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What does the acidity of water 1 mol l. PH Metry. PH and pOH equations

HYDROGEN INDICATOR (PH).One of the most important properties of aqueous solutions is their acidity (or alkalinity), which is determined by the concentration of H + and OH - ions ( cm. ELECTROLYTIC DISSOCIATION. ELECTROLYTES). The concentrations of these ions in aqueous solutions are related by a simple dependence \u003d TO w; (square brackets are used to denote the concentration in units of mol / l). The value of Kw is called the ionic product of water and is constant at a given temperature. So, at 0 ° C it is equal to 0.11 × 10 –14, at 20 ° C –0.69 × 10 –14, and at 100 ° C – 55.0 × 10 –14. Most commonly used value K w at 25 ° C, which is equal to 1.00 × 10 –14. In absolutely pure water that does not even contain dissolved gases, the concentrations of H + and OH - ions are equal (the solution is neutral). In other cases, these concentrations do not coincide: in acidic solutions, H + ions predominate, in alkaline solutions - OH - ions. But their product in any aqueous solution is constant. Therefore, if you increase the concentration of one of these ions, then the concentration of the other ion will decrease by the same amount. So, in a weak acid solution, in which \u003d 10 –5 mol / L, \u003d 10 –9 mol / L, and their product is still equal to 10 –14. Similarly, in an alkaline solution at \u003d 3.7 × 10 –3 mol / L \u003d 10 –14 / 3.7 × 10 –3 \u003d 2.7 × 10 –11 mol / L.

From the foregoing it follows that the acidity of the solution can be unambiguously expressed by indicating the concentration of only hydrogen ions in it. For example, in pure water \u003d 10 –7 mol / L. In practice, operating with such numbers is inconvenient. In addition, the concentration of H + ions in solutions can vary by hundreds of trillions of times — from about 10–15 mol / L (strong alkali solutions) to 10 mol / L (concentrated hydrochloric acid), which cannot be shown on any graph. Therefore, it has long been agreed to indicate only the exponent 10, taken with the opposite sign, for the concentration of hydrogen ions in solution; for this, the concentration should be expressed as a degree of 10x, without a multiplier, for example, 3.7 × 10 –3 \u003d 10 –2.43. (In more accurate calculations, especially in concentrated solutions, their activity is used instead of ion concentration.) This exponent is called the hydrogen exponent, and the abbreviated pH is from the hydrogen and the German word Potenz - the mathematical degree. Thus, by definition, pH \u003d –lg [H +]; this value can vary within small limits - from –1 to 15 in total (and more often from 0 to 14). In this case, a change in the concentration of H + ions by a factor of 10 corresponds to a change in pH by one unit. The designation of pH was introduced into scientific use in 1909 by the Danish physicist and biochemist S.P. L. Sørensen, who at that time was studying the processes occurring in the fermentation of beer malt and their dependence on the acidity of the medium.

At room temperature in neutral solutions pH \u003d 7, in acidic solutions pH< 7, а в щелочных рН > 7. Approximately the pH value of an aqueous solution can be determined using indicators. For example, methyl orange at pH< 3,1 имеет красный цвет, а при рН > 4.4 - yellow; litmus at pH< 6,1 красный, а при рН > 8 - blue, etc. More precisely (up to hundredths), the pH value can be determined using special instruments - pH meters. Such devices measure the electric potential of a special electrode immersed in a solution; this potential depends on the concentration of hydrogen ions in the solution, and it can be measured with high accuracy.

It is interesting to compare the pH values \u200b\u200bof solutions of various acids, bases, salts (at a concentration of 0.1 mol / l), as well as some mixtures and natural objects. For sparingly soluble compounds marked with an asterisk, the pH of saturated solutions is shown.

Table 1. Hydrogen indicators for solutions
Solution	PH
Hcl	1,0
H 2 SO 4	1,2
H 2 C 2 O 4	1,3
NaHSO 4	1,4
H 3 RO 4	1,5
Gastric juice	1,6
Wine acid	2,0
Lemon acid	2,1
HNO 2	2,2
Lemon juice	2,3
Lactic acid	2,4
Salicylic acid	2,4
Table vinegar	3,0
Grapefruit juice	3,2
CO 2	3,7
Apple juice	3,8
H 2 s	4,1
Urine	4,8–7,5
Black coffee	5,0
Saliva	7,4–8
Milk	6,7
Blood	7,35–7,45
Bile	7,8–8,6
Water of the oceans	7,9–8,4
Fe (OH) 2	9,5
MgO	10,0
Mg (OH) 2	10,5
Na 2 CO 3	11
Ca (OH) 2	11,5
NaOH	13,0

The table allows you to make a number of interesting observations. PH values, for example, immediately show the comparative strength of acids and bases. One can also clearly see the strong change in the neutral medium as a result of hydrolysis of salts formed by weak acids and bases, as well as during the dissociation of acid salts.

Natural water always has an acid reaction (pH< 7) из-за того, что в ней растворен углекислый газ; при его реакции с водой образуется кислота: СО 2 + Н 2 О « Н + + НСО 3 2– . Если насытить воду углекислым газом при атмосферном давлении, рН полученной «газировки» будет равен 3,7; такую кислотность имеет примерно 0,0007%-ный раствор соляной кислоты – желудочный сок намного кислее! Но даже если повысить давление CO 2 над раствором до 20 атм, значение pH не опускается ниже 3,3. Это значит, что газированную воду (в умеренных количествах, конечно) можно пить без вреда для здоровья, даже если она насыщена углекислым газом.

