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1. Electrolysis of sodium chloride. Fig. 32 represents a vessel in which the electrolyte is a solution of sodium chloride (NaCl). According to the dissociation theory the molecules of sodium chloride dissociate into the ions Na^{+} and Cl^{-}. The Na^{+} ions are attracted to the cathode owing to its large negative charge. On coming into contact with the cathode, the Na^{+} ions give up their positive charge and are then ordinary sodium atoms. They immediately decompose the water according to the equation
Na + H_{2}O = NaOH + H,
and hydrogen is evolved about the cathode.
The chlorine ions on being discharged at the anode in similar manner may either be given off as chlorine gas, or may attack the water, as represented in the equation
2Cl + H_{2}O = 2HCl + O.
2. Electrolysis of water. The reason for the addition of sulphuric acid to water in the preparation of oxygen and hydrogen by electrolysis will now be clear. Water itself is not an electrolyte to an appreciable extent; that is, it does not form enough ions to carry a current. Sulphuric acid dissolved in water is an electrolyte, and dissociates into the ions 2 H^{+} and SO{4}^{—}. In the process of electrolysis of the solution, the hydrogen ions travel to the cathode, and on being discharged escape as hydrogen gas. The SO{4} ions, when discharged at the anode, act upon water, setting free oxygen and once more forming sulphuric acid:
SO{4} + H{2}O = H{2}SO{4} + O.
The sulphuric acid can again dissociate and the process repeat itself as long as any water is left. Hence the hydrogen and oxygen set free in the electrolysis of water really come directly from the acid but indirectly from the water.
3. Electrolysis of sodium sulphate. In a similar way, sodium sulphate (Na{2}SO{4}), when in solution, gives the ions 2 Na^{+} and SO{4}^{—}. On being discharged, the sodium atoms decompose water about the cathode, as in the case of sodium chloride, while the SO{4} ions when discharged at the anode decompose the water, as represented in the equation
SO{4} + H{2}O = H{2}SO{4} + O
That new substances are formed at the cathode and anode may be shown in the following way. A U-tube, such as is represented in Fig. 33, is partially filled with a solution of sodium sulphate, and the liquid in one arm is colored with red litmus, that in the other with blue litmus. An electrode placed in the red solution is made to serve as cathode, while one in the blue solution is made the anode. On allowing the current to pass, the blue solution turns red, while the red solution turns blue. These are exactly the changes which would take place if sodium hydroxide and sulphuric acid were to be set free at the electrodes, as required by the theory.
The properties of electrolytes depend upon the ions present. When a substance capable of dissociating into ions is dissolved in water, the properties of the solution will depend upon two factors: (1) the ions formed from the substance; (2) the undissociated molecules. Since the ions are usually more active chemically than the molecules, most of the chemical properties of an electrolyte are due to the ions rather than to the molecules.
The solutions of any two substances which give the same ion will have certain properties in common. Thus all solutions containing the copper ion (Cu^{+}) are blue, unless the color is modified by the presence of ions or molecules having some other color.
EXERCISES
1. Distinguish clearly between the following terms: electrolysis, electrolyte, electrolytic dissociation, ions, solute, solvent, solution, saturated solution, and supersaturated solution.
2. Why does the water from some natural springs effervesce?
3. (a) Why does not the water of the ocean freeze? (b) Why will ice and salt produce a lower temperature than ice alone?
4. Why does shaking or stirring make a solid dissolve more rapidly in a liquid?
5. By experiment it was found that a certain volume of water was saturated at 100 deg. with 114 g. of potassium nitrate. On cooling to 0 deg. a portion of the substance crystallized. (a) How many grams of the substance remained in solution? (b) What was the strength of the solution at 18 deg.? (c) How much water had been used in the experiment?
6. (a) 10 g. of common salt were dissolved in water and the solution evaporated to dryness; what weight of solid was left? (b) 10 g. of zinc were dissolved in hydrochloric acid and the solution evaporated to dryness; what weight of solid was left?
7. Account for the fact that sugar sometimes deposits from molasses, even when no evaporation has taken place.
8. (a) From the standpoint of the theory of electrolytic dissociation, write the simple equation for a dilute solution of copper sulphate (CuSO_{4}); this solution is blue. (b) In the same manner, write one for sodium sulphate; this solution is colorless. (c) How would you account for the color of the copper sulphate solution?
9. (a) As in the preceding exercise, write a simple equation for a dilute solution of copper chloride (CuCl_{2}); this solution is blue. (b) In the same manner, write one for sodium chloride; this solution is colorless. To what is the blue color due?
10. What component is present in concentrated sulphuric acid that is almost wanting in very dilute sulphuric acid?
11. Why will vegetables cook faster when boiled in strong salt water than when boiled in pure water?
12. How do you explain the foaming of soda water?
CHAPTER X
ACIDS, BASES, AND SALTS; NEUTRALIZATION
Acids, bases, and salts. The three classes of compounds known respectively as acids, bases, and salts include the great majority of the compounds with which we shall have to deal. It is important, therefore, for us to consider each of these classes in a systematic way. The individual members belonging to each class will be discussed in detail in the appropriate places, but a few representatives of each class will be described in this chapter with special reference to the common properties in accordance with which they are classified.
The familiar acids. _Hydrochloric acid_ is a gas composed of hydrogen and chlorine, and has the formula HCl. The substance is very soluble in water, and it is this solution which is usually called hydrochloric acid. _Nitric acid_ is a liquid composed of hydrogen, nitrogen, and oxygen, having the formula HNO_{3}. As sold commercially it is mixed with about 32% of water. _Sulphuric acid_, whose composition is represented by the formula H_{2}SO_{4}, is an oily liquid nearly twice as heavy as water, and is commonly called _oil of vitriol_.
Characteristics of acids. (1) All acids contain hydrogen. (2) When dissolved in water the molecules of the acid dissociate into two kinds of ions. One of these is always hydrogen and is the cation (+), while the other consists of the remainder of the molecule and is the anion (-). (3) The solution tastes sour. (4) It has the power to change the color of certain substances called indicators. Thus blue litmus is changed to red, and yellow methyl orange is changed to red. Since all acids produce hydrogen cations, while the anions of each are different, the properties which all acids have in common when in solution, such as taste and action on indicators, must be attributed to the hydrogen ions.
DEFINITION: An acid is a substance which produces hydrogen ions when dissolved in water or other dissociating liquids.
Undissociated acids. When acids are perfectly free from water, or are dissolved in liquids like benzene which do not have the power of dissociating them into ions, they should have no real acid properties. This is found to be the case. Under these circumstances they do not affect the color of indicators or have any of the properties characteristic of acids.
The familiar bases. The bases most used in the laboratory are sodium hydroxide (NaOH), potassium hydroxide (KOH), and calcium hydroxide (Ca(OH)_{2}). These are white solids, soluble in water, the latter sparingly so. Some bases are very difficultly soluble in water. The very soluble ones with most pronounced basic properties are sometimes called the _alkalis_.
Characteristics of bases. (1) All bases contain hydrogen and oxygen. (2) When dissolved in water the molecules of the base dissociate into two kinds of ions. One of these is always composed of oxygen and hydrogen and is the anion. It has the formula OH and is called the hydroxyl ion. The remainder of the molecule, which usually consists of a single atom, is the cation. (3) The solution of a base has a soapy feel and a brackish taste. (4) It reverses the color change produced in indicators by acids, turning red litmus blue, and red methyl orange yellow. Since all bases produce hydroxyl anions, while the cations of each are different, the properties which all bases have in common when in solution must be due to the hydroxyl ions.
DEFINITION: A base is a substance which produces hydroxyl ions when dissolved in water or other dissociating liquids.
Undissociated bases. Bases, in the absence of water or when dissolved in liquids which do not dissociate them, should have none of the properties characteristic of this class of substances. This has been found to be the case. For example, they have no effect upon indicators under these circumstances.
Neutralization. When an acid and a base are brought together in solution in proper proportion, the characteristic properties of each disappear. The solution tastes neither sour nor brackish; it has no effect upon indicators. There can therefore be neither hydrogen nor hydroxyl ions present in the solution. A study of reactions of this kind has shown that the hydrogen ions of the acid combine with the hydroxyl ions of the base to form molecules of water, water being a substance which is not appreciably dissociated into ions. This action of an acid on a base is called neutralization. The following equations express the neutralization of the three acids by three bases, water being formed in each case.
Na^{+}, OH^{-} + H^{+}, Cl^{-} = Na^{+}, Cl^{-} + H_{2}O.
K^{+}, OH^{-} + H^{+}, NO_{3}^{-} = K^{+}, NO_{3}^{-} + H_{2}O.
Ca^{+}, (OH)_{2}^{—} + H_{2}^{+}, SO_{4}^{—} = Ca^{+}, SO_{4}^{—} + 2H_{2}O.
DEFINITION: Neutralization consists in the union of the hydrogen ion of an acid with the hydroxyl ion of a base to form water.
Salts. It will be noticed that in neutralization the anion of the acid and the cation of the base are not changed. If, however, the water is expelled by evaporation, these two ions slowly unite, and when the water becomes saturated with the substance so produced, it separates in the form of a solid called a salt.
DEFINITION: A salt is a substance formed by the union of the anion of an acid with the cation of a base.
Characteristics of salts. (1) From the definition of a salt it will be seen that there is no element or group of elements which characterize salts. (2) Salts as a class have no peculiar taste. (3) In the absence of all other substances they are without action on indicators. (4) When dissolved in water they form two kinds of ions.
Heat of neutralization. If neutralization is due to the union of hydrogen ions with hydroxyl ions, and nothing more, it follows that when a given weight of water is formed in neutralization, the heat set free should always be the same, no matter from what acid and base the two kinds of ions have been supplied. Careful experiments have shown that this is the case, provided no other reactions take place at the same time. When 18g. of water are formed in neutralization, 13,700 cal. of heat are set free. This is represented in the equations
Na^{+}, OH^{-} + H^{+}, Cl^{-} = Na^{+}, Cl^{-} + H_{2}O + 13,700 cal.
K^{+}, OH^{-} + H^{+}, NO_{3}^{-} = K^{+}, NO_{3}^{-} + H_{2}O + 13,700 cal.
