Which substances are partially dissolved in water?

Water and solutions: solubility of salts, gases, etc. Substances, temperature dependence, density of salt solutions, pH value and calculation

Updated / revised on 06/28/2009

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1. Water and solvent power

The dissolving power of water differs depending on the substance. Completely insoluble in water is not a substance, but there are many substances that are so insoluble that in practice they can be considered insoluble. On the other hand, a number of liquids are miscible with water in any proportions, e.g. B. alcohol, methanol, acetic acid, sulfuric acid, glycol and glycerin. Gases also dissolve in water. Their solubility increases with pressure and with their concentration in the air.

The solubility of all substances is more or less dependent on the temperature. With common table salt (NaCl) z. B. it changes very little from 0 ° C to 100 ° C. In contrast, cesium aluminum sulfate dissolves 86 times better at 100 ° C than at 0 ° C. In the case of gases, the solubility usually decreases with temperature; in the case of solids, however, it usually increases. An exception is e.g. B. Calcium sulphate (gypsum), which dissolves better in cold water than in hot water.

When a substance is dissolved, its molecular structure is destroyed and the molecules "swim" around in the water, detached from one another. Acids, alkalis and salts also disintegrate ("dissociate") in whole or in part into ions (positively and negatively charged). Examples:
Table salt (NaCl) breaks down into Na (+) and Cl (-).
Sulfuric acid (H2SO4) breaks down into H (+) and HSO4 (-).
The HSO4 (-) decays z. T. continue in H (+) and SO4 (2-).
Sodium hydroxide (NaOH) breaks down into Na (+) and OH (-).
Even water itself breaks down to a very small extent (0.00000018%) into ions: into H (+) and OH (-).

Other substances, such as B. sugar or alcohol, practically do not break down into ions.

If several substances are to be dissolved in water at the same time, they more or less influence each other in terms of their solubility. This effect is particularly strong with different salts that have a common type of ion, e.g. B. Sodium Chloride and Potassium Chloride (NaCl and KCl). The Cl (-) ion is formed from both when dissolved in water. In a solution that contains KCl, less NaCl dissolves than in pure water and vice versa.

Some examples of solubilities in water at different temperatures in% by weight (a solubility of 15% by weight means that a maximum of 15 g of the substance can be dissolved in 85 g of water (gives a total of 100 g of solution)):

Substance solubility in% by weight 20 ° C 80 ° C ====================================== ============== Sodium Chloride 26.5 27.5 ("Table Salt") Potassium Chloride 25.5 33.6 Ammonium Chloride 27.0 40.0 ("Salmiak") Potassium Sulphate 10.0 17 , 5 calcium sulfate 0.199 0.10 ("gypsum") calcium hydroxide 0.17 0.087 calcium carbonate 0.0015 0.002 (100 ° C) zinc chloride 78.7 84.5

If the maximum soluble amount of potassium chloride is dissolved in water at 80 ° C. and the solution is allowed to cool, the solubility is reduced and the no longer soluble portion of potassium chloride crystallizes out again.


Salt solutions usually have higher densities than pure water. Using sodium chloride at 20 ° C as an example:

Conc. (Wt .-%) Density (kg / m³) ============================= 0 998 2 1013 4 1027 6 1041 10 1071 20 1148 25 1189

The solubility of gases depends on temperature and pressure. The following table shows the solubility coefficients of pure oxygen, pure nitrogen and pure CO2 in g (gas) / kg (water) / bar:

Temp. (° C) Oxygen Nitrogen CO2 ========================================= = 0 0.0676 0.0281 3.26 10 0.0526 0.0226 2.28 20 0.0428 0.0190 1.67 30 0.0364 0.0166 1.28 50 0.0291 0.0137 0, 82 70 0.0258 0.0129 0.59 90 0.0246 0.0125 ----
The solubility of a gas at a given temperature is proportional to its pressure. In the case of pure gases, this is equal to the total gas pressure that can be measured with a manometer; in the case of gas mixtures, the partial pressure (= partial pressure) of each individual gas is decisive for the solubility of each gas. The partial pressure is calculated from the total gas pressure, multiplied by the mole fraction (for the term mole see below). Under normal ambient conditions, the mole fraction and volume fraction for air constituents are almost identical, so the volume fraction (or volume%) can also be used instead of the mole fraction.

Example: Air contains approx. 21% by volume of oxygen (O2). How much oxygen dissolves at 1.2 bar air pressure in water at a temperature of 20 ° C?

