Chapter XI: Appendix: Electrolytic Dissociation Theory

The following brief statements concerning the ionic theory and a few of its applications are intended for reference in connection with the explanations which are given in the Notes accompanying the various procedures. The reader who desires a more extended discussion of the fundamental theory and its uses is referred to such books as Talbot and Blanchard's !Electrolytic Dissociation Theory! (Macmillan Company), or Alexander Smith's !Introduction to General Inorganic Chemistry! (Century Company).

The !electrolytic dissociation theory!, as propounded by Arrhenius in 1887, assumes that acids, bases, and salts (that is, electrolytes) in aqueous solution are dissociated to a greater or less extent into !ions!. These ions are assumed to be electrically charged atoms or groups of atoms, as, for example, H^{+} and Br^{-} from hydrobromic acid, Na^{+} and OH^{-} from sodium hydroxide, 2NH_{4}^{+} and SO_{4}^{--} from ammonium sulphate. The unit charge is that which is dissociated with a hydrogen ion. Those upon other ions vary in sign and number according to the chemical character and valence of the atoms or radicals of which the ions are composed. In any solution the aggregate of the positive charges upon the positive ions (!cations!) must always balance the aggregate negative charges upon the negative ions (!anions!).

It is assumed that the Na^{+} ion, for example, differs from the sodium atom in behavior because of the very considerable electrical charge which it carries and which, as just stated, must, in an electrically neutral solution, be balanced by a corresponding negative charge on some other ion. When an electric current is passed through a solution of an electrolyte the ions move with and convey the current, and when the cations come into contact with the negatively charged cathode they lose their charges, and the resulting electrically neutral atoms (or radicals) are liberated as such, or else enter at once into chemical reaction with the components of the solution.

Two ions of identically the same composition but with different electrical charges may exhibit widely different properties. For example, the ion MnO_{4}^{-} from permanganates yields a purple-red solution and differs in its chemical behavior from the ion MnO_{4}^{--} from manganates, the solutions of which are green.

The chemical changes upon which the procedures of analytical chemistry depend are almost exclusively those in which the reacting substances are electrolytes, and analytical chemistry is, therefore, essentially the chemistry of the ions. The percentage dissociation of the same electrolyte tends to increase with increasing dilution of its solution, although not in direct proportion. The percentage dissociation of different electrolytes in solutions of equivalent concentrations (such, for example, as normal solutions) varies widely, as is indicated in the following tables, in which approximate figures are given for tenth-normal solutions at a temperature of about 18°C.

ACIDS
=========================================================================
|
SUBSTANCE | PERCENTAGE DISSOCIATION IN
| 0.1 EQUIVALENT SOLUTION
_____________________________________________|___________________________
|
HCl, HBr, HI, HNO_{3} | 90
|
HClO_{3}, HClO_{4}, HMnO_{4} | 90
|
H_{2}SO_{4} <--> H^{+} + HSO_{4}^{-} | 90
|
H_{2}C_{2}O_{4} <--> H^{+} + HC_{2}O_{4}^{-} | 50
|
H_{2}SO_{3} <--> H^{+} + HSO{_}3^{-} | 20
|
H_{3}PO_{4} <--> H^{+} + H_{2}PO_{4}^{-} | 27
|
H_{2}PO_{4}^{-} <--> H^{+} + HPO_{4}^{--} | 0.2
|
H_{3}AsO_{4} <--> H^{+} + H_{2}AsO_{4}^{-} | 20
|
HF | 9
|
HC_{2}H_{3}O_{2} | 1.4
|
H_{2}CO_{3} <--> H^{+} + HCO_{3}^{-} | 0.12
|
H_{2}S <--> H^{+} + HS^{-} | 0.05
|
HCN | 0.01
|
=========================================================================

BASES
=========================================================================
|
SUBSTANCE | PERCENTAGE DISSOCIATION IN
| 0.1 EQUIVALENT SOLUTION
_____________________________________________|___________________________
|
KOH, NaOH | 86
|
Ba(OH)_{2} | 75
|
NH_{4}OH | 1.4
|
=========================================================================

SALTS
=========================================================================
|
TYPE OF SALT | PERCENTAGE DISSOCIATION IN
| 0.1 EQUIVALENT SOLUTION
_____________________________________________|___________________________
|
R^{+}R^{-} | 86
|
R^{++}(R^{-})_{2} | 72
|
(R^{+})_{2}R^{--} | 72
|
R^{++}R^{--} | 45
|
=========================================================================

The percentage dissociation is determined by studying the electrical conductivity of the solutions and by other physico-chemical methods, and the following general statements summarize the results:

!Salts!, as a class, are largely dissociated in aqueous solution.