Ca^{+}, (OH)_{2}^{—} + H_{2}^{+}, SO_{4}^{—} = Ca^{+}, SO_{4}^{—} + 2H_{2}O + 2 x 13,700 cal.
Neutralization a quantitative act. Since neutralization is a definite chemical act, each acid will require a perfectly definite weight of each base for its neutralization. For example, a given weight of sulphuric acid will always require a definite weight of sodium hydroxide, in accordance with the equation
H_{2}, SO_{4} + 2Na, OH = Na_{2}, SO_{4} + 2H_{2}O.
Determination of the ratio in neutralization. The quantities of acid and base required in neutralization may be determined in the following way. Dilute solutions of the two substances are prepared, the sulphuric acid being placed in one of the burettes (Fig. 34) and the sodium hydroxide in the other. The levels of the two liquids are then brought to the zero marks of the burettes by means of the stopcocks. A measured volume of the acid is drawn off into a beaker, a few drops of litmus solution added, and the sodium hydroxide is run in drop by drop until the red litmus just turns blue. The volume of the sodium hydroxide consumed is then noted. If the concentrations of the two solutions are known, it is easy to calculate what weight of sodium hydroxide is required to neutralize a given weight of sulphuric acid. By evaporating the neutralized solution to dryness, the weight of the sodium sulphate formed can be determined directly. Experiment shows that the weights are always in accordance with the equation in the preceding paragraph.
Extent of dissociation. The question will naturally arise, When an acid, base, or salt dissolves in water, do all the molecules dissociate into ions, or only a part of them? The experiments by which this question can be answered cannot be described here. It has been found, however, that only a fraction of the molecules dissociate. The percentage which will dissociate in a given case depends upon several conditions, the chief of which are: (1) The concentration of the solution. In concentrated solutions only a very small percentage of dissociation occurs. As the solution is diluted the percentage increases, and in very dilute solutions it may be very large, though it is never complete in any ordinary solution. (2) The nature of the dissolved compound. At equal concentrations substances differ much among themselves in the percentage of dissociation. The great majority of salts are about equally dissociated. Acids and bases, on the contrary, show great differences. Some are freely dissociated, while others are dissociated to but a slight extent.
Strength of acids and bases. Since acid and basic properties are due to hydrogen and hydroxyl ions respectively, the acid or base which will produce the greatest percentage of these ions at a given concentration must be regarded as the strongest representative of its class. The acids and bases described in the foregoing paragraphs are all quite strong. In 10% solutions they are dissociated to about 50%, and this is also approximately the extent to which most salts are dissociated at this same concentration.
Partial neutralization. 1. _Basic salts._ The chemical action between an acid and a base is not always as complete as has been represented in the foregoing paragraphs. For example, if the base magnesium hydroxide (Mg(OH)_{2}) and hydrochloric acid (HCl) are brought together in the ratio of an equal number of molecules of each, there will be only half enough hydrogen ions for the hydroxyl ions present.
Mg, (OH){2} + H, Cl = Mg, OH, Cl + H{2}O.
Magnesium, hydroxyl, and chlorine ions are left at the close of the reaction, and under the proper conditions unite to form molecules of the compound Mg(OH)Cl. This compound, when dissolved, can form hydroxyl ions and therefore possesses basic properties; it can also form the ions of a salt (Mg and Cl), and has properties characteristic of salts. Substances of this kind are called basic salts.
DEFINITION: A basic salt is a substance which can give the ions both of a base and of a salt when dissolved in water.
2. Acid salts. In a similar way, when sulphuric acid and sodium hydroxide are brought together in the ratio of equal numbers of the molecules of each, it is possible to have a reaction expressed by the equation
Na, OH + H{2}, SO{4} = Na, H, SO{4} + H{2}O.
The ions remaining after all the hydroxyl ions have been used up are those of an acid (H) and those of a salt (Na and SO{4}). These unite to form the substance NaHSO{4}, and as the solution becomes saturated with this substance through evaporation, it separates in the form of crystals. In solution this substance can give hydrogen ions, and therefore possesses acid properties; it can also give the ions characteristic of a salt. It is therefore called an acid salt.
DEFINITION: An acid salt is one which can give the ions of an acid and of a salt when in solution.
3. Normal salts. Salts which are the products of complete neutralization, such as Na{2}SO{4}, and which in solution can give neither hydrogen nor hydroxyl ions, but only the ions of a salt, are called normal salts to distinguish them from acid and basic salts.
Methods of expressing reactions between compounds in solution. Chemical equations representing reactions between substances in solution may represent the details of the reaction, or they may simply indicate the final products formed. In the latter case the formation of ions is not indicated. Thus, if we wish to call attention to the details of the reaction between sodium hydroxide and hydrochloric acid in solution, the equation is written as follows:
Na^{+}, OH^{-} + H^{+}, Cl^{-} = Na^{+}, Cl^{-} + H_{2}O.
On the other hand, if we wish simply to represent the final products formed, the following is used.
NaOH + HCl = NaCl + H_{2}O.
Both of these methods will therefore be used:
Radicals. It has been emphasized that the hydroxyl group (OH) always forms the anion of a base, while the group NO_{3} forms the anion of nitric acid and sodium nitrate; the group SO_{4}, the anion of sulphuric acid and calcium sulphate. A group of elements which in this way constitutes a part of a molecule, acting as a unit in a chemical change, or forming ions in solution, is called a _radical_. Some of these radicals have been given special names, the names signifying the elements present in the radical. Thus we have the hydroxyl radical (OH) and the nitrate radical (NO_{3}).
DEFINITION: A radical is a group of elements forming part of a molecule, and acting as a unit in chemical reactions.
Names of acids, bases, and salts. Since acids, bases, and salts are so intimately related to each other, it is very advantageous to give names to the three classes in accordance with some fixed system. The system universally adopted is as follows:
Naming of bases. All bases are called hydroxides. They are distinguished from each other by prefixing the name of the element which is in combination with the hydroxyl group. Examples: sodium hydroxide (NaOH); calcium hydroxide (Ca(OH){2}); copper hydroxide (Cu(OH){2}).
Naming of acids. The method of naming acids depends upon whether the acid consists of two elements or three.
1. _Binary acids._ Acids containing only one element in addition to hydrogen are called _binary acids_. They are given names consisting of the prefix _hydro-_, the name of the second element present, and the termination _-ic_. Examples: hydrochloric acid (HCl); hydrosulphuric acid (H_{2}S).
2. Ternary acids. In addition to the two elements present in binary acids, the great majority of acids also contain oxygen. They therefore consist of three elements and are called ternary acids. It usually happens that the same three elements can unite in different proportions to make several different acids. The most familiar one of these is given a name ending in the suffix -ic, while the one with less oxygen is given a similar name, but ending in the suffix -ous. Examples: nitric acid (HNO{3}); nitrous acid (HNO{2}). In cases where more than two acids are known, use is made of prefixes in addition to the two suffixes -ic and -ous. Thus the prefix per- signifies an acid still richer in oxygen; the prefix hypo- signifies one with less oxygen.
Naming of salts. A salt derived from a binary acid is given a name consisting of the names of the two elements composing it, with the termination -ide. Example: sodium chloride (NaCl). All other binary compounds are named in the same way.
A salt of a ternary acid is named in accordance with the acid from which it is derived. A ternary acid with the termination -ic gives a salt with the name ending in -ate, while an acid with termination -ous gives a salt with the name ending in -ite. The following table will make the application of these principles clear:
ACIDS SYMBOL SALTS SYMBOL
Hydrochloric HCl Sodium chloride NaCl Hypochlorous HClO Sodium hypochlorite NaClO Chlorous HClO{2} Sodium chlorite NaClO{2} Chloric HClO{3} Sodium chlorate NaClO{3} Perchloric HClO{4} Sodium perchlorate NaClO{4}
EXERCISES
1. 25 cc. of a solution containing 40 g. of sodium hydroxide per liter was found to neutralize 25 cc. of a solution of hydrochloric acid. What was the strength of the acid solution?
2. After neutralizing a solution of sodium hydroxide with nitric acid, there remained after evaporation 100 g. of sodium nitrate. How much of each substance had been used?
3. A solution contains 18 g. of hydrochloric acid per 100 cc. It required 25 cc. of this solution to neutralize 30 cc. of a solution of sodium hydroxide. What was the strength of the sodium hydroxide solution in parts per hundred?
4. When perfectly dry sulphuric acid is treated with perfectly dry sodium hydroxide, no chemical change takes place. Explain.
5. When cold, concentrated sulphuric acid is added to zinc, no change takes place. Recall the action of dilute sulphuric acid on the same metal. How do you account for the difference?
6. A solution of hydrochloric acid in benzene does not conduct the electric current. When this solution is treated with zinc, will hydrogen be evolved? Explain.
7. (a) Write equation for preparation of hydrogen from zinc and dilute sulphuric acid. (b) Rewrite the same equation from the standpoint of the theory of electrolytic dissociation, (c) Subtract the common SO_{4} ion from both members of the equation, (d) From the resulting equation, explain in what the preparation of hydrogen consists when examined from the standpoint of this theory.
8. In the same manner as in the preceding exercise, explain in what the action of sodium on water to give hydrogen consists.
CHAPTER XI
VALENCE
Definition of valence. A study of the formulas of various binary compounds shows that the elements differ between themselves in the number of atoms of other elements which they are able to hold in combination. This is illustrated in the formulas
HCl, H_{2}O, H_{3}N, H_{4}C. (hydrochloric acid) (water) (ammonia) (marsh gas)
It will be noticed that while one atom of chlorine combines with one atom of hydrogen, an atom of oxygen combines with two, an atom of nitrogen with three, one of carbon with four. The number which expresses this combining ratio between atoms is a definite property of each element and is called its valence.
DEFINITION: The valence of an element is that property which determines the number of the atoms of another element which its atom can hold in combination.
Valence a numerical property. Valence is therefore merely a numerical relation and does not convey any information in regard to the intensity of the affinity between atoms. Judging by the heat liberated in their union, oxygen has a far stronger affinity for hydrogen than does nitrogen, but an atom of oxygen can combine with two atoms only of hydrogen, while an atom of nitrogen can combine with three.