Certain pH values \u200b\u200bare extremely important for the life of living organisms. Biochemical processes in them should proceed at strictly specified acidity. Biological catalysts - enzymes are able to work only within certain pH ranges, and when going beyond these limits their activity can sharply decrease. For example, the activity of the pepsin enzyme, which catalyzes the hydrolysis of proteins and thus promotes the digestion of protein foods in the stomach, is maximum at pH values \u200b\u200bof about 2. Therefore, for normal digestion, it is necessary that the gastric juice has fairly low pH values: normal 1.53–1, 67. With gastric ulcer, the pH decreases on average to 1.48, and with a duodenal ulcer it can even reach 105. The exact pH value of gastric juice is determined by intragastric examination (pH probe). If a person has low acidity, the doctor can prescribe a weak hydrochloric acid solution with food, and with increased acidity, take anti-acid agents, for example, magnesium or aluminum hydroxides. Interestingly, if you drink lemon juice, the acidity of the gastric juice ... will decrease! Indeed, a citric acid solution will only dilute the stronger hydrochloric acid contained in the gastric juice.

In the cells of the body, the pH is about 7, in the extracellular fluid - 7.4. Nerve endings that are outside the cells are very sensitive to changes in pH. With mechanical or thermal damage to tissues, cell walls are destroyed and their contents enter the nerve endings. As a result, a person feels pain. The Scandinavian researcher Olaf Lindahl did this experiment: using a special needleless injector, a very thin stream of solution was injected through the skin of a person, which did not damage the cells, but acted on the nerve endings. It was shown that it is precisely hydrogen cations that cause pain, and with a decrease in the pH of the solution, the pain intensifies. Likewise, a solution of formic acid, which stinging insects or nettles inject under the skin, directly “acts on the nerves”. The different pH of the tissues also explains why with some inflammations a person feels pain, and with some - no.

Interestingly, injecting clean water under the skin gave particularly severe pain. This strange at first glance phenomenon is explained as follows: when in contact with pure water, cells break as a result of osmotic pressure and their contents affect the nerve endings.

In a very narrow range, the pH of the blood should remain; even its slight acidification (acidosis) or alkalization (alkalosis) can lead to the death of the body. Acidosis is observed in diseases such as bronchitis, circulatory failure, lung tumors, pneumonia, diabetes, fever, damage to the kidneys and intestines. Alkolosis is observed with hyperventilation of the lungs (or by inhalation of pure oxygen), with anemia, CO poisoning, hysteria, a brain tumor, excessive consumption of drinking soda or alkaline mineral water, and diuretic drugs. It is interesting that the pH of arterial blood should normally be in the range 7.37–7.45, and venous - 7.34–7.43. Various microorganisms are also very sensitive to the acidity of the medium. So, pathogenic microbes quickly develop in a slightly alkaline environment, while they cannot withstand an acidic environment. Therefore, for preservation (pickling, salting) products are used, as a rule, acidic solutions, adding vinegar or food acids to them. The correct selection of pH is also of great importance for chemical-technological processes.

It is possible to maintain the desired pH value, to prevent it from deviating noticeably in one direction or another when the conditions change, using the so-called buffer (from English buff - to soften shocks) solutions. Such solutions are often a mixture of a weak acid and its salt or a weak base and its salt. Such solutions "resist" within certain limits (called the buffer capacity) to attempts to change their pH. For example, if you try to acidify the mixture of acetic acid and sodium acetate a little, the acetate ions will bind the excess H + ions to poorly dissociated acetic acid, and the pH of the solution will not change much (there are a lot of acetate ions in the buffer solution, as they form as a result of complete dissociation sodium acetate). On the other hand, if a little alkali is introduced into such a solution, the excess OH - ions will be neutralized with acetic acid while maintaining the pH value. Other buffer solutions act in a similar manner, with each of them maintaining a specific pH value. Solutions of acid salts of phosphoric acid and weak organic acids - oxalic, tartaric, citric, phthalic, etc. also have a buffering effect. The specific pH of the buffer solution depends on the concentration of the components of the buffer. So, acetate buffer allows you to maintain the pH of the solution in the range of 3.8-6.3; phosphate (a mixture of KH 2 PO 4 and Na 2 HPO 4) - in the range of 4.8 - 7.0, borate (a mixture of Na 2 B 4 O 7 and NaOH) - in the range of 9.2–11, etc.

Many natural fluids have buffering properties. An example is water in the ocean, the buffering properties of which are largely due to dissolved carbon dioxide and HCO 3 - bicarbonate ions. The source of the latter, in addition to CO 2, is huge amounts of calcium carbonate in the form of shells, chalk and limestone deposits in the ocean. Interestingly, the photosynthetic activity of plankton, one of the main suppliers of oxygen to the atmosphere, leads to an increase in the pH of the medium. This happens in accordance with the Le Chatelier principle as a result of a shift in equilibrium during the absorption of dissolved carbon dioxide: 2Н + + СО 3 2– “Н + + НСО 3 -“ Н 2 СО 3 “Н 2 О + СО 2. When CO 2 is removed from the solution during photosynthesis of CO 2 + H 2 O + hv ® 1 / n (CH 2 O) n + O 2, the equilibrium shifts to the right and the medium becomes more alkaline. In body cells, CO 2 hydration is catalyzed by the carbonic anhydrase enzyme.