!Acids! yield H^{+} ions in water solution, and the comparative !strength!, that is, the activity, of acids is proportional to the concentration of the H^{+} ions and is measured by the percentage dissociation in solutions of equivalent concentration. The common mineral acids are largely dissociated and therefore give a relatively high concentration of H^{+} ions, and are commonly known as "strong acids." The organic acids, on the other hand, belong generally to the group of "weak acids."

!Bases! yield OH^{-} ions in water solution, and the comparative strength of the bases is measured by their relative dissociation in solutions of equivalent concentration. Ammonium hydroxide is a weak base, as shown in the table above, while the hydroxides of sodium and potassium exhibit strongly basic properties.

Ionic reactions are all, to a greater or less degree, !reversible reactions!. A typical example of an easily reversible reaction is that representing the changes in ionization which an electrolyte such as acetic acid undergoes on dilution or concentration of its solutions, !i.e.!, HC_{2}H_{3}O_{2} <--> H^{+} + C_{2}H_{3}O_{2}^{-}. As was stated above, the ionization increases with dilution, the reaction then proceeding from left to right, while concentration of the solution occasions a partial reassociation of the ions, and the reaction proceeds from right to left. To understand the principle underlying these changes it is necessary to consider first the conditions which prevail when a solution of acetic acid, which has been stirred until it is of uniform concentration throughout, has come to a constant temperature. A careful study of such solutions has shown that there is a definite state of equilibrium between the constituents of the solution; that is, there is a definite relation between the undissociated acetic acid and its ions, which is characteristic for the prevailing conditions. It is not, however, assumed that this is a condition of static equilibrium, but rather that there is continual dissociation and association, as represented by the opposing reactions, the apparent condition of rest resulting from the fact that the amount of change in one direction during a given time is exactly equal to that in the opposite direction. A quantitative study of the amount of undissociated acid, and of H^{+} ions and C_{2}H_{3}O_{2}^{-} ions actually to be found in a large number of solutions of acetic acid of varying dilution (assuming them to be in a condition of equilibrium at a common temperature), has shown that there is always a definite relation between these three quantities which may be expressed thus:

(!Conc'n H^{+} x Conc'n C_{2}H_{3}O_{2}^{-})/Conc'n HC_{2}H_{3}O_{2} = Constant!.

In other words, there is always a definite and constant ratio between the product of the concentrations of the ions and the concentration of the undissociated acid when conditions of equilibrium prevail.

It has been found, further, that a similar statement may be made regarding all reversible reactions, which may be expressed in general terms thus: The rate of chemical change is proportional to the product of the concentrations of the substances taking part in the reaction; or, if conditions of equilibrium are considered in which, as stated, the rate of change in opposite directions is assumed to be equal, then the product of the concentrations of the substances entering into the reaction stands in a constant ratio to the product of the concentrations of the resulting substances, as given in the expression above for the solutions of acetic acid. This principle is called the !Law of Mass Action!.

It should be borne in mind that the expression above for acetic acid applies to a wide range of dilutions, provided the temperature remains constant. If the temperature changes the value of the constant changes somewhat, but is again uniform for different dilutions at that temperature. The following data are given for temperatures of about 18°C.[1]

==========================================================================
| | | |
MOLAL | FRACTION | MOLAL CONCENTRA- | MOLAL CONCENTRA- | VALUE OF
CONCENTRATION | IONIZED | TION OF H^{+} AND| TION OF UNDIS- | CONSTANT
CONSTANT | | ACETATE^{-} IONS | SOCIATED ACID |
______________|__________|__________________|__________________|__________
| | | |
1.0 | .004 | .004 | .996 | .0000161
| | | |
0.1 | .013 | .0013 | .0987 | .0000171
| | | |
0.01 | .0407 | .000407 | .009593 | .0000172
| | | |
===========================================================================

[Footnote 1: Alexander Smith, !General Inorganic Chemistry!, p. 579.]

The molal concentrations given in the table refer to fractions of a gram-molecule per liter of the undissociated acid, and to fractions of the corresponding quantities of H^{+} and C_{2}H_{3}O_{2}^{-} ions per liter which would result from the complete dissociation of a gram-molecule of acetic acid. The values calculated for the constant are subject to some variation on account of experimental errors in determining the percentage ionized in each case, but the approximate agreement between the values found for molal and centimolal (one hundredfold dilution) is significant.