1. Calculate the partial pressure of O2:
p (O2) = 1.2 bar * 21% by volume / 100% by volume = 0.252 bar

2. Calculate the solubility of O2 using the relevant solubility coefficient from the table above:
L (O2) = 0.252 bar * 0.0428 g (O2) / kg (water) / bar = 0.0108 g (O2) / kg (water)

The solubility of CO2 is strongly dependent on the pH value. This in turn is influenced by a large number of substances when they are dissolved in the water.






2. Water, solutions and the pH

The pH value is a (logarithmic) measure of the concentration of hydrogen ions in an aqueous solution. Its calculation is not that simple. You need a special unit of measurement, the mole. 1 mole means 602300000000000000000000 particles, e.g. B. water molecules, sodium ions or hydrogen ions (H (+)). Despite the huge number, a mole is usually less than 1 kg. This is because atoms and molecules are so incredibly small and light.

The chemist likes to calculate in moles because chemical reactions always require a certain number of the types of molecules involved. In the formation of water from hydrogen and oxygen z. B. 2 hydrogen molecules and 1 oxygen molecule always react with each other to form 2 water molecules:

2 H2 + 1 O2 = 2 H2O

If you do the whole thing 602300000000000000000000 times, then 2 moles of H2 react with 1 mole of O2 to form 2 moles of H2O. In the end, however, you have to convert into "weight" units (strictly speaking, they are mass units!), You just have to know how much a mole of which atomic type "weighs" or what mass it has. Since there are only about 110 types of atoms (most of which only play a role in very special cases), this is easier than it looks. Examples:

1 mol (= 602300000000000000000000 atoms) of an atomic type has the following mass:
Hydrogen (H): approx. 1.0 g Oxygen (O): approx. 16.0 g Nitrogen (N): approx. 14.0 g Carbon (C): approx. 12.0 g Sodium (Na): approx 23.0 g potassium (K): approx. 39.1 g chlorine (Cl): approx 35.5 g sulfur (S): approx 32.1 g magnesium (Mg): approx 24.3 g calcium (Ca): approx. 40.1 g fluorine (F): approx. 19.0 g phosphorus (P): approx. 31.0 g iron (Fe): approx. 55.8 g selenium (Se): approx. 79.0 g silicon (Si): approx. 28.1 g zinc (Zn): approx. 65.4 g boron (B): approx. 10.8 g copper (Cu): approx. 63.5 g
According to this, 1 mole of CO2 molecules (each of them consists of 1 C atom and 2 O atoms) have the following mass (in colloquial language: "weigh"):

1 mole of C corresponds to 12.0 g + 2 moles of O corresponds to 32.0 g ------------------------------ 1 mole of CO2 corresponds to 44.0 g ==============================
The mass that 1 mole of a certain type of molecule has is called its molar mass, or often also its molecular weight.

In the same way, the molar mass M (HCl) of hydrochloric acid gas results:

M (HCl) = 1 * 1.0 g / mol + 1 * 35.5 g / mol = 36.5 g / mol

In pure water there are always just enough water molecules disintegrating into H (+) and OH (-) (the disintegration into ions is also called dissociation) that the concentration of H (+) is approximately 1/10000000 mol / l. In other words, this is 1/10 ^ 7 mol / l or 10 ^ -7 mol / l. The logarithm of ten of 10 ^ -7 is = -7. The pH value, in turn, is the negative logarithm of the H (+) concentration in mol / l and thus - (- 7) = 7. In pure water, the pH value = 7.

Acids or acidic salts dissolved in water usually disintegrate to a much greater extent than water and therefore give off many additional H (+) ions, which increases their concentration in the water. Because of the increased H (+) concentration, so many OH (-) ions with H (+) rearrange to form H2O molecules at the same time until the concentrations of H (+) and OH (-), i.e. c (H (+)) and c (OH (-)), multiplied together again the value 1 * 10 ^ -14 mol ^ 2 / l ^ 2 result:

c (H (+)) * c (OH (-)) = 1 * 10 ^ -14 mol ^ 2 / l ^ 2

Example: Hydrochloric acid (HCl) breaks down practically completely into its ions in water:

HCl = H (+) + Cl (-).