Cellular fluid, blood are also examples of natural buffers. So, the blood contains about 0.025 mol / L of carbon dioxide, and its content in men is about 5% higher than in women. The concentration of bicarbonate ions in blood is approximately the same (there are also more of them in men).

In soil testing, pH is one of the most important characteristics. Different soils can have a pH from 4.5 to 10. According to the pH value, in particular, it is possible to judge the nutrient content in the soil, as well as which plants can grow successfully on this soil. For example, the growth of beans, lettuce, blackcurrant is difficult when the soil pH is below 6.0; cabbage - below 5.4; apple trees - below 5.0; potatoes - below 4.9. Acidic soils are usually less rich in nutrients, because they hold less metal cations, which are necessary for plants. For example, hydrogen ions entering the soil displace bound Ca 2+ ions from it. And aluminum ions displaced from clay (aluminosilicate) rocks in high concentrations are toxic to crops.

For deoxidation of acidic soils, their liming is used - the introduction of substances that gradually bind excess acid. Natural minerals such as chalk, limestone, dolomite, as well as lime, slag from metallurgical plants can serve as such a substance. The amount of deoxidant added depends on the buffer capacity of the soil. For example, liming clay soil requires more deoxidizing agents than sand.

Of great importance are the pH measurements of rainwater, which can be quite acidic due to the presence of sulfuric and nitric acids in it. These acids are formed in the atmosphere from nitrogen and sulfur (IV) oxides, which are emitted from waste from numerous industries, vehicles, boiler houses and thermal power plants. It is known that acid rains with a low pH (less than 5.6) destroy vegetation, the living world of water bodies. Therefore, the pH of rainwater is constantly monitored.

Ilya Leenson

The hydrogen indicator - pH - is a measure of the activity (in the case of dilute solutions reflects the concentration) of hydrogen ions in the solution, quantitatively expressing its acidity, is calculated as the negative (taken with the opposite sign) decimal logarithm of the activity of hydrogen ions, expressed in moles per liter.

pН \u003d - log

This concept was introduced in 1909 by the Danish chemist Sørensen. The indicator is called pH, according to the first letters of the Latin words potentia hydrogeni - the strength of hydrogen, or pondus hydrogenii - the weight of hydrogen.

The inverse pH value, a measure of the basicity of the solution, pOH, equal to the negative decimal logarithm of the concentration of OH ions in the solution, was somewhat less widespread:

rON \u003d - lg

In pure water at 25 ° C, the concentrations of hydrogen ions () and hydroxide ions () are the same and amount to 10 -7 mol / L, this directly follows from the autoprotolysis constant of water K w, which is also called the ionic product of water:

To w \u003d · \u003d 10 –14 [mol 2 / l 2] (at 25 ° C)

pH + pOH \u003d 14

When the concentrations of both types of ions in the solution are the same, they say that the solution has a neutral reaction. When acid is added to water, the concentration of hydrogen ions increases, and the concentration of hydroxide ions decreases accordingly, when adding a base, on the contrary, the content of hydroxide ions increases, and the concentration of hydrogen ions decreases. When\u003e they say that the solution is acidic, and when\u003e - alkaline.

PH determination

Several methods are widely used to determine the pH of solutions.

1) The hydrogen index can be approximately estimated using indicators, accurately measured with a pH meter, or determined analytically by acid-base titration.

For a rough estimate of the concentration of hydrogen ions, acid-base indicators are widely used - organic dyes, the color of which depends on the pH of the medium. The most famous indicators include litmus, phenolphthalein, methyl orange (methyl orange) and others. Indicators can exist in two differently colored forms - either in acidic or in basic. The color change of each indicator occurs in its range of acidity, usually 1–2 units (see Table 1, session 2).

To extend the working range of pH measurement, the so-called universal indicator is used, which is a mixture of several indicators. The universal indicator successively changes color from red through yellow, green, blue to violet upon transition from the acidic to alkaline region. Determination of pH by the indicator method is difficult for turbid or colored solutions.

2) The analytical volumetric method - acid-base titration - also gives accurate results for determining the total acidity of solutions. A solution of known concentration (titrant) is added dropwise to the test solution. When mixed, a chemical reaction proceeds. The equivalence point — the moment when the titrant is precisely enough to completely complete the reaction — is fixed with an indicator. Further, knowing the concentration and volume of the added titrant solution, the total acidity of the solution is calculated.

The acidity of the medium is important for many chemical processes, and the possibility of occurrence or the result of a particular reaction often depends on the pH of the medium. To maintain a certain pH value in the reaction system during laboratory tests or in production, buffer solutions are used that allow you to maintain a practically constant pH value when diluted or when small amounts of acid or alkali are added to the solution.

The hydrogen pH is widely used to characterize the acid-base properties of various biological media (Table 2).