Measure of valence. In expressing the valence of an element we must select some standard for comparison, just as in the measurement of any other numerical quantity. It has been found that an atom of hydrogen is never able to hold in combination more than one atom of any other element. Hydrogen is therefore taken as the standard, and other elements are compared with it in determining their valence. A number of other elements are like hydrogen in being able to combine with at most one atom of other elements, and such elements are called univalent. Among these are chlorine, iodine, and sodium. Elements such as oxygen, calcium, and zinc, which can combine with two atoms of hydrogen or other univalent elements, are said to be divalent. Similarly, we have trivalent, tetravalent, pentavalent elements. None have a valence of more than 8.
Indirect measure of valence. Many elements, especially among the metals, do not readily form compounds with hydrogen, and their valence is not easy to determine by direct comparison with the standard element. These elements, however, combine with other univalent elements, such as chlorine, and their valence can be determined from the compounds so formed.
Variable valence. Many elements are able to exert different valences under differing circumstances. Thus we have the compounds Cu{2}O and CuO, CO and CO{2}, FeCl{2} and FeCl{3}. It is not always possible to assign a fixed valence to an element. Nevertheless each element tends to exert some normal valence, and the compounds in which it has a valence different from this are apt to be unstable and easily changed into compounds in which the valence of the element is normal. The valences of the various elements will become familiar as the elements are studied in detail.
Valence and combining ratios. When elements combine to form compounds, the ratio in which they combine will be determined by their valences. In those compounds which consist of two elements directly combined, the union is between such numbers of the two atoms as have equal valences. Elements of the same valence will therefore combine atom for atom. Designating the valence of the atoms by Roman numerals placed above their symbols, we have the formulas
II II II III III IV IV HCl, ZnO, BN, CSi.
A divalent element, on the other hand, will combine with two atoms of a univalent element. Thus we have
II II II II ZnCl{2} and H{2}O
(the numerals above each symbol representing the sum of the valences of the atoms of the element present). A trivalent atom will combine with three atoms of a univalent element, as in the compound
III III H_{3}N.
If a trivalent element combines with a divalent element, the union will be between two atoms of the trivalent element and three of the divalent element, since these numbers are the smallest which have equal valences. Thus the oxide of the trivalent metal aluminium has the formula Al{2}O{3}. Finally one atom of a tetravalent element such as carbon will combine with four atoms of a univalent element, as in the compound CH{4}, or with two atoms of a divalent element, as in the compound CO{2}.
We have no knowledge as to why elements differ in their combining power, and there is no way to determine their valences save by experiment.
Valence and the structure of compounds. Compounds will be met from time to time which are apparent exceptions to the general statements just made in regard to valence. Thus, from the formula for hydrogen dioxide (H_{2}O_{2}), it might be supposed that the oxygen is univalent; yet it is certainly divalent in water (H_{2}O). That it may also be divalent in H_{2}O_{2} may be made clear as follows: The unit valence of each element may be represented graphically by a line attached to its symbol. Univalent hydrogen and divalent oxygen will then have the symbols H- and -O-. When atoms combine, each unit valence of one atom combines with a unit valence of another atom. Thus the composition of water may be expressed by the formula H-O-H, which is meant to show that each of the unit valences of oxygen is satisfied with the unit valence of a single hydrogen atom.
The chemical conduct of hydrogen dioxide leads to the conclusion that the two oxygen atoms of its molecule are in direct combination with each other, and in addition each is in combination with a hydrogen atom. This may be expressed by the formula H-O-O-H. The oxygen in the compound is therefore divalent, just as it is in water. It will thus be seen that the structure of a compound must be known before the valences of the atoms making up the compound can be definitely decided upon.
Such formulas as H-O-H and H-O-O-H are known as structural formulas, because they are intended to show what is known in regard to the arrangement of the atoms in the molecules.
Valence and the replacing power of atoms. Just as elements having the same valence combine with each other atom for atom, so if they replace each other in a chemical reaction they will do so in the same ratio. This is seen in the following equations, in which a univalent hydrogen atom is replaced by a univalent sodium atom:
NaOH + HCl = NaCl + H_{2}O.
2NaOH + H_{2}SO_{4} = Na_{2}SO_{4} + 2H_{2}O.
Na + H_{2}O = NaOH + H.
Similarly, one atom of divalent calcium will replace two atoms of univalent hydrogen or one of divalent zinc:
Ca(OH)_{2} + 2 HCl = CaCl_{2} + 2H_{2}O.
CaCl{2} + ZnSO{4} = CaSO{4} + ZnCl{2}.
In like manner, one atom of a trivalent element will replace three of a univalent element, or two atoms will replace three atoms of a divalent element.
Valence and its applications to formulas of salts. While the true nature of valence is not understood and many questions connected with the subject remain unanswered, yet many of the main facts are of much help to the student. Thus the formula of a salt, differs from that of the acid from which it is derived in that the hydrogen of the acid has been replaced by a metal. If, then, it is known that a given metal forms a normal salt with a certain acid, the formula of the salt can at once be determined if the valence of the metal is known. Since sodium is univalent, the sodium salts of the acids HCl and H{2}SO{4} will be respectively NaCl and Na{2}SO{4}. One atom of divalent zinc will replace 2 hydrogen atoms, so that the corresponding zinc salts will be ZnCl{2} and ZnSO{4}.
The formula for aluminium sulphate is somewhat more difficult to determine. Aluminium is trivalent, and the simplest ratio in which the aluminium atom can replace the hydrogen in sulphuric acid is 2 atoms of aluminium (6 valences) to 3 molecules of sulphuric acid (6 hydrogen atoms). The formula of the sulphate will then be Al_{2}(SO_{4})_{3}.
Valence and its application to equation writing. It will be readily seen that a knowledge of valence is also of very great assistance in writing the equations for reactions of double decomposition. Thus, in the general reaction between an acid and a base, the essential action is between the univalent hydrogen ion and the univalent hydroxyl ion. The base and the acid must always be taken in such proportions as to secure an equal number of each of these ions. Thus, in the reaction between ferric hydroxide (Fe(OH)_{3}) and sulphuric acid (H_{2}SO_{4}), it will be necessary to take 2 molecules of the former and 3 of the latter in order to have an equal number of the two ions, namely, 6. The equation will then be
2Fe(OH)_{3} + 3H_{2}SO_{4} = Fe_{2}(SO_{4})_{3} + 6H_{2}O.
Under certain conditions the salts Al{2}(SO{4}){3} and CaCl{2} undergo double decomposition, the two metals, aluminium and calcium, exchanging places. The simplest ratio of exchange in this case is 2 atoms of aluminium (6 valences) and 3 atoms of calcium (6 valences). The reaction will therefore take place between 1 molecule of Al{2}(SO{4}){3} and 3 of CaCl{2}, and the equation is as follows:
Al{2}(SO{4}){3} + 3 CaCl{2} = 3CaSO{4} + 2AlCl{3}.
EXERCISES
1. Sodium, calcium, and aluminium have valences of 1, 2, and 3 respectively; write the formulas of their chlorides, sulphates, and phosphates (phosphoric acid = H{3}PO{4}), on the supposition that they form salts having the normal composition.
2. Iron forms one series of salts in which it has a valence of 2, and another series in which it has a valence of 3; write the formulas for the two chlorides of iron, also for the two sulphates, on the supposition that these have the normal composition.
3. Write the equation representing the neutralization of each of the following bases by each of the acids whose formulas are given:
NaOH HCl Ba(OH){2} H{2}SO{4} Al(OH){3} H{3}PO{4}
4. Silver acts as a univalent element and calcium as a divalent element in the formation of their respective nitrates and chlorides. (a) Write the formula for silver nitrate; for calcium chloride. (b) When solutions of these two salts are mixed, the two metals, silver and calcium, exchange places; write the equation for the reaction.
_5._ Antimony acts as a trivalent element in the formation of a chloride. (a) What is the formula for antimony chloride? (b) When hydrosulphuric acid (H_{2}S) is passed into a solution of this chloride the hydrogen and antimony exchange places; write the equation for the reaction.
6. Lead has a valence of 2 and iron of 3 in the compounds known respectively as lead nitrate and ferric sulphate. (a) Write the formulas for these two compounds. (b) When their solutions are mixed the two metals exchange places; write the equation for the reaction.
CHAPTER XII
COMPOUNDS OF NITROGEN
Occurrence. As has been stated in a former chapter, nitrogen constitutes a large fraction of the atmosphere. The compounds of nitrogen, however, cannot readily be obtained from this source, since at any ordinary temperature nitrogen is able to combine directly with very few of the elements.
In certain forms of combination nitrogen occurs in the soil from which it is taken up by plants and built into complex substances composed chiefly of carbon, hydrogen, oxygen, and nitrogen. Animals feeding on these plants assimilate the nitrogenous matter, so that this element is an essential constituent of both plants and animals.
Decomposition of organic matter by bacteria. When living matter dies and undergoes decay complicated chemical reactions take place, one result of which is that the nitrogen of the organic matter is set free either as the element nitrogen, or in the form of simple compounds, such as ammonia (NH_{3}) or oxides of nitrogen. Experiment has shown that all such processes of decay are due to the action of different kinds of bacteria, each particular kind effecting a different change.
Decomposition of organic matter by heat. When organic matter is strongly heated decomposition into simpler substances takes place in much the same way as in the case of bacterial decomposition. Coal is a complex substance of vegetable origin, consisting largely of carbon, but also containing hydrogen, oxygen, and nitrogen. When this is heated in a closed vessel so that air is excluded, about one seventh of the nitrogen is converted into ammonia, and this is the chief source from which ammonia and its compounds are obtained.
COMPOUNDS OF NITROGEN WITH HYDROGEN
Ammonia (NH_{3}). Several compounds consisting exclusively of nitrogen and hydrogen are known, but only one, ammonia, need be considered here.
Preparation of ammonia. Ammonia is prepared in the laboratory by a different method from the one which is used commercially.
1. Laboratory method. In the laboratory ammonia is prepared from ammonium chloride, a compound having the formula NH{4}Cl, and obtained in the manufacture of coal gas. As will be shown later in the chapter, the group NH{4} in this compound acts as a univalent radical and is known as ammonium. When ammonium chloride is warmed with sodium hydroxide, the ammonium and sodium change places, the reaction being expressed in the following equation.
NH{4}Cl + NaOH = NaCl + NH{4}OH.
The ammonium hydroxide (NH_{4}OH) so formed is unstable and breaks down into water and ammonia.
NH_{4}OH = NH_{3} + H_{2}O.
