) H2CO3 (aq)
H2CO3 (aq) + H2O (liq) H3O+ (aq) + HCO3- (aq)
K for the second reaction is Ka(1) for H2CO3 which is 4.4 x 10-7 and pKa(1) = 6.36. The ratio of [HCO3-] to [H2CO3] can be calculated in the same way as for the phosphate ion buffer system:
7.40 = 6.36 + log ( [HCO3-]/[H2CO3] )
log ( [HCO3-]/[H2CO3] ) = 1.04
log ( [HCO3-]/[H2CO3] ) = 101.04 = 11.0
Note that the high ratio of [HCO3-] to [H2CO3] indicates that the buffer would have limited effectiveness since addition of acid would build up [H2CO3] and destroy the balance. However, the body has a mechanism to combat the buildup of [H2CO3] through decomposition of H2CO3 to H2O and CO2 which takes place very rapidly under the influence of the enzyme catalyst, carbonic anhydrase:
H2CO3(aq) 
H2O (liq) + CO2(g)
The CO2 is exhaled from the lungs. Failure to adequately remove CO2 due to respiratory disfunction results in a buildup of H2CO3 in the blood and a lowering of the pH. Thus, monitoring of blood pH can be used in the diagnosis of impaired respiration in premature babies and coronary care patients.
CHEMBytes: Additional Problems on Acids/Bases
- A saturated solution of lanthanum iodate, La(IO3)3, in pure water has a concentration of iodate ion equal to 2.07 X 10-3 mol L-1 at 25°C. Find [La3+] and the solubility product of lanthanum iodate.
- Calculate the solubility of lead sulfate in a solution whose pH is 1.0, given the solubility product for PbSO4 to be 1.8 X 10-8.
- An unknown student takes an unknown weight of an unknown acid, dissolves it in an unknown amount of water, and titrates it with a strong base of unknown concentration. When he has added 10.00 mL of base, he notices that the pH is 5.0. He continues the titration until he reaches the equivalence point for removal of on proton. At this point, his buret reads 22.22 mL. What is the dissociation constant of the acid?
- What is the pH of a solution prepared by adding 300 mL of 0.500 M H3PO4 to 250 mL of 0.300 M NaOH.