The acidity of the reaction medium is of particular importance for biochemical reactions occurring in living systems. The concentration of hydrogen ions in a solution often affects the physicochemical properties and biological activity of proteins and nucleic acids, therefore, for the normal functioning of the body, maintaining an acid-base homeostasis is an extremely important task. Dynamic maintenance of optimal pH of biological fluids is achieved due to the action of buffer systems.

3) Using a special device - a pH meter - allows you to measure pH over a wider range and more accurately (up to 0.01 pH units) than with indicators, is convenient and highly accurate, allows you to measure the pH of opaque and color solutions, and therefore it is wide is used.

Using a pH meter, the concentration of hydrogen ions (pH) in solutions, drinking water, food products and raw materials, environmental objects and production systems for continuous monitoring of technological processes, including in aggressive environments, is measured.

the pH meter is indispensable for the hardware monitoring of pH of uranium and plutonium separation solutions, when the requirements for the correctness of the readings of the equipment without its calibration are extremely high.

The device can be used in stationary and mobile laboratories, including field, as well as clinical diagnostic, forensic, research, production, including meat, dairy and baking industries.

Recently, pH meters have also been widely used in aquarium farms, monitoring water quality in domestic conditions, agriculture (especially in hydroponics), and also for monitoring the diagnosis of health conditions.

Table 2. pH values \u200b\u200bfor some biological systems and other solutions

System (solution)
Duodenum
Gastric juice
Human blood


Muscle
Pancreatic juice


Protoplasm of cells



Small intestine

Sea water
Chicken Egg Protein
Orange juice
Tomato juice

Water is a very weak electrolyte, slightly dissociates, forming hydrogen ions (H +) and hydroxide ions (OH -),

This process corresponds to the dissociation constant:

Since the degree of dissociation of water is very small, the equilibrium concentration of undissociated water molecules with sufficient accuracy is equal to the total concentration of water, i.e. 1000/18 \u003d 5.5 mol / dm 3.
In dilute aqueous solutions, the water concentration varies little and can be considered a constant value. Then the expression of the dissociation constant of water is transformed as follows:

The constant equal to the product of the concentration of H + and OH - ions is a constant and is called ionic product of water. In pure water at 25 ºС, the concentrations of hydrogen ions and hydroxide ions are equal and amount to

Solutions in which the concentrations of hydrogen ions and hydroxide ions are the same are called neutral solutions.

So, at 25 ºС

- neutral solution;

\u003e - acidic solution;

< – щелочной раствор.

Instead of H + and OH - it is more convenient to use their decimal logarithms taken with the opposite sign; are indicated by pH and pOH:

;

The decimal logarithm of the concentration of hydrogen ions, taken with the opposite sign, is called hydrogen indicator(pH) .

Water ions in some cases can interact with ions of a dissolved substance, which leads to a significant change in the composition of the solution and its pH.

table 2

Hydrogen Index Formulas (pH)

* Values \u200b\u200bof dissociation constants ( K) are indicated in Appendix 3.

p K \u003d - lg K;

HAn is an acid; KtOH is the base; KtAn is the salt.

When calculating the pH of aqueous solutions, it is necessary:

1. Determine the nature of the substances that make up the solutions, and select a formula for calculating the pH (table 2).

2. If a weak acid or base is present in the solution, find in the manual or in Appendix 3 p K of this compound.

3. Determine the composition and concentration of the solution ( WITH).

4. Substitute the numerical values \u200b\u200bof the molar concentration ( WITH) and p K
into the calculation formula and calculate the pH of the solution.

Table 2 shows the pH calculation formulas in solutions of strong and weak acids and bases, buffer solutions and solutions of salts subjected to hydrolysis.

If only strong acid (HAn) is present in the solution, which is a strong electrolyte and almost completely dissociates into ions then the hydrogen index (pH) will depend on the concentration of hydrogen ions (H +) in a given acid and is determined by the formula (1).

If only a strong base is present in the solution, which is a strong electrolyte and almost completely dissociates into ions, then the pH (pH) will depend on the concentration of hydroxide ions (OH -) in the solution and is determined by the formula (2).

If only a weak acid or only a weak base is present in the solution, then the pH of such solutions is determined by formulas (3), (4).

If a mixture of strong and weak acids is present in the solution, then the ionization of the weak acid is practically suppressed by the strong acid, therefore, when calculating the pH in such solutions, the presence of weak acids is neglected and the calculation formula used for strong acids is used (1). The same reasoning is true for the case when a mixture of strong and weak bases is present in the solution. PH Calculations lead by the formula (2).

If a solution contains a mixture of strong acids or strong bases, then the pH is calculated using the pH formulas for strong acids (1) or bases (2), after summing up the concentrations of the components.

If the solution contains a strong acid and its salt or a strong base and its salt, then the pH depends only on the concentration of a strong acid or a strong base and is determined by the formulas (1) or (2).

If a weak acid and its salt (for example, CH 3 COOH and CH 3 COONa; HCN and KCN) or a weak base and its salt (for example, NH 4 OH and NH 4 Cl) are present in the solution, this mixture is buffer solution and pH is determined by formulas (5), (6).

If a salt is present in the solution formed by a strong acid and a weak base (hydrolyzes to a cation) or a weak acid and a strong base (hydrolyzes to an anion), a weak acid and a weak base (hydrolyzes to a cation and anion), then these salts undergo hydrolysis pH value, and the calculation is carried out according to formulas (7), (8), (9).