Calcium hydroxide (Ca(OH)_{2}) is frequently used in place of the more expensive sodium hydroxide, the equations being
2NH{4}Cl + Ca(OH){2} = CaCl{2} + 2NH{4}OH,
2NH_{4}OH = 2H_{2}O + 2NH_{3}.
In the preparation, the ammonium chloride and calcium hydroxide are mixed together and placed in a flask arranged as shown in Fig. 35. The mixture is gently warmed, when ammonia is evolved as a gas and is collected by displacement of air.
2. Commercial method. Nearly all the ammonia of commerce comes from the gasworks. Ordinary illuminating gas is made by distilling coal, as will be explained later, and among the products of this distillation a solution of ammonia in water is obtained. This solution, known as gas liquor, contains not only ammonia but other soluble substances. Most of these combine chemically with lime, while ammonia does not; if then lime is added to the gas liquor and the liquor is heated, the ammonia is driven out from the mixture. It may be dissolved again in pure, cold water, forming aqua ammonia, or the ammonia water of commerce.
Preparation from hydrogen and nitrogen. When electric sparks are passed for some time through a mixture of hydrogen and nitrogen, a small percentage of the two elements in the mixture is changed into ammonia. The action soon ceases, however, for the reason that ammonia is decomposed by the electric discharge. The reaction expressed in the equation
N + 3H = NH_{3}
can therefore go in either direction depending upon the relative quantities of the substances present. This recalls the similar change from oxygen into ozone, which soon ceases because the ozone is in turn decomposed into oxygen.
Physical properties. Under ordinary conditions ammonia is a gas whose density is 0.59. It is therefore little more than half as heavy as air. It is easily condensed into a colorless liquid, and can now be purchased in liquid form in steel cylinders. The gas is colorless and has a strong, suffocating odor. It is extremely soluble in water, 1 l. of water at 0 deg. and 760 mm. pressure dissolving 1148 l. of the gas. In dissolving this large volume of gas the water expands considerably, so that the density of the solution is less than that of water, the strongest solutions having a density of 0.88.
Chemical properties. Ammonia will not support combustion, nor will it burn under ordinary conditions. In an atmosphere of oxygen it burns with a feeble, yellowish flame. When quite dry it is not a very active substance, but when moist it combines with a great many substances, particularly with acids.
Uses. It has been stated that ammonia can be condensed to a liquid by the application of pressure. If the pressure is removed from the liquid so obtained, it rapidly passes again into the gaseous state and in so doing absorbs a large amount of heat. Advantage is taken of this fact in the preparation of artificial ice. Large quantities of ammonia are also used in the preparation of ammonium compounds.
The manufacture of artificial ice. Fig. 36 illustrates the method of preparing artificial ice. The ammonia gas is liquefied in the pipes X by means of the pump Y. The heat generated is absorbed by water flowing over the pipes. The pipes lead into a large brine tank, a cross section of which is shown in the figure. Into the brine (concentrated solution of common salt) contained in this tank are dipped the vessels A, B, C, filled with pure water. The pressure is removed from the liquid ammonia as it passes into the pipes immersed in the brine, and the heat absorbed by the rapid evaporation of the liquid lowers the temperature of the brine below zero. The water in A, B, C is thereby frozen into cakes of ice. The gaseous ammonia resulting from the evaporation of the liquid ammonia is again condensed, so that the process is continuous.
Ammonium hydroxide (NH_{4}OH). The solution of ammonia in water is found to have strong basic properties and therefore contains hydroxyl ions. It turns red litmus blue; it has a soapy feel; it neutralizes acids, forming salts with them. It seems probable, therefore, that when ammonia dissolves in water it combines chemically with it according to the equation
NH_{3} + H_{2}O = NH_{4}OH,
and that it is the substance NH{4}OH, called ammonium hydroxide, which has the basic properties, dissociating into the ions NH{4} and OH. Ammonium hydroxide has never been obtained in a pure state. At every attempt to isolate it the substance breaks up into water and ammonia,—
NH_{4}OH = NH_{3} + H_{2}O.
The ammonium radical. The radical NH_{4} plays the part of a metal in many chemical reactions and is called ammonium. The ending _-ium_ is given to the name to indicate the metallic properties of the substance, since the names of the metals in general have that ending. The salts formed by the action of the base ammonium hydroxide on acids are called ammonium salts. Thus, with hydrochloric acid, ammonium chloride is formed in accordance with the equation
NH_{4}OH + HCl = NH_{4}Cl + H_{2}O.
Similarly, with nitric acid, ammonium nitrate (NH_{4}NO_{3}) is formed, and with sulphuric acid, ammonium sulphate ((NH_{4})_{2}S0_{4}).
It will be noticed that in the neutralization of ammonium hydroxide by acids the group NH_{4} replaces one hydrogen atom of the acid, just as sodium does. The group therefore acts as a univalent metal.
Combination of nitrogen with hydrogen by volume. Under suitable conditions ammonia can be decomposed into nitrogen and hydrogen by passing electric sparks through the gas. Accurate measurement has shown that when ammonia is decomposed, two volumes of the gas yield one volume of nitrogen and three volumes of hydrogen. Consequently, if the two elements were to combine directly, one volume of nitrogen would combine with three volumes of hydrogen to form two volumes of ammonia. Here, as in the formation of steam from hydrogen and oxygen, small whole numbers serve to indicate the relation between the volumes of combining gases and that of the gaseous product.
COMPOUNDS OF NITROGEN WITH OXYGEN AND HYDROGEN
In addition to ammonium hydroxide, nitrogen forms several compounds with hydrogen and oxygen, of which nitric acid (HNO{3}) and nitrous acid (HNO{2}) are the most familiar.
Nitric acid (HNO_{3}). Nitric acid is not found to any extent in nature, but some of its salts, especially sodium nitrate (NaNO_{3}) and potassium nitrate (KNO_{3}) are found in large quantities. From these salts nitric acid can be obtained.
Preparation of nitric acid. When sodium nitrate is treated with concentrated cold sulphuric acid, no chemical action seems to take place. If, however, the mixture is heated in a retort, nitric acid is given off as a vapor and may be easily condensed to a liquid by passing the vapor into a tube surrounded by cold water, as shown in Fig. 37. An examination of the liquid left in the retort shows that it contains sodium acid sulphate (NaHSO_{4}), so that the reaction may be represented by the equation
NaNO_{3} + H_{2}SO_{4} = NaHSO_{4} + HNO_{3}.
If a smaller quantity of sulphuric acid is taken and the mixture is heated to a high temperature, normal sodium sulphate is formed:
2NaNO{3} + H{2}SO{4} = Na{2}SO{4} + 2HNO{3}.
In this case, however, the higher temperature required decomposes a part of the nitric acid.
The commercial preparation of nitric acid. Fig. 38 illustrates a form of apparatus used in the preparation of nitric acid on a large scale. Sodium nitrate and sulphuric acid are heated in the iron retort A. The resulting acid vapors pass in the direction indicated by the arrows, and are condensed in the glass tubes B, which are covered with cloth kept cool by streams of water. These tubes are inclined so that the liquid resulting from the condensation of the vapors runs back into C and is drawn off into large vessels (D).
Physical properties of nitric acid. Pure nitric acid is a colorless liquid, which boils at about 86 deg. and has a density of 1.56. The concentrated acid of commerce contains about 68% of the acid, the remainder being water. Such a mixture has a density of 1.4. The concentrated acid fumes somewhat in moist air, and has a sharp choking odor.
Chemical properties. The most important chemical properties of nitric acid are the following.
1. Acid properties. As the name indicates, this substance is an acid, and has all the properties of that class of substances. It changes blue litmus red and has a sour taste in dilute solutions. It forms hydrogen ions in solution and neutralizes bases forming salts. It also acts upon the oxides of most metals, forming a salt and water. It is one of the strongest acids.
2. Decomposition on heating. When boiled, or exposed for some time to sunlight, it suffers a partial decomposition according to the equation
2HNO_{3} = H_{2}O + 2NO_{2} + O.
The substance NO_{2}, called nitrogen peroxide, is a brownish gas, which is readily soluble in water and in nitric acid. It therefore dissolves in the undecomposed acid, and imparts a yellowish or reddish color to it. Concentrated nitric acid highly charged with this substance is called _fuming nitric acid_.
3. Oxidizing action. According to its formula, nitric acid contains a large percentage of oxygen, and the reaction just mentioned shows that the compound is not a very stable one, easily undergoing decomposition. These properties should make it a good oxidizing agent, and we find that this is the case. Under ordinary circumstances, when acting as an oxidizing agent, it is decomposed according to the equation
2HNO{3} = H{2}O + 2NO + 3O.
The oxygen is taken up by the substance oxidized, and not set free, as is indicated in the equation. Thus, if carbon is oxidized by nitric acid, the oxygen combines with carbon, forming carbon dioxide (CO_{2}):
C + 2O = CO_{2}.
4. Action on metals. We have seen that when an acid acts upon a metal hydrogen is set free. Accordingly, when nitric acid acts upon a metal, such as copper, we should expect the reaction to take place which is expressed in the equation
Cu + 2HNO_{3} = Cu(NO_{3})_{2} + 2H.
This reaction does take place, but the hydrogen set free is immediately oxidized to water by another portion of the nitric acid according to the equation
HNO{3} + 3H = 2H{2}O + NO.
As these two equations are written, two atoms of hydrogen are given off in the first equation, while three are used up in the second. In order that the hydrogen may be equal in the two equations, we must multiply the first by 3 and the second by 2. We shall then have
3Cu + 6HNO_{3} = 3Cu(NO_{3})_{2} + 6H,
2HNO{3} + 6H = 4H{2}O + 2NO.
The two equations may now be combined into one by adding the quantities on each side of the equality sign, canceling the hydrogen which is given off in the one reaction and used up in the other. We shall then have the equation
3Cu + 8HNO{3} = 3Cu(NO{3}){2} + 2NO + 4H{2}O.
A number of other reactions may take place when nitric acid acts upon metals, resulting in the formation of other oxides of nitrogen, free nitrogen, or even ammonia. The reaction just given is, however, the usual one.