The figures given also illustrate the general principle, that the !relative! ionization of an electrolyte increases with the dilution of its solution. If we consider what happens during the (usually) brief period of dilution of the solution from molal to 0.1 molal, for example, it will be seen that on the addition of water the conditions of concentration which led to equality in the rate of change, and hence to equilibrium in the molal solution, cease to exist; and since the dissociating tendency increases with dilution, as just stated, it is true at the first instant after the addition of water that the concentration of the undissociated acid is too great to be permanent under the new conditions of dilution, and the reaction, HC_{2}H_{3}O_{2} <--> H^{+} + C_{2}H_{3}O_{2}^{-}, will proceed from left to right with great rapidity until the respective concentrations adjust themselves to the new conditions.

That which is true of this reaction is also true of all reversible reactions, namely, that any change of conditions which occasions an increase or a decrease in concentration of one or more of the components causes the reaction to proceed in one direction or the other until a new state of equilibrium is established. This principle is constantly applied throughout the discussion of the applications of the ionic theory in analytical chemistry, and it should be clearly understood that whenever an existing state of equilibrium is disturbed as a result of changes of dilution or temperature, or as a consequence of chemical changes which bring into action any of the constituents of the solution, thus altering their concentrations, there is always a tendency to re-establish this equilibrium in accordance with the law. Thus, if a base is added to the solution of acetic acid the H^{+} ions then unite with the OH^{-} ions from the base to form undissociated water. The concentration of the H^{+} ions is thus diminished, and more of the acid dissociates in an attempt to restore equilbrium, until finally practically all the acid is dissociated and neutralized.

Similar conditions prevail when, for example, silver ions react with chloride ions, or barium ions react with sulphate ions. In the former case the dissociation reaction of the silver nitrate is AgNO_{3} <--> Ag^{+} + NO_{3}^{-}, and as soon as the Ag^{+} ions unite with the Cl^{-} ions the concentration of the former is diminished, more of the AgNO_{3} dissociates, and this process goes on until the Ag^{+} ions are practically all removed from the solution, if the Cl^{-} ions are present in sufficient quantity.

For the sake of accuracy it should be stated that the mass law cannot be rigidly applied to solutions of those electrolytes which are largely dissociated. While the explanation of the deviation from quantitative exactness in these cases is not known, the law is still of marked service in developing analytical methods along more logical lines than was formerly practicable. It has not seemed wise to qualify each statement made in the Notes to indicate this lack of quantitative exactness. The student should recognize its existence, however, and will realize its significance better as his knowledge of physical chemistry increases.

If we apply the mass law to the case of a substance of small solubility, such as the compounds usually precipitated in quantitative analysis, we derive what is known as the !solubility product!, as follows: Taking silver chloride as an example, and remembering that it is not absolutely insoluble in water, the equilibrium expression for its solution is:

(!Conc'n Ag^{+} x Conc'n Cl^{-})/Conc'n AgCl = Constant!.

But such a solution of silver chloride which is in contact with the solid precipitate must be saturated for the existing temperature, and the quantity of undissociated AgCl in the solution is definite and constant for that temperature. Since it is a constant, it may be eliminated, and the expression becomes !Conc'n Ag^{+} x Conc'n Cl^{-} = Constant!, and this is known as the solubility product. No precipitation of a specific substance will occur until the product of the concentrations of its ions in a solution exceeds the solubility product for that substance; whenever that product is exceeded precipitation must follow.

It will readily be seen that if a substance which yields an ion in common with the precipitated compound is added to such a solution as has just been described, the concentration of that ion is increased, and as a result the concentration of the other ion must proportionately decrease, which can only occur through the formation of some of the undissociated compound which must separate from the already saturated solution. This explains why the addition of an excess of the precipitant is often advantageous in quantitative procedures. Such a case is discussed at length in Note 2 on page 113.

Similarly, the ionization of a specific substance in solution tends to diminish on the addition of another substance with a common ion, as, for instance, the addition of hydrochloric acid to a solution of hydrogen sulphide. Hydrogen sulphide is a weak acid, and the concentration of the hydrogen ions in its aqueous solutions is very small. The equilibrium in such a solution may be represented as:

(!(Conc'n H^{+})^{2} x Conc'n S^{--})/Conc'n H_{2}S = Constant!, and a marked increase in the concentration of the H^{+} ions, such as would result from the addition of even a small amount of the highly ionized hydrochloric acid, displaces the point of equilibrium and some of the S^{--} ions unite with H^{+} ions to form undissociated H_{2}S. This is of much importance in studying the reactions in which hydrogen sulphide is employed, as in qualitative analysis. By a parallel course of reasoning it will be seen that the addition of a salt of a weak acid or base to solutions of that acid or base make it, in effect, still weaker because they decrease its percentage ionization.