Example 1 Calculate the pH of an aqueous solution of NH 4 Br salt with a concentration.

Decision. 1. In an aqueous solution, a salt formed by a weak base and a strong acid is hydrolyzed according to the cation according to the equations:

In an aqueous solution, hydrogen ions (H +) remain in excess.

2. To calculate the pH, we use the formula for calculating the hydrogen index for the salt undergoing cation hydrolysis:

Weak base dissociation constant
(R K = 4,74).

3. Substitute the numerical values \u200b\u200bin the formula and calculate the hydrogen index:

Example 2 Calculate the pH of an aqueous solution consisting of a mixture of sodium hydroxide, mol / dm 3 and potassium hydroxide, mol / dm 3.

Decision.1. Sodium hydroxide (NaOH) and potassium hydroxide (KOH) are strong bases that almost completely dissociate in aqueous solutions into metal cations and hydroxide ions:

2. The hydrogen index will be determined by the sum of the hydroxide ions. To do this, we summarize the concentration of alkalis:

3. We substitute the calculated concentration in formula (2) for calculating the pH of strong bases:

Example 3 Calculate the pH of a buffer solution consisting of a 0.10 M solution of formic acid and a 0.10 M solution of sodium formate diluted 10 times.

Decision.1. Formic acid HCOOH is a weak acid, in an aqueous solution it only partially dissociates into ions, in Appendix 3 we find formic acid :

2. Sodium formate HCOONa - salt formed by weak acid and strong base; hydrolyzed by anion, an excess of hydroxide ions appears in the solution:

3. To calculate the pH, we use the formula for calculating the hydrogen parameters of buffer solutions formed by weak acid and its salt, according to the formula (5)

Substitute the numerical values \u200b\u200bin the formula and get

4. The hydrogen index of the buffer solutions does not change upon dilution. If the solution is diluted 10 times, its pH will remain equal to 3.76.

Example 4 Calculate the pH of a solution of acetic acid with a concentration of 0.01 M, the degree of dissociation of which is 4.2%.

Decision. Acetic acid is a weak electrolyte.

In a weak acid solution, the ion concentration is less than the concentration of the acid itself and is defined as a∙C.

To calculate the pH, we use the formula (3):

Example 5 To 80 cm 3 0.1 n CH 3 COOH solution was added 20 cm 3 0.2
n solution of CH 3 COONa. Calculate the pH of the resulting solution if K(CH 3 COOH) \u003d 1.75 ∙ 10 –5.

Decision.1. If the solution contains weak acid (CH 3 COOH) and its salt (CH 3 COONa), then this is a buffer solution. We calculate the pH of the buffer solution of this composition according to the formula (5):

2. The volume of the solution obtained after draining the initial solutions is 80 + 20 \u003d 100 cm 3, hence the concentration of acid and salt will be equal to:

3. The obtained values \u200b\u200bof acid and salt concentrations are substituted
into the formula

Example 6 To 200 cm 3 of a 0.1 N hydrochloric acid solution was added 200 cm 3 of a 0.2 N potassium hydroxide solution, to determine the pH of the resulting solution.

Decision.1. A neutralization reaction takes place between hydrochloric acid (HCl) and potassium hydroxide (KOH), resulting in the formation of potassium chloride (KCl) and water:

HCl + KOH → KCl + H 2 O.

2. Determine the concentration of acid and base:

According to the reaction, HCl and KOH react as 1: 1, therefore, KOH with a concentration of 0.10-0.05 \u003d 0.05 mol / dm 3 remains in excess in such a solution. Since the KCl salt does not undergo hydrolysis and does not change the pH of water, the potassium hydroxide in excess in this solution will influence the pH value. KOH is a strong electrolyte, to calculate the pH we use the formula (2):

135. How many grams of potassium hydroxide is contained in a 10 dm 3 solution, the hydrogen index of which is 11?

136. The hydrogen index (pH) of one solution is 2 and the other is 6. In 1 dm 3 of which solution is the concentration of hydrogen ions greater and by how many times?

137. Indicate the reaction of the medium and find the concentration of ions in solutions for which the pH is: a) 1.6; b) 10.5.

138. Calculate the pH of solutions in which the concentration is (mol / dm 3): a) 2.0 ∙ 10 –7; b) 8.1 ∙ 10 –3; c) 2.7 ∙ 10 –10.

139. Calculate the pH of solutions in which the concentration of ions is (mol / dm 3): a) 4.6 ∙ 10 –4; b) 8.1 ∙ 10 –6; c) 9.3 ∙ 10 –9.

140. Calculate the molar concentration of monobasic acid (HAn) in solution if: a) pH \u003d 4, α \u003d 0.01; b) pH \u003d 3, α \u003d 1%; c) pH \u003d 6,
α \u003d 0.001.

141. Calculate the pH of 0.01 n acetic acid solution in which the degree of acid dissociation is 0.042.

142. Calculate the pH of the following solutions of weak electrolytes:
a) 0.02 M NH 4 OH; b) 0.1 M HCN; c) 0.05 n HCOOH; d) 0.01 M CH 3 COOH.