Importance of steps in a reaction. This complete equation has the advantage of making it possible to calculate very easily the proportions in which the various substances enter into the reaction or are formed in it. It is unsatisfactory in that it does not give full information about the way in which the reaction takes place. For example, it does not suggest that hydrogen is at first formed, and subsequently transformed into water. It is always much more important to remember the steps in a chemical reaction than to remember the equation expressing the complete action; for if these steps in the reaction are understood, the complete equation is easily obtained in the manner just described.
Salts of nitric acid,—nitrates. The salts of nitric acid are called nitrates. Many of these salts will be described in the study of the metals. They are all soluble in water, and when heated to a high temperature undergo decomposition. In a few cases a nitrate on being heated evolves oxygen, forming a nitrite:
NaNO{3} = NaNO{2} + O.
In other cases the decomposition goes further, and the metal is left as oxide:
Cu(NO_{3})_{2} = CuO + 2NO_{2} + O.
Nitrous acid (HNO{2}). It is an easy matter to obtain sodium nitrite (NaNO{2}), as the reaction given on the previous page indicates. Instead of merely heating the nitrate, it is better to heat it together with a mild reducing agent, such as lead, when the reaction takes place which is expressed by the equation
NaNO{3} + Pb = PbO + NaNO{2}.
When sodium nitrite is treated with an acid, such as sulphuric acid, it is decomposed and nitrous acid is set free:
NaNO_{2} + H_{2}SO_{4} = NaHSO_{4} + HNO_{2}.
The acid is very unstable, however, and decomposes readily into water and nitrogen trioxide (N{2}O{3}):
2HNO{2} = H{2}O + N{2}O{3}.
Dilute solutions of the acid, however, can be obtained.
COMPOUNDS OF NITROGEN WITH OXYGEN
Nitrogen combines with oxygen to form five different oxides. The formulas and names of these are as follows:
N{2}O nitrous oxide. NO nitric oxide. NO{2} nitrogen peroxide. N{2}O{3} nitrogen trioxide, or nitrous anhydride. N{2}O{5} nitrogen pentoxide, or nitric anhydride.
These will now be briefly discussed.
Nitrous oxide (_laughing gas_) (N_{2}O). Ammonium nitrate, like all nitrates, undergoes decomposition when heated; and owing to the fact that it contains no metal, but does contain both oxygen and hydrogen, the reaction is a peculiar one. It is represented by the equation
NH{4}NO{3} = 2H{2}O + N{2}O.
The oxide of nitrogen so formed is called nitrous oxide or laughing gas. It is a colorless gas having a slight odor. It is somewhat soluble in water, and in solution has a slightly sweetish taste. It is easily converted into a liquid and can be purchased in this form. When inhaled it produces a kind of hysteria (hence the name "laughing gas"), and even unconsciousness and insensibility to pain if taken in large amounts. It has long been used as an anaesthetic for minor surgical operations, such as those of dentistry, but owing to its unpleasant after effects it is not so much in use now as formerly.
Chemically, nitrous oxide is remarkable for the fact that it is a very energetic oxidizing agent. Substances such as carbon, sulphur, iron, and phosphorus burn in it almost as brilliantly as in oxygen, forming oxides and setting free nitrogen. Evidently the oxygen in nitrous oxide cannot be held in very firm combination by the nitrogen.
Nitric oxide (NO). We have seen that when nitric acid acts upon metals, such as copper, the reaction represented by the following equation takes place:
3Cu + 8HNO{3} = 3Cu(NO{3}){3} + 2NO + 4H{2}O.
Nitric oxide is most conveniently prepared in this way. The metal is placed in the flask A (Fig. 39) and the acid added slowly through the funnel tube B. The gas escapes through C and is collected over water.
Pure nitric oxide is a colorless gas, slightly heavier than air, and is practically insoluble in water. It is a difficult gas to liquefy. Unlike nitrous oxide, nitric oxide does not part with its oxygen easily, and burning substances introduced into this gas are usually extinguished. A few substances like phosphorus, which have a very strong affinity for oxygen and which are burning energetically in the air, will continue to burn in an atmosphere of nitric oxide. In this case the nitric oxide loses all of its oxygen and the nitrogen is set free as gas.
Action of nitric oxide with oxygen. When nitric oxide comes into contact with oxygen or with the air, it at once combines with the oxygen even at ordinary temperatures, forming a reddish-yellow gas of the formula NO_{2}, which is called nitrogen peroxide. This action is not energetic enough to produce a flame, though considerable heat is set free.
Nitrogen peroxide (NO_{2}). This gas, as we have just seen, is formed by allowing nitric oxide to come into contact with oxygen. It can also be made by heating certain nitrates, such as lead nitrate:
Pb(NO_{3})_{2} = PbO + 2NO_{2} + O.
It is a reddish-yellow gas of unpleasant odor, which is quite poisonous when inhaled. It is heavier than air and is easily condensed to a liquid. It dissolves in water, but this solution is not a mere physical solution; the nitrogen peroxide is decomposed, forming a mixture of nitric and nitrous acids:
2NO{2} + H{2}O = HNO{2} + HNO{3}.
Nitrogen peroxide will not combine with more oxygen; it will, however, give up a part of its oxygen to burning substances, acting as an oxidizing agent:
NO_{2} = NO + O.
Acid anhydrides. The oxides N{2}O{3} (nitrogen trioxide) and N{2}O{5} (nitrogen pentoxide) are rarely prepared and need not be separately described. They bear a very interesting relation to the acids of nitrogen. When dissolved in water they combine with the water, forming acids:
N{2}O{3} + H{2}O = 2HNO{2},
N{2}O{5} + H{2}O = 2HNO{3}.
On the other hand, nitrous acid very easily decomposes, yielding water and nitrogen trioxide, and by suitable means nitric acid likewise may be decomposed into water and nitrogen pentoxide:
2HNO{2} = H{2}O + N{2}O{3},
2HNO{3} = H{2}O + N{2}O{5}.
In view of the close relation between these oxides and the corresponding acids, they are called anhydrides of the acids, N{2}O{3} being nitrous anhydride and N{2}O{5} nitric anhydride.
DEFINITION: Any oxide which will combine with water to form an acid, or which together with water is formed by the decomposition of an acid, is called an anhydride of that acid.
EXERCISES
1. Perfectly dry ammonia does not affect litmus paper. Explain.
2. Can ammonia be dried by passing the gas through concentrated sulphuric acid? Explain.
3. Ammonium hydroxide is a weak base, i.e. it is not highly dissociated. When it is neutralized by strong acids the heat of reaction is less than when strong bases are so neutralized. Suggest some possible cause for this.
4. Why is brine used in the manufacture of artificial ice?
5. Discuss the energy changes which take place in the manufacture of artificial ice.
6. What weight of ammonium chloride is necessary to furnish enough ammonia to saturate 1 l. of water at 0 deg. and 760 mm.?
7. What weight of sodium nitrate is necessary to prepare 100 cc. of commercial nitric acid? What weight of potassium nitrate is necessary to furnish the same weight of acid?
8. 100 l. of nitrogen peroxide were dissolved in water and neutralized with sodium hydroxide. What substances were formed and how much of each?(1 l. nitrogen peroxide weighs 2.05 grams.)
9. How many liters of nitrous oxide, measured under standard conditions, can be prepared from 10 g. of ammonium nitrate?
10. What weight of copper is necessary to prepare 50 l. of nitric oxide under standard conditions?
11. (a) Calculate the percentage composition of the oxides of nitrogen. (b) What important law does this series of substances illustrate?
12. Write the equations representing the reactions between ammonium hydroxide, and sulphuric acid and nitric acid respectively, in accordance with the theory of electrolytic dissociation.
13. In the same way, write the equations representing the reactions between nitric acid and each of the following bases: NaOH, KOH, NH{4}OH, Ca(OH){2}.
CHAPTER XIII
REVERSIBLE REACTIONS AND CHEMICAL EQUILIBRIUM
Reversible reactions. The reactions so far considered have been represented as continuing, when once started, until one or the other substance taking part in the reaction has been used up. In some reactions this is not the case. For example, we have seen that when steam is passed over hot iron the reaction is represented by the equation
3Fe + 4H_{2}O = Fe_{3}O_{4} + 8H.
On the other hand, when hydrogen is passed over hot iron oxide the reverse reaction takes place:
Fe_{3}O_{4} +8H = 3Fe + 4H_{2}O.
The reaction can therefore go in either direction, depending upon the conditions of the experiment. Such a reaction is called a reversible reaction. It is represented by an equation with double arrows in place of the equality sign, thus:
3Fe + 4H_{2}O Fe_{3}O_{4} + 8H.
In a similar way, the equation
N + 3H NH_{3}
expresses the fact that under some conditions nitrogen may unite with hydrogen to form ammonia, while under other conditions ammonia decomposes into nitrogen and hydrogen.
The conversion of oxygen into ozone is also reversible and may be represented thus:
oxygen ozone.
Chemical equilibrium. Reversible reactions do not usually go on to completion in one direction unless the conditions under which the reaction takes place are very carefully chosen. Thus, if iron and steam are confined in a heated tube, the steam acts upon the iron, producing iron oxide and hydrogen. But these substances in turn act upon each other to form iron and steam once more. When these two opposite reactions go on at such rates that the weight of the iron changed into iron oxide is just balanced by the weight of the iron oxide changed into iron, there will be no further change in the relative weights of the four substances present in the tube. The reaction is then said to have reached an equilibrium.
Factors which determine the point of equilibrium. There are two factors which have a great deal of influence in determining the point at which a given reaction will reach equilibrium.
1. Influence of the chemical nature of the substances. If two reversible reactions of the same general kind are selected, it has been found that the point of equilibrium is different in the two cases. For example, in the reactions represented by the equations
3Fe + 4H_{2}O Fe_{3}O_{4} + 8H,
Zn + H_{2}O ZnO + 2H,
the equilibrium will be reached when very different quantities of the iron and zinc have been changed into oxides. The individual chemical properties of the iron and zinc have therefore marked influence upon the point at which equilibrium will be reached.
2. Influence of relative mass. If the tube in which the reaction
3Fe + 4H_{2}O Fe_{3}O_{4} + 8H
has come to an equilibrium is opened and more steam is admitted, an additional quantity of the iron will be changed into iron oxide. If more hydrogen is admitted, some of the oxide will be reduced to metal. The point of equilibrium is therefore dependent upon the relative masses of the substances taking part in the reaction. When one of the substances is a solid, however, its mass has little influence, since it is only the extent of its surface which can affect the reaction.