So 1 mole of H (+) is formed from 1 mole of HCl dissolved in water. The following is an example calculation: 0.1 mol of HCl is dissolved in so much water that exactly 1 liter of solution comes out. Such a solution is called a 0.1 molar hydrochloric acid (solution) (short: 0.1M-HCl). When the HCl molecules completely disintegrate, 0.1 mol of H (+) is formed. The ions resulting from the disintegration of water can be neglected compared to this large amount or high concentration: At 0.1 mol / l H (+) the OH (-) concentration (and thus also the concentration of the H (+) Ions, which can be assigned to the self-decay of water) only 10 ^ -13 mol / l and therefore no longer plays a role. How high is the pH of the solution? Very easily:

1. Form the logarithm of the H (+) concentration, whereby the concentration value in mol / l must always be used, other units are not permitted here:

Log (0.1) = -1

2. Reverse the sign ("negate"):

pH (0.1M-HCl) = - (Log (0.1)) = - (- 1) = 1

The pH value of a "0.1 molar" hydrochloric acid is consequently 1.

Now you can convert the concentration of hydrochloric acid into the more common units of mass:

1 mole of HCl corresponds to 36.5 g

0.1 mol of HCl corresponds to 3.65 g

Thus, 1 l of 0.1M HCl solution contains 3.65 g of HCl.


Substances that do not release H (+) but OH (-) ions when they break down into ions ("bases") have an indirect influence on the pH value. The OH (-) ions react with H (+) ions again to form water (H2O) and thereby lower the H (+) concentration. The whole thing happens to the extent that the product of the H (+) and OH (-) concentration in water is always = 1/100000000000000 mol ^ 2 / l ^ 2, i.e. 10 ^ -14 mol ^ 2 / l ^ 2. For example, if 0.01 mol of NaOH is dissolved in water, it will almost completely break down into Na (+) and OH (-). The OH (-) - concentration is then = 0.01 mol / l (= 10 ^ -2 mol / l) and the H (+) - concentration consequently = 10 ^ -12 mol / l (10 ^ -2 mol / l * 10 ^ -12 mol / l = 10 ^ -14 mol ^ 2 / l ^ 2) and the pH = 12.

In the case of "weak" acids and bases, which only partially decompose into ions, the so-called "degree of dissociation" must be taken into account. Their influence on the pH is significantly less than that of strong acids or bases.

If the pH value of 1 l of pure water is to be reduced with hydrochloric acid, the following quantities of hydrochloric acid are required (in practice, however, it must be noted that hydrochloric acid is never pure HCl, but a solution of HCl (HCl is a gas!) In Water in a concentration of up to approx. 36% by weight):

pH reduction Mol HCl = g HCl corresponds to g 25% hydrochloric acid (*) ================================ ====================== From 7 to 5 0.00001 0.000365 0.00146 From 7 to 3 0.001 0.0365 0.146 From 7 to 1 0, 1 3.65 14.6
(*): Unless otherwise stated, "% ig" for liquids and solid mixtures usually always means% by weight (=% by mass).

In the pH range around 7, even the slightest addition of acid can lead to strong pH changes. It is therefore extremely difficult to try to bring pure water to a pH of 6, for example, by adding hydrochloric acid.

Strictly speaking, the very small amounts of H (+) ions that result from the disintegration of the water must also be taken into account when calculating the pH. In most cases, however, these can be neglected because the amount of acid or base is much larger.

In water that contains dissolved salts, a possibly existing "buffer effect" must also be taken into account. If the salt of a weak acid is already dissolved in the water (e.g. a carbonate, acetate or phosphate), then the H (+) ions produced by the decomposition of the added (strong) acid react with the anion of the weak acid (the one produced by the Decomposition of the salt was formed) to a considerable extent to the weak acid: The pH changes only slightly.

Example: In a sodium acetate solution (NaCH3COO) in which the salt has broken down into its ions:

NaCH3COO = Na (+) + CH3COO (-)

a little HCl is added, which in turn completely breaks down into its ions ("dissociates"):

HCl = H (+) + Cl (-)

Most of the H (+) ions formed, however, are "trapped" by the acetate ions:

H (+) + CH3COO (-) = HCH3COO

The result is that only a few of the H (+) ions formed actually remain in the solution as such, so that the pH value only decreases slightly.

The same applies to the salts of weak bases (e.g. ammonium salts) and added strong bases.

Since tap water, spring water, river and sea water almost always contain significant amounts of carbonates and in some cases also phosphates, it is not possible to calculate the amount of acid or base required to adjust the pH to a certain value from the measured pH alone. If an exact water analysis is available, the calculation of the required amount of acid or base is still not an easy matter.





3. Sources:

D'Ans-Lax: Pocket book for chemists and physicists, Volume 1: Macroscopic chem.-phys. Properties, Springer-Verlag, 3rd edition, 1967

Weast, Handbook of Chemistry and Physics, 64th Edition, 1983-1984, CRC Press Inc., Boca Raton, Florida






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