To understand the changes which occur when solids are dissolved where chemical action is involved, it should be remembered that no substance is completely insoluble in water, and that those products of a chemical change which are least dissociated will first form. Consider, for example, the action of hydrochloric acid upon magnesium hydroxide. The minute quantity of dissolved hydroxide dissociates thus: Mg(OH)_{2} <--> Mg^{++} + 2OH^{-}. When the acid is introduced, the H^{+} ions of the acid unite with the OH^{-} ions to form undissociated water. The concentration of the OH^{-} ions is thus diminished, more Mg(OH)_{2} dissociates, the solution is no longer saturated with the undissociated compound, and more of the solid dissolves. This process repeats itself with great rapidity until, if sufficient acid is present, the solid passes completely into solution.

Exactly the same sort of process takes place if calcium oxalate, for example, is dissolved in hydrochloric acid. The C_{2}O_{4}^{--} ions unite with the H^{+} ions to form undissociated oxalic acid, the acid being less dissociated than normally in the presence of the H^{+} ions from the hydrochloric acid (see statements regarding hydrogen sulphide above). As the undissociated oxalic acid forms, the concentration of the C_{2}O_{4}^{--} ions lessens and more CaC_{2}O_{4} dissolves, as described for the Mg(OH)_{2} above. Numerous instances of the applications of these principles are given in the Notes.

Water itself is slightly dissociated, and although the resulting H^{+} and OH^{-} ions are present only in minute concentrations (1 mol. of dissociated water in 10^{7} liters), yet under some conditions they may give rise to important consequences. The term !hydrolysis! is applied to the changes which result from the reaction of these ions. Any salt which is derived from a weak base or a weak acid (or both) is subject to hydrolytic action. Potassium cyanide, for example, when dissolved in water gives an alkaline solution because some of the H^{+} ions from the water unite with CN^{-} ions to form (HCN), which is a very weak acid, and is but very slightly dissociated. Potassium hydroxide, which might form from the OH^{-} ions, is so largely dissociated that the OH^{-} ions remain as such in the solution. The union of the H^{+} ions with the CN^{-} ions to form the undissociated HCN diminishes the concentration of the H^{+} ions, and more water dissociates (H_{2}O <--> H^{+} + OH^{-}) to restore the equilibrium. It is clear, however, that there must be a gradual accumulation of OH^{-} ions in the solution as a result of these changes, causing the solution to exhibit an alkaline reaction, and also that ultimately the further dissociation of the water will be checked by the presence of these ions, just as the dissociation of the H_{2}S was lessened by the addition of HCl.

An exactly opposite result follows the solution of such a salt as Al_{2}(SO_{4})_{3} in water. In this case the acid is strong and the base weak, and the OH^{-} ions form the little dissociated Al(OH)_{3}, while the H^{+} ions remain as such in the solution, sulphuric acid being extensively dissociated. The solution exhibits an acid reaction.

Such hydrolytic processes as the above are of great importance in analytical chemistry, especially in the understanding of the action of indicators in volumetric analysis. (See page 32.)

The impelling force which causes an element to pass from the atomic to the ionic condition is termed !electrolytic solution pressure!, or ionization tension. This force may be measured in terms of electrical potential, and the table below shows the relative values for a number of elements.

In general, an element with a greater solution pressure tends to cause the deposition of an element of less solution pressure when placed in a solution of its salt, as, for instance, when a strip of zinc or iron is placed in a solution of a copper salt, with the resulting precipitation of metallic copper.

Hydrogen is included in the table, and its position should be noted with reference to the other common elements. For a more extended discussion of this topic the student should refer to other treatises.

POTENTIAL SERIES OF THE METALS

__________________________________________________________________ | | | | POTENTIAL | | POTENTIAL | IN VOLTS | | IN VOLTS _____________________|___________|____________________|___________ | | | Sodium Na^{+} | +2.44 | Lead Pb^{++} | -0.13 Calcium Ca^{++} | | Hydrogen H^{+} | -0.28 Magnesium Mg^{++} | | Bismuth Bi^{+++}| Aluminum A1^{+++} | +1.00 | Antimony | -0.75 Manganese Mn^{++} | | Arsenic | Zinc Zn^{++} | +0.49 | Copper Cu^{++} | -0.61 Cadmium Cd^{++} | +0.14 | Mercury Hg^{+} | -1.03 Iron Fe^{++} | +0.063 | Silver Ag^{+} | -1.05 Cobalt Co^{++} | -0.045 | Platinum | Nickel Ni^{++} | -0.049 | Gold | Tin Sn^{++} | -0.085(?) | | _____________________|___________|____________________|__________