Conditions under which reversible reactions are complete. If, when the equilibrium between iron and steam has been reached, the tube is opened and a current of steam is passed in, the hydrogen is swept away as fast as it is formed. The opposing reaction of hydrogen upon iron oxide must therefore cease, and the action of steam on the iron will go on until all of the iron has been transformed into iron oxide.
On the other hand, if a current of hydrogen is admitted into the tube, the steam will be swept away by the hydrogen, and all of the iron oxide will be reduced to iron. A reversible reaction can therefore be completed in either direction when one of the products of the reaction is removed as fast as it is formed.
Equilibrium in solution. When reactions take place in solution in water the same general principles hold good. The matter is not so simple, however, as in the case just described, owing to the fact that many of the reactions in solution are due to the presence of ions. The substances most commonly employed in solution are acids, bases, or salts, and all of these undergo dissociation. Any equilibrium which may be reached in solutions of these substances must take place between the various ions formed, on the one hand, and the undissociated molecules, on the other. Thus, when nitric acid is dissolved in water, equilibrium is reached in accordance with the equation
H^{+} + NO{3}^{-} HNO{3}.
Conditions under which reversible reactions in solution are complete. The equilibrium between substances in solution may be disturbed and the reaction caused to go on in one direction to completion in either of three ways.
1. _A gas may be formed which escapes from the solution._ When sodium nitrate and sulphuric acid are brought together in solution all four ions, Na^{+}, NO_{3}^{-}, H^{+}, SO_{4}^{—}, are formed. These ions are free to rearrange themselves in various combinations. For example, the H^{+} and the NO_{3}^{-} ions will reach the equilibrium
H^{+} + NO{3}^{-} HNO{3}.
If the experiment is performed with very little water present, as is the case in the preparation of nitric acid, the equilibrium will be reached when most of the H^{+} and the NO{3}^{-} ions have combined to form undissociated HNO{3}.
Finally, if the mixture is now heated above the boiling point of nitric acid, the acid distills away as fast as it is formed. More and more H^{+} and NO{3}^{-} ions will then combine, and the process will continue until one or the other of them has all been removed from the solution. The substance remaining is sodium acid sulphate (NaHSO{4}), and the reaction can therefore be expressed by the equation
NaNO_{3} + H_{2}SO_{4} = NaHSO_{4} + HNO_{3}.
2. An insoluble solid may be formed. When hydrochloric acid (HCl) and silver nitrate (AgNO{3}) are brought together in solution the following ions will be present: H^{+}, Cl^{-}, Ag^{+}, NO{3}^{-}. The ions Ag^{+} and Cl^{-} will then set up the equilibrium
Ag^{+} + Cl^{-} AgCl.
But silver chloride (AgCl) is almost completely insoluble in water, and as soon as a very little of it has formed the solution becomes supersaturated, and the excess of the salt precipitates. More silver and chlorine ions then unite, and this continues until practically all of the silver or the chlorine ions have been removed from the solution. We then say that the following reaction is complete:
AgNO{3} + HCl = AgCl + HNO{3}.
3. Two different ions may form undissociated molecules. In the neutralization of sodium hydroxide by hydrochloric acid the ions H^{+} and OH^{-} come to the equilibrium
H^{+} + OH^{-} H_{2}O.
But since water is almost entirely undissociated, equilibrium can only be reached when there are very few hydroxyl or hydrogen ions present. Consequently the two ions keep uniting until one or the other of them is practically removed from the solution. When this occurs the neutralization expressed in the following equation is complete:
NaOH + HCl = H_{2}O + NaCl.
Preparation of acids. The principle of reversible reactions finds practical application in the preparation of most of the common acids. An acid is usually prepared by treating the most common of its salts with some other acid of high boiling point. The mixture is then heated until the lower boiling acid desired distills out. Owing to its high boiling point (338 deg.), sulphuric acid is usually employed for this purpose, most other acids boiling below that temperature.
EXERCISES
1. What would take place when solutions of silver nitrate and sodium chloride are brought together? What other chlorides would act in the same way?
2. Is the reaction expressed by the equation NH_{3} + H_{2}O = NH_{4}OH reversible? If so, state the conditions under which it will go in each direction.
3. Is the reaction expressed by the equation 2H + O = H_{2}O reversible? If so, state the conditions under which it will go in each direction.
4. Suggest a method for the preparation of hydrochloric acid.
CHAPTER XIV
SULPHUR AND ITS COMPOUNDS
Occurrence. The element sulphur has been known from the earliest times, since it is widely distributed in nature and occurs in large quantities in the uncombined form, especially in the neighborhood of volcanoes. Sicily has long been famous for its sulphur mines, and smaller deposits are found in Italy, Iceland, Mexico, and especially in Louisiana, where it is mined extensively. In combination, sulphur occurs abundantly in the form of sulphides and sulphates. In smaller amounts it is found in a great variety of minerals, and it is a constituent of many animal and vegetable substances.
Extraction of sulphur. Sulphur is prepared from the native substance, the separation of crude sulphur from the rock and earthy materials with which it is mixed being a very simple process. The ore from the mines is merely heated until the sulphur melts and drains away from the earthy impurities. The crude sulphur obtained in this way is distilled in a retort-shaped vessel made of iron, the exit tube of which opens into a cooling chamber of brickwork. When the sulphur vapor first enters the cooling chamber it condenses as a fine crystalline powder called flowers of sulphur. As the condensing chamber becomes warm, the sulphur collects as a liquid in it, and is drawn off into cylindrical molds, the product being called roll sulphur or brimstone.
Physical properties. Roll sulphur is a pale yellow, crystalline solid, without marked taste and with but a faint odor. It is insoluble in water, but is freely soluble in a few liquids, notably in carbon disulphide. Roll sulphur melts at 114.8 deg.. Just above the melting point it forms a rather thin, straw-colored liquid. As the temperature is raised, this liquid turns darker in color and becomes thicker, until at about 235 deg. it is almost black and is so thick that the vessel containing it can be inverted without danger of the liquid running out. At higher temperatures it becomes thin once more, and boils at 448 deg., forming a yellowish vapor. On cooling the same changes take place in reverse order.
Varieties of sulphur. Sulphur is known in two general forms, crystalline and amorphous. Each of these forms exists in definite modifications.
Crystalline sulphur. Sulphur occurs in two crystalline forms, namely, rhombic sulphur and monoclinic sulphur.
1. Rhombic sulphur. When sulphur crystallizes from its solution in carbon disulphide it separates in crystals which have the same color and melting point as roll sulphur, and are rhombic in shape. Roll sulphur is made up of minute rhombic crystals.
2. Monoclinic sulphur. When melted sulphur is allowed to cool until a part of the liquid has solidified, and the remaining liquid is then poured off, it is found that the solid sulphur remaining in the vessel has assumed the form of fine needle-shaped crystals. These differ much in appearance from the rhombic crystals obtained by crystallizing sulphur from its solution in carbon disulphide. The needle-shaped form is called monoclinic sulphur. The two varieties differ also in density and in melting point, the monoclinic sulphur melting at 120 deg..
Monoclinic and rhombic sulphur remain unchanged in contact with each other at 96 deg.. Above this temperature the rhombic changes into monoclinic; at lower temperatures the monoclinic changes into rhombic. The temperature 96 deg. is therefore called the transition point of sulphur. Heat is set free when monoclinic sulphur changes into rhombic.
Amorphous sulphur. Two varieties of amorphous sulphur can be readily obtained. These are white sulphur and plastic sulphur.
1. White sulphur. Flowers of sulphur, the preparation of which has been described, consists of a mixture of rhombic crystals and amorphous particles. When treated with carbon disulphide, the crystals dissolve, leaving the amorphous particles as a white residue.
2. Plastic sulphur. When boiling sulphur is poured into cold water it assumes a gummy, doughlike form, which is quite elastic. This can be seen in a very striking manner by distilling sulphur from a small, short-necked retort, such as is represented in Fig. 40, and allowing the liquid to run directly into water. In a few days it becomes quite brittle and passes over into ordinary rhombic sulphur.
Chemical properties of sulphur. When sulphur is heated to its kindling temperature in oxygen or in the air it burns with a pale blue flame, forming sulphur dioxide (SO{2}). Small quantities of sulphur trioxide (SO{3}) may also be formed in the combustion of sulphur. Most metals when heated with sulphur combine directly with it, forming metallic sulphides. In some cases the action is so energetic that the mass becomes incandescent, as has been seen in the case of iron uniting with sulphur. This property recalls the action of oxygen upon metals, and in general the metals which combine readily with oxygen are apt to combine quite readily with sulphur.
Uses of sulphur. Large quantities of sulphur are used as a germicide in vineyards, also in the manufacture of gunpowder, matches, vulcanized rubber, and sulphuric acid.
COMPOUNDS OF SULPHUR WITH HYDROGEN
Hydrosulphuric acid (H{2}S). This substance is a gas having the composition expressed by the formula H{2}S and is commonly called hydrogen sulphide. It is found in the vapors issuing from volcanoes, and in solution in the so-called sulphur waters of many springs. It is formed when organic matter containing sulphur undergoes decay, just as ammonia is formed under similar circumstances from nitrogenous matter.
Preparation. Hydrosulphuric acid is prepared in the laboratory by treating a sulphide with an acid. Iron sulphide (FeS) is usually employed:
FeS + 2HCl = FeCl{2} + H{2}S.
A convenient apparatus is shown in Fig. 41. A few lumps of iron sulphide are placed in the bottle A, and dilute acid is added in small quantities at a time through the funnel tube B, the gas escaping through the tube C.
Explanation of the reaction. Iron sulphide is a salt of hydrosulphuric acid, and this reaction is therefore similar to the one which takes place when sulphuric acid acts upon a nitrate. In both cases a salt and an acid are brought together, and there is a tendency for the reaction to go on until a state of equilibrium is reached. This equilibrium is constantly disturbed by the escape of the gaseous acid set free, so that the reaction goes on until all of the original salt has been decomposed. The two reactions differ in that the first one is complete at ordinary temperatures, while in the case of sulphuric acid acting upon sodium nitrate, the reacting substances must be heated so as to secure a temperature at which nitric acid is a gas.
