Hydrogencarbonates

With excess carbon dioxide, i,e, if the gas is passed through a solution of the hydroxide, a hydrogencarbonate is formed ... 

A further peculiarity of the Group I and II carbonates is the ability to form the hydrogencarbonate or bicarbonate ion HCOj ... 

Group II hydrogencarbonates have insufficient thermal stability for them to be isolated as solids. However, in areas where natural deposits of calcium and magnesium carbonates are found a reaction between the carbonate, water and carbon dioxide occurs ... 

When heated, sodium hydrogencarbonate readily decomposes evolving carbon dioxide, a reaction which leads to its use as baking powder when the carhon dioxide evolved aerates the dough. In the soda-ammonia process the carbon dioxide evolved is used to supplement the main carbon dioxide supply obtained by heating calcium carbonate ... 

In one process the carbon dioxide is removed using potassium carbonate solution, potassium hydrogencarbonate being produced ... 

This reaction is used as a test for carbon dioxide. Passage of an excess of carbon dioxide produces the soluble hydrogencarbonate ... 

The hydrogencarbonate ion, produced in nature by this reaction, is one of the main causes of temporary hardness in water. Carbon dioxide is fairly soluble in water, 1 cm dissolving 1.7 cm of the gas at stp. The variation of solubility with pressure does not obey Henry s law, since the reaction... 

The amount of carbonic acid present, undissociated or dissociated, is only about 1 of the total concentration of dissolved carbon dioxide. Carbonic acid, in l especi of its dissociation into hydrogen and hydrogencarbonate ions, is actually a stronger acid than acetic acid the dissociation constant is ... 

Hydrogencarbonates of sodium, potassium and ammonium are known in the solid state and show hydrogen bonding in the crystal ... 

Magnesium and calcium hydrogencarbonates are known in solution and are responsible for temporary hardness in water. 

If the hydrogencarbonate is in solution and the cation is Ca or Mg. the insoluble carbonate is precipitated this reaction may be used, therefore, to remove hardness in water by precipitation of Ca or Mg ions.) The ease of decomposition of hydrogencar-bonates affords a test to distinguish between a hydrogencarbonate and a carbonate carbon dioxide is evolved by a hydrogencarbonate, but not by a carbonate, if it is heated, either as the solid or in solution, on a boiling water bath. 

Lead(II) carbonate occurs naturally as cerussite. It is prepared in the laboratory by passing carbon dioxide through, or adding sodium hydrogencarbonate to. a cold dilute solution of leadfll) nitrate or lead(II) ethanoate ... 

In liquid nitric acid, hydrogen bonding gives a loose structure similar to that of hydrogencarbonate ions. However, although pure nitric acid does not attack metals readily and does not evolve carbon dioxide from a carbonate, it is a conducting liquid, and undergoes auto-ionisation thus ... 

Such water, and also that containing salts of multipositive metals, (usually sulphates), is said to be hard since it does not readily produce a lather with soap. Experiments with alkali metal salts can be performed to verify that the hardness is due to the presence of the multipositive metal ions and not to any of the anions present. The hardness due to calcium and magnesium hydrogencarbonates is said to be temporary since it can be removed by boiling ... 

The alkali metal hydroxides are also readily absorb CO2 and H2S to form carbonates (or hydrogencarbonates) and sulfides (or hydrogen-sulfides), and are extensively used to remove mercaptans from petroleum products. Amphoteric oxides such as those of Al, Zn, Sn and Pb react with MOH to form aluminates, zincates, stannates and plumbates, and even SiC>2 (and silicate glasses) are attacked. 

Flgure 4.7 Solubilities of alkali carbonates and bicarbonales (hydrogencarbonates). (H. Stephen and T. Stephen, Solubilities of Inorganic and Organic Compounds, Vol. 1, Part 1, Macmillan, New York.). 

A mixture of 4.0 grams of N-methyl-3-toluidine and 2.8 grams of sodium hydrogencarbon-ate in 50 cc of acetone was stirred at 0° to 10°C and 7.4 grams of 2-naphthyl chlorothiono-formate was added in small portions thereto and the mixture was heated under reflux for 30 minutes. The cooled mixture was poured into about 150 cc of cold water and 2-naphthvl-N-methvl-N-(3-tolyl)thionocarbamate was obtained as white crystals. Yield is 9.1 grams (90 o). Recrystallization from alcohol gave colorless needle crystals, MP 110.5° to 111.5°C. 

Some solids are either too soluble, or the solubility does not vary sufficiently with temperature, in a given solvent for direct crystallisation to be practicable. In many cases, the solid can be precipitated from, say, a concentrated aqueous solution by the addition of a liquid, miscible with water, in which it is less soluble. Ethanol, in which many inorganic compounds are almost insoluble, is generally used. Care must be taken that the amount of ethanol or other solvent added is not so large that the impurities are also precipitated. Potassium hydrogencarbonate and antimony potassium tartrate may be purified by this method. 

Commercial samples of water are frequently alkaline due to the presence of hydrogencarbonates, carbonates, or hydroxides. The alkalinity is determined by titrating a 100.0 mL sample with 0.02M hydrochloric acid using screened methyl orange as indicator (or to a pH of 3.8). To obtain the total cation content in terms of CaC03, the total methyl orange alkalinity is added to the EMA. 