THE FOLDING OF A FILTER PAPER

If a filter paper is folded along its diameter, and again folded along the radius at right angles to the original fold, a cone is formed on opening, the angle of which is 60°. Funnels for analytical use are supposed to have the same angle, but are rarely accurate. It is possible, however, with care, to fit a filter thus folded into a funnel in such a way as to prevent air from passing down between the paper and the funnel to break the column of liquid in the stem, which aids greatly, by its gentle suction, in promoting the rate of filtration.

Such a filter has, however, the disadvantage that there are three thicknesses of paper back of half of its filtering surface, as a consequence of which one half of a precipitate washes or drains more slowly. Much time may be saved in the aggregate by learning to fold a filter in such a way as to improve its effective filtering surface. The directions which follow, though apparently complicated on first reading, are easily applied and easily remembered. Use a 6-inch filter for practice. Place four dots on the filter, two each on diameters which are at right angles to each other. Then proceed as follows: (1) Fold the filter evenly across one of the diameters, creasing it carefully; (2) open the paper, turn it over, rotate it 90° to the right, bring the edges together and crease along the other diameter; (3) open, and rotate 45° to the right, bring edges together, and crease evenly; (4) open, and rotate 90° to the right, and crease evenly; (5) open, turn the filter over, rotate 22-(1/2)° to the right, and crease evenly; (6) open, rotate 45° to the right and crease evenly; (7) open, rotate 45° to the right and crease evenly; (8) open, rotate 45° to the right and crease evenly; (9) open the filter, and, starting with one of the dots between thumb and forefinger of the right hand, fold the second crease to the left over on it, and do the same with each of the other dots. Place it, thus folded, in the funnel, moisten it, and fit to the side of the funnel. The filter will then have four short segments where there are three thicknesses and four where there is one thickness, but the latter are evenly distributed around its circumference, thus greatly aiding the passage of liquids through the paper and hastening both filtration and washing of the whole contents of the filter.

!SAMPLE PAGES FOR LABORATORY RECORDS!

!Page A!

Date

CALIBRATION OF BURETTE No.

___________________________________________________________________________
| | | |
BURETTE | DIFFERENCE | OBSERVED | DIFFERENCE | CALCULATED
READINGS | | WEIGHTS | | CORRECTION
_______________|______________|______________|______________|______________
0.02 | | 16.27 | |
10.12 | 10.10 | 26.35 | 10.08 | -.02
20.09 | 9.97 | 36.26 | 9.91 | -.06
30.16 | 10.07 | 46.34 | 10.08 | +.01
40.19 | 10.03 | 56.31 | 9.97 | -.06
50.00 | 9.81 | 66.17 | 9.86 | +.05
_______________|______________|______________|______________|______________

These data to be obtained in duplicate for each burette.

!Page B!

Date

DETERMINATION OF COMPARATIVE STRENGTH HCl vs. NaOH

___________________________________________________________________________
| |
DETERMINATION | I | II
_________________________|________________________|________________________
| |
| Corrected | Corrected
Final Reading HCl | 48.17 48.08 | 43.20 43.14
Initial Reading HCl | 0.12 .12 | .17 .17
| ----- ----- | ----- -----
| 47.96 | 42.97
| |
| Corrected | Corrected
Final Reading HCl | 46.36 46.29 | 40.51 40.37
Initial Reading HCl | 1.75 1.75 | .50 .50
| ----- ----- | ----- -----
| 44.54 | 39.87
| |
log cc. NaOH | 1.6468 | 1.6008
colog cc. HCl | 8.3192 | 8.3668
| ------ | ------
| 9.9680 - 10 | 9.9676 - 10
1 cc. HCl | .9290 cc. NaOH | .9282 cc. NaOH
Mean | .9286 |
_________________________|________________________|________________________