Physical properties. Hydrosulphuric acid is a colorless gas, having a weak, disagreeable taste and an exceedingly offensive odor. It is rather sparingly soluble in water at ordinary temperatures, about three volumes dissolving in one of water. In boiling water it is not soluble at all. In pure form it acts as a violent poison, and even when diluted largely with air produces headache, dizziness, and nausea. It is a little heavier than air, having a density of 1.18.
Chemical properties. The most important chemical properties of hydrosulphuric acid are the following:
1. Acid properties. Hydrosulphuric acid is a weak acid. In solution in water it turns blue litmus red and neutralizes bases, forming salts called sulphides.
2. Action on oxygen. The elements composing hydrosulphuric acid have each a strong affinity for oxygen, and are not held together very firmly. Consequently the gas burns readily in oxygen or the air, according to the equation
H_{2}S + 3O = H_{2}O + SO_{2}.
When there is not enough oxygen for both the sulphur and the hydrogen, the latter element combines with the oxygen and the sulphur is set free:
H{2}S + O = H{2}O + S.
3. Reducing action. Owing to the ease with which hydrosulphuric acid decomposes and the strong affinity of both sulphur and hydrogen for oxygen, the substance is a strong reducing agent, taking oxygen away from many substances which contain it.
4. Action on metals. Hydrosulphuric acid acts towards metals in a way very similar to water. Thus, when it is passed over heated iron in a tube, the reaction is represented by the equation
3Fe + 4H_{2}S = Fe_{3}S_{4} + 8H.
Water in the form of steam, under similar circumstances, acts according to the equation
3Fe + 4H_{2}O = Fe_{3}O_{4} + 8H.
Salts of hydrosulphuric acid,—sulphides. The salts of hydrosulphuric acid, called sulphides, form an important class of salts. Many of them are found abundantly in nature, and some of them are important ores. They will be frequently mentioned in connection with the metals.
Most of the sulphides are insoluble in water, and some of them are insoluble in acids. Consequently, when hydrosulphuric acid is passed into a solution of a salt, it often happens that a sulphide is precipitated. With copper chloride the equation is
CuCl{2} + H{2}S = CuS + 2HCl.
Because of the fact that some metals are precipitated in this way as sulphides while others are not, hydrosulphuric acid is extensively used in the separation of the metals in the laboratory.
Explanation of the reaction. When hydrosulphuric acid and copper chloride are brought together in solution, both copper and sulphur ions are present, and these will come to an equilibrium, as represented in the equation
Cu^{+} + S^{-} CuS.
Since copper sulphide is almost insoluble in water, as soon as a very small quantity has formed the solution becomes supersaturated, and the excess keeps precipitating until nearly all the copper or sulphur ions have been removed from the solution. With some other ions, such as iron, the sulphide formed does not saturate the solution, and no precipitate results.
OXIDES OF SULPHUR
Sulphur forms two well-known compounds with oxygen: sulphur dioxide (SO{2}), sometimes called sulphurous anhydride; and sulphur trioxide (SO{3}), frequently called sulphuric anhydride.
Sulphur dioxide (SO_{2}). Sulphur dioxide occurs in nature in the gases issuing from volcanoes, and in solution in the water of many springs. It is likely to be found wherever sulphur compounds are undergoing oxidation.
Preparation. Three general ways may be mentioned for the preparation of sulphur dioxide:
1. By the combustion of sulphur. Sulphur dioxide is readily formed by the combustion of sulphur in oxygen or the air:
S + 2O = SO_{2}.
It is also formed when substances containing sulphur are burned:
ZnS + 3O = ZnO + SO_{2}.
2. By the reduction of sulphuric acid. When concentrated sulphuric acid is heated with certain metals, such as copper, part of the acid is changed into copper sulphate, and part is reduced to sulphurous acid. The latter then decomposes into sulphur dioxide and water, the complete equation being
Cu + 2H_{2}SO_{4} = CuSO_{4} + SO_{2} + 2H_{2}O.
3. By the action of an acid on a sulphite. Sulphites are salts of sulphurous acid (H{2}SO{3}). When a sulphite is treated with an acid, sulphurous acid is set free, and being very unstable, decomposes into water and sulphur dioxide. These reactions are expressed in the equations
Na{2}SO{3} + 2HCl = 2NaCl + H{2}SO{3},
H{2}SO{3} = H{2}O + SO{2}.
Explanation of the reaction. In this case we have two reversible reactions depending on each other. In the first reaction,
(1) Na{2}SO{3} + 2HCl 2NaCl + H{2}SO{3},
we should expect an equilibrium to result, for none of the four substances in the equation are insoluble or volatile when water is present to hold them in solution. But the quantity of the H{2}SO{3} is constantly diminishing, owing to the fact that it decomposes, as represented in the equation
(2) H{2}SO{3} H{2}O + SO{2},
and the sulphur dioxide, being a gas, escapes. No equilibrium can therefore result, since the quantity of the sulphurous acid is constantly being diminished because of the escape of sulphur dioxide.
Physical properties. Sulphur dioxide is a colorless gas, which at ordinary temperatures is 2.2 times as heavy as air. It has a peculiar, irritating odor. The gas is very soluble in water, one volume of water dissolving eighty of the gas under standard conditions. It is easily condensed to a colorless liquid, and can be purchased in this condition stored in strong bottles, such as the one represented in Fig. 42.
Chemical properties. Sulphur dioxide has a marked tendency to combine with other substances, and is therefore an active substance chemically. It combines with oxygen gas, but not very easily. It can, however, take oxygen away from some other substances, and is therefore a good reducing agent. Its most marked chemical property is its ability to combine with water to form sulphurous acid (H{2}SO{3}).
Sulphurous acid (H{2}SO{3}). When sulphur dioxide dissolves in water it combines chemically with it to form sulphurous acid, an unstable substance having the formula H{3}SO{3}. It is impossible to prepare this acid in pure form, as it breaks down very easily into water and sulphur dioxide. The reaction is therefore reversible, and is expressed by the equation
H{2}O + SO{2} H{2}SO{3}.
Solutions of the acid in water have a number of interesting properties.
1. Acid properties. The solution has all the properties typical of an acid. When neutralized by bases, sulphurous acid yields a series of salts called sulphites.
2. Reducing properties. Solutions of sulphurous acid act as good reducing agents. This is due to the fact that sulphurous acid has the power of taking up oxygen from the air, or from substances rich in oxygen, and is changed by this reaction into sulphuric acid:
H{2}SO{3} + O = H{2}SO{4},
H_{2}SO_{3} + H_{2}O_{2} = H_{2}S0_{4} + H_{2}O.
3. Bleaching properties. Sulphurous acid has strong bleaching properties, acting upon many colored substances in such a way as to destroy their color. It is on this account used to bleach paper, straw goods, and even such foods as canned corn.
4. Antiseptic properties. Sulphurous acid has marked antiseptic properties, and on this account has the power of arresting fermentation. It is therefore used as a preservative.
Salts of sulphurous acid,—sulphites. The sulphites, like sulphurous acid, have the power of taking up oxygen very readily, and are good reducing agents. On account of this tendency, commercial sulphites are often contaminated with sulphates. A great deal of sodium sulphite is used in the bleaching industry, and as a reagent for softening paper pulp.
Sulphur trioxide (SO{3}). When sulphur dioxide and oxygen are heated together at a rather high temperature, a small amount of sulphur trioxide (SO{3}) is formed, but the reaction is slow and incomplete. If, however, the heating takes place in the presence of very fine platinum dust, the reaction is rapid and nearly complete.
~ Experimental preparation of sulphur trioxide.~ The experiment can be performed by the use of the apparatus shown in Fig. 43, the fine platinum being secured by moistening asbestos fiber with a solution of platinum chloride and igniting it in a flame. The fiber, covered with fine platinum, is placed in a tube of hard glass, which is then heated with a burner to about 350 deg., while sulphur dioxide and air are passed into the tube. Union takes place at once, and the strongly fuming sulphur trioxide escapes from the jet at the end of the tube, and may be condensed by surrounding the receiving tube with a freezing mixture.
Properties of sulphur trioxide. Sulphur trioxide is a colorless liquid, which solidifies at about 15 deg. and boils at 46 deg.. A trace of moisture causes it to solidify into a mass of silky white crystals, somewhat resembling asbestos fiber in appearance. In contact with the air it fumes strongly, and when thrown upon water it dissolves with a hissing sound and the liberation of a great deal of heat. The product of this reaction is sulphuric acid, so that sulphur trioxide is the anhydride of that acid:
SO{3} + H{2}O = H{2}SO{4}.
Catalysis. It has been found that many chemical reactions, such as the union of sulphur dioxide with oxygen, are much influenced by the presence of substances which do not themselves seem to take a part in the reaction, and are left apparently unchanged after it has ceased. These reactions go on very slowly under ordinary circumstances, but are greatly hastened by the presence of the foreign substance. Substances which hasten very slow reactions in this way are said to act as catalytic agents or catalyzers, and the action is called catalysis. Just how the action is brought about is not well understood.
DEFINITION: A catalyzer is a substance which changes the velocity of a reaction, but does not change its products.
Examples of Catalysis. We have already had several instances of such action. Oxygen and hydrogen combine with each other at ordinary temperatures in the presence of platinum powder, while if no catalytic agent is present they do not combine in appreciable quantities until a rather high temperature is reached. Potassium chlorate, when heated with manganese dioxide, gives up its oxygen at a much lower temperature than when heated alone. Hydrogen dioxide decomposes very rapidly when powdered manganese dioxide is sifted into its concentrated solution.
On the other hand, the catalytic agent sometimes retards chemical action. For example, a solution of hydrogen dioxide decomposes more slowly when it contains a little phosphoric acid than when perfectly pure. For this reason commercial hydrogen dioxide always contains phosphoric acid.
Many reactions are brought about by the catalytic action of traces of water. For example, phosphorus will not burn in oxygen in the absence of all moisture. Hydrochloric acid will not unite with ammonia if the reagents are perfectly dry. It is probable that many of the chemical transformations in physiological processes, such as digestion, are assisted by certain substances acting as catalytic agents. The principle of catalysis is therefore very important.
Sulphuric acid (oil of vitriol) (H{2}SO{4}). Sulphuric acid is one of the most important of all manufactured chemicals. Not only is it one of the most common reagents in the laboratory, but enormous quantities of it are used in many of the industries, especially in the refining of petroleum, the manufacture of nitroglycerin, sodium carbonate, and fertilizers.