143. What is the concentration of acetic acid solution, whose pH is 5.2?

144. Determine the molar concentration of formic acid (HCOOH) solution, whose pH is 3.2 ( K UNSOO \u003d 1.76 ∙ 10 –4).

145. Find the degree of dissociation (%) and a 0.1 M solution of CH 3 COOH if the dissociation constant of acetic acid is 1.75 ∙ 10 –5.

146. Calculate and pH of 0.01 M and 0.05 n solutions of H 2 SO 4.

147. Calculate the pH of the solution of H 2 SO 4 with a mass fraction of acid of 0.5% ( ρ \u003d 1.00 g / cm 3).

148. Calculate the pH of the potassium hydroxide solution if 2 dm 3 of the solution contains 1.12 g of KOH.

149. Calculate the pH of a 0.5 M solution of ammonium hydroxide. \u003d 1.76 ∙ 10 –5.

150. Calculate the pH of the solution obtained by mixing 500 cm 3 of 0.02 M CH 3 COOH with an equal volume of 0.2 M CH 3 COOK.

151. Determine the pH of the buffer mixture containing equal volumes of solutions of NH 4 OH and NH 4 Cl with a mass fraction of 5.0%.

152. Calculate the ratio of sodium acetate and acetic acid to obtain a buffer solution with pH \u003d 5.

153. In which aqueous solution the degree of dissociation is greatest: a) 0.1 M CH 3 COOH; b) 0.1 M UNS; c) 0.1 M HCN?

154. Derive the formula for calculating the pH of: a) acetate buffer mixture; b) ammonia buffer mixture.

155. Calculate the molar concentration of the HCOOH solution having a pH \u003d 3.

156. How will the pH change if diluted in half with water: a) 0.2 M HCl solution; b) 0.2 M solution of CH 3 COOH; c) a solution containing 0.1 M CH 3 COOH and 0.1 M CH 3 COONa?

157 *. A 0.1 N acetic acid solution was neutralized with a 0.1 N sodium hydroxide solution at 30% of its initial concentration. Determine the pH of the resulting solution.

158 *. To 300 cm 3 0.2 M solution of formic acid ( K \u003d 1.8 ∙ 10 –4) 50 cm 3 of a 0.4 M NaOH solution were added. The pH was measured and then the solution was diluted 10 times. Calculate the pH of the diluted solution.

159 *. To 500 cm 3 0.2 M solution of acetic acid ( K \u003d 1.8 ∙ 10 –5) 100 cm 3 of a 0.4 M NaOH solution was added. The pH was measured and then the solution was diluted 10 times. Calculate the pH of the diluted solution, write the equations of the chemical reaction.

160 *. To maintain the required pH value, the chemist prepared a solution: to 200 cm 3 of a 0.4 M solution of formic acid added 10 cm 3 of a 0.2% KOH solution ( p \u003d 1 g / cm 3) and the resulting volume was diluted 10 times. With what pH is the solution obtained? ( K HCOOH \u003d 1.8 ∙ 10 –4).

Water is a weak electrolyte; it weakly dissociates by the equation

At 25 ° С in 1 liter of water it decomposes into 10-7 mol H2O ions. The concentration of H + and OH- ions (in mol / l) will be equal to

Pure water has a neutral reaction. When acid is added to it, the concentration of H + ions increases, i.e. \u003e 10-7 mol / l; the concentration of OH– ions decreases, i.e. less than 10-7 mol / l. When alkali is added, the concentration of OH- ions increases:\u003e 10-7 mol / L, therefore, less than 10-7 mol / L.

In practice, to express the acidity or alkalinity of a solution, its negative decimal logarithm, which is called the pH value of pH, is used instead of concentration

In neutral water, pH \u003d 7. The pH values \u200b\u200band the corresponding concentrations of H + and OH- ions are given in table. 4.

Buffer solutions

Many analytical reactions are carried out at a strictly defined pH value, which should be maintained throughout the reaction time. In some reactions, the pH may change as a result of the binding or release of H + ions. To maintain a constant pH value, buffers are used.

Buffer solutions are most often mixtures of weak acids with salts of these acids or mixtures of weak bases with salts of the same bases. If, for example, a certain amount of a strong acid such as HCl is added to the acetate buffer solution consisting of acetic acid CH3COOH and sodium acetate CH3COONa, it will react with acetate ions to form slightly dissociating CH3COOH:

Thus, the H + ions added to the solution will not remain free, but will be bound by CH3COO- ions, and therefore the pH of the solution will hardly change. When an alkali solution is added to the acetate buffer solution, OH– ions will be bound by undissociated molecules of acetic acid CH3COOH:

Therefore, the pH of the solution in this case also will not change much.

Buffer solutions retain their buffer action to a certain limit, i.e. they have a certain buffer capacity. If there are more H + or OH- ions in the solution than the buffer capacity of the solution allows, then the pH will change to a large extent, as in a non-buffer solution.