Signed

!Page C! Date

STANDARDIZATION OF HYDROCHLORIC ACID
=====================================================================
| |
Weight sample and tube| 9.1793 | 8.1731
| 8.1731 | 6.9187
| ------ | ------
Weight sample | 1.0062 | 1.2544
| |
Final Reading HCl | 39.97 39.83 | 49.90 49.77
Initial Reading HCl | .00 .00 | .04 .04
| ----- ----- | ----- -----
| 39.83 | 49.73
| |
Final Reading NaOH | .26 .26 | .67 .67
Initial Reading NaOH | .12 .12 | .36 .36
| --- --- | --- ---
| .14 | .31
| |
| .14 | .31
Corrected cc. HCl | 39.83 - ----- = 39.68 | 49.73 - ----- = 49.40
| .9286 | .9286
| |
log sample | 0.0025 | 0.0983
colog cc | 8.4014 - 10 | 8.3063 - 10
colog milli equivalent| 1.2757 | 1.2757
| ------ | ------
| 9.6796 - 10 | 9.6803 - 10
| |
Normal value HCl | .4782 | .4789
Mean | .4786 |
| |
=====================================================================

Signed

!Page D! Date

DETERMINATION OF CHLORINE IN CHLORIDE, SAMPLE No.
=====================================================================
| |
Weight sample and tube| 16.1721 | 15.9976
| 15.9976 | 15.7117
| ------- | -------
Weight sample | .1745 | .2859
| |
Weight crucible | |
+ precipitate | 14.4496 | 15.6915
Constant weights | 14.4487 | 15.6915
| 14.4485 |
| |
Weight crucible | 14.2216 | 15.3196
Constant weight | 14.2216 | 15.3194
| |
Weight AgCl | .2269 | .3721
| |
log Cl | 1.5496 | 1.5496
log weight AgCl | 9.3558 - 10 | 9.5706 - 10
log 100 | 2.0000 | 2.0000
colog AgCl | 7.8438 - 10 | 7.7438 - 10
colog sample | 0.7583 | 0.5438
| ------- | -------
| 1.5075 | 1.5078
| |
Cl in sample No. | 32.18% | 32.20%
| |
=====================================================================

Signed

STRENGTH OF REAGENTS

The concentrations given in this table are those suggested for use in the procedures described in the foregoing pages. It is obvious, however, that an exact adherence to these quantities is not essential.

Approx. Approx.
Grams relation relation
per to normal to molal
liter. solution solution

Ammonium oxalate, (NH_{4})_{2}C_{2}O_{4}.H_{2}O 40 0.5N 0.25 Barium chloride, BaCl_{2}.2H_{2}O 25 0.2N 0.1 Magnesium ammonium chloride (of MgCl_{2}) 71 1.5N 0.75 Mercuric chloride, HgCl_{2} 45 0.33N 0.66 Potassium hydroxide, KOH (sp. gr. 1.27) 480 Potassium thiocyanate, KSCN 5 0.05N 0.55 Silver nitrate, AgNO_{3} 21 0.125N 0.125 Sodium hydroxide, NaOH 100 2.5N 2.5 Sodium carbonate. Na_{2}CO_{3} 159 3N 1.5 Sodium phosphate, Na_{2}HPO_{4}.12H_{2}O 90 0.5N or 0.75N 0.25

Stannous chloride, SnCl_{2}, made by saturating hydrochloric acid (sp. gr. 1.2) with tin, diluting with an equal volume of water, and adding a slight excess of acid from time to time. A strip of metallic tin is kept in the bottle.

A solution of ammonium molybdate is best prepared as follows: Stir 100 grams of molybdic acid (MoO_{3}) into 400 cc. of cold, distilled water. Add 80 cc. of concentrated ammonium hydroxide (sp. gr. 0.90). Filter, and pour the filtrate slowly, with constant stirring, into a mixture of 400 cc. concentrated nitric acid (sp. gr. 1.42) and 600 cc. of water. Add to the mixture about 0.05 gram of microcosmic salt. Filter, after allowing the whole to stand for 24 hours.

The following data regarding the common acids and aqueous ammonia are based upon percentages given in the Standard Tables of the Manufacturing Chemists' Association of the United States [!J.S.C.I.!, 24 (1905), 787-790]. All gravities are taken at 15.5°C. and compared with water at the same temperature.

Aqueous ammonia (sp. gr. 0.96) contains 9.91 per cent NH_{3} by weight, and corresponds to a 5.6 N and 5.6 molal solution.

Aqueous ammonia (sp. gr. 0.90) contains 28.52 per cent NH_{3} by weight, and corresponds to a 5.6 N and 5.6 molal solution.

Hydrochloric acid (sp. gr. 1.12) contains 23.81 per cent HCl by weight, and corresponds to a 7.3 N and 7.3 molal solution.

Hydrochloric acid (sp. gr. 1.20) contains 39.80 per cent HCl by weight, and corresponds to a 13.1 N and 13.1 molal solution.