Manufacture of sulphuric acid. 1. Contact process. The reactions taking place in this process are represented by the following equations:
SO{2} + O = SO{3},
SO{3} + H{2}O = H{2}SO{4}.
To bring about the first of these reactions rapidly, a catalyzer is employed, and the process is carried out in the following way: Large iron tubes are packed with some porous material, such as calcium and magnesium sulphates, which contains a suitable catalytic substance scattered through it. The catalyzers most used are platinum powder, vanadium oxide, and iron oxide. Purified sulphur dioxide and air are passed through the tubes, which are kept at a temperature of about 350 deg.. Sulphur trioxide is formed, and as it issues from the tube it is absorbed in water or dilute sulphuric acid. The process is continued until all the water in the absorbing vessel has been changed into sulphuric acid, so that a very concentrated acid is made in this way. An excess of the trioxide may dissolve in the strong sulphuric acid, forming what is known as fuming sulphuric acid.
2. Chamber process. The method of manufacture exclusively employed until recent years, and still in very extensive use, is much more complicated. The reactions are quite involved, but the conversion of water, sulphur dioxide, and oxygen into sulphuric acid is accomplished by the catalytic action of oxides of nitrogen. The reactions are brought about in large lead-lined chambers, into which oxides of nitrogen, sulphur dioxide, steam, and air are introduced in suitable proportions.
Reactions of the chamber process. In a very general way, the various reactions which take place in the lead chambers may be expressed in two equations. In the first reaction sulphur dioxide, nitrogen peroxide, steam, and oxygen unite, as shown in the equation
(1) 2SO_{2} + 2NO_{2} + H_{2}O + O = 2SO_{2} (OH) (NO_{2}).
The product formed in this reaction is called nitrosulphuric acid or "chamber crystals." It actually separates on the walls of the chambers when the process is not working properly. Under normal conditions, it is decomposed as fast as it is formed by the action of excess of steam, as shown in the equation
(2) 2SO{2} (OH) (NO{2}) + H{2}O + O = 2H{2}SO{4} + 2NO{2}.
The nitrogen dioxide formed in this reaction can now enter into combination with a new quantity of sulphur dioxide, steam, and oxygen, and the series of reactions go on indefinitely. Many other reactions occur, but these two illustrate the principle of the process.
The relation between sulphuric acid and nitrosulphuric acid can be seen by comparing their structural formulas:
O= -OH O= -OH S S O= -OH O= -NO_{2}
The latter may be regarded as derived from the former by the substitution of the nitro group (NO_{2}) for the hydroxyl group (OH).
The sulphuric acid plant. Fig. 44 illustrates the simpler parts of a plant used in the manufacture of sulphuric acid by the chamber process. Sulphur or some sulphide, as FeS_{2}, is burned in furnace A. The resulting sulphur dioxide, together with air and some nitrogen peroxide, are conducted into the large chambers, the capacity of each chamber being about 75,000 cu. ft. Steam is also admitted into these chambers at different points. These compounds react to form sulphuric acid, according to the equations given above. The nitrogen left after the withdrawal of the oxygen from the admitted air escapes through the Gay-Lussac tower X. In order to prevent the escape of the oxides of nitrogen regenerated in the reaction, the tower is filled with lumps of coke, over which trickles concentrated sulphuric acid admitted from Y. The nitrogen peroxide dissolves in the acid and the resulting solution collects in H. This is pumped into E, where it is mixed with dilute acid and allowed to trickle down through the chamber D (Glover tower), which is filled with some acid-resisting rock. Here the nitrogen peroxide is expelled from the solution by the action of the hot gases entering from A, and together with them enters the first chamber again. The acid from which the nitrogen peroxide is expelled collects in F. Theoretically, a small amount of nitrogen peroxide would suffice to prepare an unlimited amount of sulphuric acid; practically, some of it escapes, and this is replaced by small amounts admitted at B.
The sulphuric acid so formed, together with the excess of condensed steam, collect upon the floor of the chambers in the form of a liquid containing from 62% to 70% of sulphuric acid. The product is called chamber acid and is quite impure; but for many purposes, such as the manufacture of fertilizers, it needs no further treatment. It can be concentrated by boiling it in vessels made of iron or platinum, which resist the action of the acid, nearly all the water boiling off. Pure concentrated acid can be made best by the contact process, while the chamber process is cheaper for the dilute impure acid.
Physical properties. Sulphuric acid is a colorless, oily liquid, nearly twice as heavy as water. The ordinary concentrated acid contains about 2% of water, has a density of 1.84, and boils at 338 deg.. It is sometimes called oil of vitriol, since it was formerly made by distilling a substance called green vitriol.
Chemical properties. Sulphuric acid possesses chemical properties which make it one of the most important of chemical substances.
1. Action as an acid. In dilute solution sulphuric acid acts as any other acid, forming salts with oxides and hydroxides.
2. Action as an oxidizing agent. Sulphuric acid contains a large percentage of oxygen and is, like nitric acid, a very good oxidizing agent. When the concentrated acid is heated with sulphur, carbon, and many other substances, oxidation takes place, the sulphuric acid decomposing according to the equation
H{2}SO{4} = H{2}SO{3} + O.
3. Action on metals. In dilute solution sulphuric acid acts upon many metals, such as zinc, forming a sulphate and liberating hydrogen. When the concentrated acid is employed the hydrogen set free is oxidized by a new portion of the acid, with the liberation of sulphur dioxide. With copper the reactions are expressed by the equations
(1) Cu + H_{2}SO_{4} = CuSO_{4} + 2H,
(2) H_{2}SO_{4} + 2H = H_{2}SO_{3} + H_{2}O,
(3) H{2}SO{3} = H{2}O + SO{2}.
By combining these equations the following one is obtained:
Cu + 2H_{2}SO_{4} = CuSO_{4} + SO_{2} + 2H_{2}O.
4. Action on salts. We have repeatedly seen that an acid of high boiling point heated with the salt of some acid of lower boiling point will drive out the low boiling acid. The boiling point of sulphuric acid (338 deg.) is higher than that of almost any common acid; hence it is used largely in the preparation of other acids.
5. Action on water. Concentrated sulphuric acid has a very great affinity for water, and is therefore an effective dehydrating agent. Gases which have no chemical action upon sulphuric acid can be freed from water vapor by bubbling them through the strong acid. When the acid is diluted with water much heat is set free, and care must be taken to keep the liquid thoroughly stirred during the mixing, and to pour the acid into the water,—never the reverse.
Not only can sulphuric acid absorb water, but it will often withdraw the elements hydrogen and oxygen from a compound containing them, decomposing the compound, and combining with the water so formed. For this reason most organic substances, such as sugar, wood, cotton, and woolen fiber, and even flesh, all of which contain much oxygen and hydrogen in addition to carbon, are charred or burned by the action of the concentrated acid.
Salts of sulphuric acid,—sulphates. The sulphates form a very important class of salts, and many of them have commercial uses. Copperas (iron sulphate), blue vitriol (copper sulphate), and Epsom salt (magnesium sulphate) serve as examples. Many sulphates are important minerals, prominent among these being gypsum (calcium sulphate) and barytes (barium sulphate).
Thiosulphuric acid (H_{2}S_{2}O_{3}); Thiosulphates. Many other acids of sulphur containing oxygen are known, but none of them are of great importance. Most of them cannot be prepared in a pure state, and are known only through their salts. The most important of these is thiosulphuric acid.
When sodium sulphite is boiled with sulphur the two substances combine, forming a salt which has the composition represented in the formula Na_{2}S_{2}O_{3}:
Na_{2}SO_{3} + S = Na_{2}S_{2}O_{3}.
The substance is called sodium thiosulphate, and is a salt of the easily decomposed acid H_{2}S_{2}O_{3}, called thiosulphuric acid. This reaction is quite similar to the action of oxygen upon sulphites:
Na{2}SO{3} + O = Na{2}SO{4}.
More commonly the salt is called sodium hyposulphite, or merely "hypo." It is a white solid and is extensively used in photography, in the bleaching industry, and as a disinfectant.
Monobasic and dibasic acids. Such acids as hydrochloric and nitric acids, which have only one replaceable hydrogen atom in the molecule, or in other words yield one hydrogen ion in solution, are called monobasic acids. Acids yielding two hydrogen ions in solution are called dibasic acids. Similarly, we may have tribasic and tetrabasic acids. The three acids of sulphur are dibasic acids. It is therefore possible for each of them to form both normal and acid salts. The acid salts can be made in two ways: the acid may be treated with only half enough base to neutralize it,—
NaOH + H{2}SO{4} = NaHSO{4} + H{2}O;
or a normal salt may be treated with the free acid,—
Na_{2}SO_{4} + H_{2}SO_{4} = 2NaHSO_{4}.
Acid sulphites and sulphides may be made in the same ways.
Carbon disulphide (CS_{2}). When sulphur vapor is passed over highly heated carbon the two elements combine, forming carbon disulphide (CS_{2}), just as oxygen and carbon unite to form carbon dioxide (CO_{2}). The substance is a heavy, colorless liquid, possessing, when pure, a pleasant ethereal odor. On standing for some time, especially when exposed to sunlight, it undergoes a slight decomposition and acquires a most disagreeable, rancid odor. It has the property of dissolving many substances, such as gums, resins, and waxes, which are insoluble in most liquids, and it is extensively used as a solvent for such substances. It is also used as an insecticide. It boils at a low temperature (46 deg.), and its vapor is very inflammable, burning in the air to form carbon dioxide and sulphur dioxide, according to the equation
CS_{2} + 6O = CO_{2} + 2SO_{2}.
Commercial preparation of carbon disulphide. In the preparation of carbon disulphide an electrical furnace is employed, such as is represented in Fig. 45. The furnace is packed with carbon C, and this is fed in through the hoppers B, as fast as that which is present in the hearth of the furnace is used up. Sulphur is introduced at A, and at the lower ends of the tubes it is melted by the heat of the furnace and flows into the hearth as a liquid. An electrical current is passed through the carbon and melted sulphur from the electrodes E, heating the charge. The vapors of carbon disulphide pass up through the furnace and escape at D, from which they pass to a suitable condensing apparatus. |
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