Typically, analysis methods indicate which buffer solution should be used when performing this analysis and how to prepare it. Buffer mixtures with an exact pH value are released in ampoules to prepare 500 ml of solution.

pH \u003d 1.00. Composition: 0.084 g of glycocol (amino acetic acid NH2CH2COOH), 0.066 g of sodium chloride NaCl and 2.228 g of hydrochloric acid HCl.

pH \u003d 2.00. Composition: 3.215 g of citric acid C6H8O7-H2O, 1.224 g of sodium hydroxide NaOH and 1.265 g of hydrochloric acid HCl.

pH \u003d 3.00. Composition: 4.235 g of citric acid C6H8O7-H2O, 1.612 g of sodium hydroxide NaOH and 1.088 g of hydrochloric acid HCl.

pH \u003d 4.00. Composition: 5.884 g of citric acid C6H8O7-H2O, 2.240 g of sodium hydroxide NaOH and 0.802 g of hydrochloric acid HCl.

pH \u003d 5.00. Composition: 10.128 g of citric acid C6H8O7-H2O and 3.920 g of sodium hydroxide NaOH.

pH \u003d 6.00. Composition: 6.263 g of citric acid C6H8O7-H2O and 3.160 g of sodium hydroxide NaOH.

pH \u003d 7.00. Composition: 1.761 g of potassium dihydrogen phosphate KH2PO4 and 3.6325 g of sodium hydrogen phosphate Na2HPO4-2H2O.

pH \u003d 8.00. Composition: 3.464 g of boric acid H3BO3, 1.117 g of sodium hydroxide NaOH and 0.805 g of hydrochloric acid HCl.

pH \u003d 9.00. Composition: 1.546 g of boric acid H3BO3, 1.864 g of potassium chloride, KCl and 0.426 g of sodium hydroxide NaOH.

pH \u003d 10.00. Composition: 1.546 g of boric acid H3BO3, 1.864 g of potassium chloride KCl and 0.878 g of sodium hydroxide NaOH.

pH \u003d 11.00. Composition: 2.225 g of sodium hydrogen phosphate Na2HPO4-2H2O and 0.068 g of sodium hydroxide NaOH.

pH \u003d 12.00. Composition: 2.225 g of sodium hydrogen phosphate Na2HPO4-2H2O and 0.446 g of sodium hydroxide NaOH.

pH \u003d 13.00. Composition: 1.864 g of potassium chloride KCl and 0.942 g of sodium hydroxide NaOH.

Deviations from the nominal pH value reach ± \u200b\u200b0.02 for solutions at pH from 1 to 10 and ± 0.05 at pH from 11 to 13. Such accuracy is quite sufficient for practical work.

Standard pH buffers are used to adjust pH meters.

1. Acetate buffer solution with pH \u003d 4.62: 6.005 g of acetic acid CH3COOH and 8.204 g of sodium acetate CH3COONa in 1 liter of solution.

2. Phosphate buffer solution with pH \u003d 6.88: 4.450 g of sodium hydrogen phosphate Na2HPO4-2H2O and 3.400 g of potassium dihydrogen phosphate KH2PO4 in 1 liter of solution.

3. Borate buffer solution with pH \u003d 9.22: 3.81 g of sodium tetraborate Na2B4O7-10H2O in 1 liter of solution.

4. Phosphate buffer solution with pH \u003d 11.00: 4.450 g of sodium hydrogen phosphate Na2HPO4-2H2O and 0.136 g of sodium hydroxide NaOH in 1 liter of solution.

For the preparation of buffer solutions for agrochemical and biochemical analysis with pH values \u200b\u200bfrom 1.1 to 12.9 with an interval of 0.1, 7 basic stock solutions are used.

Solution 1. 11.866 g of sodium hydrogen phosphate Na2HPO4-2H2O is dissolved in water and diluted in a volumetric flask with water to 1 L (solution concentration 1/15 M).

Solution 2. Dissolve 9.073 potassium dihydrogen phosphate KH2PO4 in 1 L of water in a volumetric flask (concentration 1/15 M).

Solution 3. 7.507 g of glycol (aminoacetic acid) NH2CH2COOH and 5.84 g of sodium chloride NaCl are dissolved in 1 l of water in a volumetric flask. From this solution by mixing with 0.1 n. HCl solution prepare buffer solutions with a pH of from 1.1 to 3.5; mixing with 0.1 N. NaOH solution prepare solutions with a pH of from 8.6 to 12.9.

Solution 4. 21.014 g of C6H8O7-H2O citric acid are dissolved in water, 200 ml of 1N are added to the solution. NaOH solution and diluted to 1 liter with water in a volumetric flask. By mixing this solution with 0.1 N. HCl solution prepare buffer solutions with a pH of from 1.1 to 4.9; mixing with 0.1 N. NaOH solution prepare buffer solutions with a pH of from 5.0 to 6.6.

Solution 5. 12.367 g of boric acid H3BO3 are dissolved in water, 100 ml of 1N are added. NaOH solution and diluted with water to 1 liter in a volumetric flask. By mixing this solution with 0.1 N. HCl solution prepare buffer solutions with a pH of from 7.8 to 8.9; mixing with 0.1 N. NaOH solution prepare buffer solutions with a pH of from 9.3 to 11.0.

Solution 6. Prepare exactly 0.1 n. HCl solution;

Solution 7. Prepare exactly 0.1 n. NaOH solution; distilled water to prepare a solution is boiled for 2 hours to remove CO2. During storage, the solution is protected from CO2 from the air by a calcium chloride tube.