Nitric acid (sp. gr. 1.20) contains 32.25 per cent HNO_{3} by weight, and corresponds to a 6.1 N and 6.1 molal solution:

Nitric acid (sp. gr. 1.42) contains 69.96 per cent HNO_{3} by weight, and corresponds to a 15.8 N and 15.8 molal solution.

Sulphuric acid (sp. gr. 1.8354) contains 93.19 per cent H_{2}SO_{4} by weight, and corresponds to a 34.8 N or 17.4 molal solution.

Sulphuric acid (sp. gr. 1.18) contains 24.74 per cent H_{2}SO_{4} by weight, and corresponds to a 5.9 N or 2.95 molal solution.

The term !normal! (N), as used above, has the same significance as in volumetric analyses. The molal solution is assumed to contain one molecular weight in grams in a liter of solution.

DENSITIES AND VOLUMES OF WATER AT TEMPERATURES FROM 15-30°C.

Temperature Density. Volume. Centigrade.

4° 1.000000 1.000000
15° 0.999126 1.000874
16° 0.998970 1.001031
17° 0.998801 1.001200
18° 0.998622 1.001380
19° 0.998432 1.001571
20° 0.998230 1.001773
21° 0.998019 1.001985
22° 0.997797 1.002208
23° 0.997565 1.002441
24° 0.997323 1.002685
25° 0.997071 1.002938
26° 0.996810 1.003201
27° 0.996539 1.003473
28° 0.996259 1.003755
29° 0.995971 1.004046
30° 0.995673 1.004346

Authority: Landolt, Börnstein, and Meyerhoffer's !Tabellen!, third edition.

CORRECTIONS FOR CHANGE OF TEMPERATURE OF STANDARD SOLUTIONS

The values below are average values computed from data relating to a considerable number of solutions. They are sufficiently accurate for use in chemical analyses, except in the comparatively few cases where the highest attainable accuracy is demanded in chemical investigations. The expansion coefficients should then be carefully determined for the solutions employed. For a compilation of the existing data, consult Landolt, Börnstein, and Meyerhoffer's !Tabellen!, third edition.

Corrections for 1 cc.
Concentration. of solution between
15° and 35°C.

Normal .00029
0.5 Normal .00025
0.1 Normal or more dilute solutions .00020

The volume of solution used should be multiplied by the values given, and that product multiplied by the number of degrees which the temperature of the solution varies from the standard temperature selected for the laboratory. The total correction thus found is subtracted from the observed burette reading if the temperature is higher than the standard, or added, if it is lower. Corrections are not usually necessary for variations of temperature of 2°C. or less.

INTERNATIONAL ATOMIC WEIGHTS

========================================================== | | | | 1920 | | 1920 _________________|_________|___________________|__________ | | | Aluminium Al | 27.1 | Molybdenum Mo | 96.0 Antimony Sb | 120.2 | Neodymium Nd | 144.3 Argon A | 39.9 | Neon Ne | 20.2 Arsenic As | 74.96 | Nickel Ni | 58.68 Barium Ba | 137.37 | Nitrogen N | 14.008 Bismuth Bi | 208.0 | Osmium Os | 190.9 Boron B | 11.0 | Oxygen O | 16.00 Bromine Br | 79.92 | Palladium Pd | 106.7 Cadmium Cd | 112.40 | Phosphorus P | 31.04 Caesium Cs | 132.81 | Platinum Pt | 195.2 Calcium Ca | 40.07 | Potassium K | 39.10 Carbon C | 12.005 | Praseodymium Pr | 140.9 Cerium Ce | 140.25 | Radium Ra | 226.0 Chlorine Cl | 35.46 | Rhodium Rh | 102.9 Chromium Cr | 52.0 | Rubidium Rb | 85.45 Cobalt Co | 58.97 | Ruthenium Ru | 101.7 Columbium Cb | 93.1 | Samarium Sm | 150.4 Copper Cu | 63.57 | Scandium Sc | 44.1 Dysprosium Dy | 162.5 | Selenium Se | 79.2 Erbium Er | 167.7 | Silicon Si | 28.3 Europium Eu | 152.0 | Silver Ag | 107.88 Fluorine Fl | 19.0 | Sodium Na | 23.00 Gadolinium Gd | 157.3 | Strontium Sr | 87.63 Gallium Ga | 69.9 | Sulphur S | 32.06 Germanium Ge | 72.5 | Tantalum Ta | 181.5 Glucinum Gl | 9.1 | Tellurium Te | 127.5 Gold Au | 197.2 | Terbium Tb | 159.2 Helium He | 4.00 | Thallium Tl | 204.0 Hydrogen H | 1.008 | Thorium Th | 232.4 Indium In | 114.8 | Thulium Tm | 168.5 Iodine I | 126.92 | Tin Sn | 118.7 Iridium Ir | 193.1 | Titanium Ti | 48.1 Iron Fe | 55.84 | Tungsten W | 184.0 Krypton Kr | 82.92 | Uranium U | 238.2 Lanthanum La | 139.0 | Vanadium V | 51.0 Lead Pb | 207.2 | Xenon Xe | 130.2 Lithium Li | 6.94 | Ytterbium Yb | 173.5 Lutecium Lu | 175.0 | Yttrium Y | 88.7 Magnesium Mg | 24.32 | Zinc Zn | 65.37 Manganese Mn | 54.93 | Zirconium Zr | 90.6 Mercury Hg | 200.6 | | ==========================================================