Titration of carbonate ion with a strong acid. A solution of sodium carbonate may be titrated to the hydrogencarbonate stage (i.e. with one mole of hydrogen ions), when the net reaction is ... 

The same end point is reached by titrating sodium hydrogencarbonate solution with hydrochloric acid ... 

The end point with 100 mL of 0.2M sodium hydrogencarbonate and 0.2M hydrochloric acid may be deduced as follows from the known dissociation constant and concentration of the weak acid. The end point will obviously occur when 100 mL of hydrochloric acid has been added, i.e. the solution now has a total volume of 200 mL. Consequently since the carbonic acid liberated from the sodium hydrogencarbonate (0.02 moles) is now contained in a volume of 200 mL, its concentration is 0.1 M. Kl for carbonic acid has a value of 4.3 x 10 7, and hence we can say ... 

The two methods available for this determination are modifications of those described in Section 10.32 for hydroxide/carbonate mixtures. In the first procedure, which is particularly valuable when the sample contains relatively large amounts of carbonate and small amounts of hydrogencarbonate, the total alkali is first determined in one portion of the solution by titration with standard 0.1M hydrochloric acid using methyl orange, methyl orange-indigo carmine, or bromophenol blue as indicator ... 

Let this volume correspond to V mL 1M HC1. To another sample, a measured excess of standard 0.1M sodium hydroxide (free from carbonate) over that required to transform the hydrogencarbonate to carbonate is added ... 

A slight excess of 10 per cent barium chloride solution is added to the hot solution to precipitate the carbonate as barium carbonate, and the excess of sodium hydroxide solution immediately determined, without filtering off the precipitate, by titration with the same standard acid phenolphthalein or thymol blue is used as indicator. If the volume of excess of sodium hydroxide solution added corresponds to timL of 1M sodium hydroxide and u mL 1M acid corresponds to the excess of the latter, then v — v = hydrogencarbonate, and V— v — v ) = carbonate. 

Another sample of equal volume is then titrated with the same standard acid using methyl orange, methyl orange-indigo carmine or bromophenol blue as indicator. The volume of acid used (say, ymL) corresponds to carbonate+ hydrogencarbonate. Hence 2y = carbonate, and y — 2Y = hydrogencarbonate. 

Either the Mohr titration or the adsorption indicator method may be used for the determination of chlorides in neutral solution by titration with standard 0.1M silver nitrate. If the solution is acid, neutralisation may be effected with chloride-free calcium carbonate, sodium tetraborate, or sodium hydrogencarbonate. Mineral acid may also be removed by neutralising most ofthe acid with ammonia solution and then adding an excess of ammonium acetate. Titration of the neutral solution, prepared with calcium carbonate, by the adsorption indicator method is rendered easier by the addition of 5 mL of 2 per cent dextrin solution this offsets the coagulating effect of the calcium ion. If the solution is basic, it may be neutralised with chloride-free nitric acid, using phenolphthalein as indicator. 

As the flask cools, the hydrogencarbonate solution is automatically drawn in until the pressure of the carbon dioxide inside the flask is equal to the atmospheric pressure. 

Procedure. Wdgh out accurately about 2.5g of finely powdered arsenic(III) oxide, transfer to a 500 mL beaker, and dissolve it in a concentrated solution of sodium hydroxide, prepared from 2g of iron-free sodium hydroxide and 20 mL of water. Dilute to about 200 mL, and neutralise the solution with 1M hydrochloric add, using a pH meter. When the solution is faintly add transfer the contents of the beaker quantitatively to a 500 mL graduated flask, add 2 g of pure sodium hydrogencarbonate, and, when all the salt has dissolved, dilute to the mark and shake well. 

Using a burette or a pipette with a safety pump (this is necessary owing to the poisonous properties of the solution) measure out 25.0 mL of the arsenite solution into a 250 mL conical flask, add 25-50 mL of water, 5g of sodium hydrogencarbonate, and 2 mL of starch solution. Swirl the solution carefully until the hydrogencarbonate has dissolved. Then titrate slowly with the iodine solution, contained in a burette, to the first blue colour. 

If it is desired to base the standardisation directly upon arsenic(III) oxide, proceed as follows. Weigh out accurately about 0.20 g of pure arsenic(III) oxide into a conical flask, dissolve it in 10 mL of 1M sodium hydroxide, and add a small excess of dilute sulphuric acid (say, 12-15 mL of 0.5M acid). Mix thoroughly and cautiously. Then add carefully a solution of 2 g of sodium hydrogencarbonate in 50 mL of water, followed by 2 mL of starch solution. Titrate slowly with the iodine solution to the first blue colour. Repeat with two other similar quantities of the oxide. 

This reaction is subject to a number of errors (1) the hydriodic acid (from excess of iodide and acid) is readily oxidised by air, especially in the presence of chromium(III) salts, and (2) it is not instantaneous. It is accordingly best to pass a current of carbon dioxide through the reaction flask before and during the titration (a more convenient but less efficient method is to add some solid sodium hydrogencarbonate to the acid solution, and to keep the flask covered as much as possible), and to allow 5 minutes for its completion. 

---