INDEX

Acidimetry Acid solutions, normal standard Acids, definition of Acids, weak, action of other acids on action of salts on Accuracy demanded Alkalimetry Alkali solutions, normal standard Alumina, determination of in stibnite Ammonium nitrate, acid Analytical chemistry, subdivisions of Antimony, determination of, in stibnite Apatite, analysis of Asbestos filters Atomic weights, table of

Balances, essential features of use and care of Barium sulphate, determination of sulphur in Bases, definition of Bichromate process for iron Bleaching powder, analysis of Brass, analysis of Burette, description of calibration of cleaning of reading of

Calcium, determination of, in limestone Calibration, definition of of burettes of flasks Carbon dioxide, determination of, in limestone Chlorimetry Chlorine, gravimetric determination of Chrome iron ore, analysis of Coin, determination of silver in Colloidal solution of precipitates Colorimetric analyses, definition of Copper, determination of, in brass determination of in copper ores Crucibles, use of Crystalline precipitates

Densities of water Deposition potentials Desiccators Direct methods Dissociation, degree of

Economy of time Electrolytic dissociation, theory of Electrolytic separations, principles of End-point, definition of Equilibrium, chemical Evaporation of liquids

Faraday's law Feldspar, analysis of Ferrous ammonium sulphate, analysis of Filters, folding of how fitted Filtrates, testing of Filtration Flasks, graduation of Funnels Fusions, removal of from crucibles

General directions for gravimetric analysis volumetric analysis Gooch filter Gravimetric analysis, definition of

Hydrochloric acid, standardization of Hydrolysis

Ignition of precipitates Indicators, definition of for acidimetry preparation of Indirect methods Insoluble matter, determination of in limestone Integrity Iodimetry Ions, definition of Iron, gravimetric determination of volumetric determination of

Jones reductor

Lead, determination of in brass Limestone, analysis of Limonite, determination of iron in Liquids, evaporation of transfer of Litmus Logarithms

Magnesium, determination of Mass action, law of Measuring instruments Methyl orange Moisture, determination of in limestone

Neutralization methods Normal solutions, acid and alkali oxidizing agents reducing agents Notebooks, sample pages of

Oxalic acid, determination of strength of Oxidation processes Oxidizing power of pyrolusite

Permanganate process for iron Phenolphthalein Phosphoric anhydride, determination of Pipette, calibration of description of Platinum crucibles, care of Precipitates, colloidal crystalline ignition of separation from filter washing of Precipitation Precipitation methods (volumetric) Problems Pyrolusite, oxidizing power of

Quantitative Analyses, subdivisions of

Reagents, strength of Reducing solution, normal Reductor, Jones Reversible reactions

Silica, determination of, in limestone determination of, in silicates purification of Silicic acid, dehydration of Silver, determination of in coin Soda ash, alkaline strength of Sodium chloride, determination of chlorine in Solubility product Solution pressure Solutions, normal standard Standardization, definition of Standard solutions, acidimetry and alkalimetry chlorimetry iodimetry oxidizing and reducing agents thiocyanate Starch solutions Stibnite, determination of antimony in Stirring rods Stoichiometry Strength of reagents Suction, use of Sulphur, determination of in ferrous ammonium sulphate in barium sulphate

Temperature, corrections for Testing of washings Theory of electrolytic dissociation Thiocyanate process for silver Titration, definition of Transfer of liquids

Volumetric analysis, definition of
general directions

Wash-bottles Washed filters Washing of precipitates Washings, testing of Water, ionization of densities of Weights, care of

Zimmermann-Reinhardt method for iron Zinc, determination of, in brass