In some solutions, mold deposits form during storage, to prevent this, a few drops of thymol are added to the solution as a preservative. To prepare a buffer solution of the required pH, these solutions are mixed in a certain ratio (Table 5). Volume is measured using a burette with a capacity of 100.0 ml. All pH values \u200b\u200bof the buffer solutions in the table are shown at a temperature of 20 ° C.

For the preparation of the initial solutions, reagents of qualification hch are used. Sodium hydrogen phosphate Na2HPO4-2H2O is previously recrystallized twice. During the second recrystallization, the temperature of the solution should not exceed 90 ° C. The resulting preparation is slightly moistened and dried in an thermostat at 36 ° C for two days. Potassium dihydrogen phosphate KH2PO4 is also recrystallized twice and dried at 110-120 ° C. Sodium chloride NaCl was recrystallized twice and dried at 120 ° C. C6H8O7-H2O citric acid is recrystallized twice. During the second recrystallization, the temperature of the solution should not be higher than 60 ° C. Boric acid H3BO3 is recrystallized twice from boiling water and dried at a temperature not exceeding 80 ° C.

The pH is affected by the temperature of the buffer solution. In the table. Figure 6 shows the deviations of pH depending on the temperature of standard buffer solutions.

To create a given pH in the analyzed solution during complexometric titrations, buffer solutions of the following composition are used.

pH \u003d 1. Hydrochloric acid, 0.1 N. solution.

pH \u003d 2. A mixture of glycolol NH2-CH2-COOH and its hydrochloric acid salt NH2-CH2-COOH-HCl. Solid glycocol (0.2-0.3 g) is added to 100 ml of hydrochloric acid salt solution.

pH \u003d 4-6.5. Acetate mixture 1 N. a solution of sodium acetate and 1 N. acetic acid solution. The solutions are mixed before use in equal volumes.

pH \u003d 5. A mixture of a solution of 27.22 g of crystalline sodium acetate and 60 ml of 1 N. HCl solution is diluted to 1 liter with water.

pH \u003d 5.5. Acetate mixture. 540 g of sodium acetate are dissolved in water and diluted to 1 liter. To the resulting solution was added 500 ml of 1 N. acetic acid solution.

pH \u003d 6.5-8. Triethanolamine and its hydrochloric acid salt. Mix a 1 M solution of triethanolamine N (C2H4OH) 3 and a 1 M solution of HCl in equal volumes before use.

pH \u003d 8.5-9.0. Ammonium acetate mixture. 300 ml of glacial acetic acid are added to 500 ml of concentrated ammonia and diluted with water to 1 liter.

pH \u003d 9. Borate mixture. 100 ml of a 0.3 M solution of boric acid are mixed with 45 ml of 0.5 N. caustic soda solution.

pH \u003d 8-11. Ammonia is ammonium chloride. Mix 1 N. NH4OH solution and 1 N. solution of NH4Cl in equal volumes before use.

pH \u003d 10. To 570 ml of concentrated ammonia solution, 70 g of ammonium chloride are added and diluted with water to 1 liter.

pH \u003d 11-13. Caustic soda, 0.1 N solution.

In the complexometric determination of the total hardness of water, gray-brown buffer tablets are used, prepared together with the indicator (black Eriochrom T). To a water sample (100 ml), it is enough to add a few drops of sodium sulfide solution (to mask heavy metals), two buffer tablets and 1 ml of concentrated ammonia. After dissolution of the tablets, the solution turns red; it is titrated with a 0.02 M EDTA solution until a stable green color is obtained. 1 ml of a 0.02 M EDTA solution corresponds to 0.02 eq / L of water hardness. Issued in the GDR.

PH measurement

To determine the pH of solutions use special reagents - indicators, as well as instruments - pH meters (electrometric determination of pH).

Indicator determination of pH. Most often in analytical practice, the pH of solutions is determined approximately using reactive indicator paper (in the range of 0.5-2.0 pH units). Using indicator universal paper, you can determine the pH more accurately (in the range of 0.2-0.3 pH units). In the table. 7 and 8 show data on reactive and universal indicator papers.

The color transition of universal indicator paper is given in table. 8 and 9. The resulting intermediate colors are compared with the attached comparison scale and the pH of the test solution is found from it. Indicator papers can be used to determine the pH of aqueous solutions with a low salt concentration and in the absence of strong oxidizing agents. Having determined the pH using universal indicator paper with an interval of pH \u003d 1.0-11.0 or 0-12, refine the result using paper "Rifan" with a narrower pH range.

Electrometric pH measurement. This method is convenient for measuring the pH of color solutions in which it is practically impossible. For measurements, instruments are used - pH meters with a glass electrode, which is usually replaced by a hydrogen electrode. Very rarely, an antimony or chinhydron electrode is used for this purpose.

Glass electrodes are used to determine the pH of solutions containing heavy metals, oxidizing agents and reducing agents, as well as colloidal solutions and emulsions. The determination of pH with a glass electrode is based on a change in the emf an element reversible with respect to hydrogen ions.

The potential of the glass surface in contact with the acid solution depends on the pH of the solution. This property of glass is used in glass electrodes - pH indicators. The glass electrode is usually in the form of a test tube, the bottom of which is made in the form of a thin-walled glass plate or in the form of a ball with a wall thickness of not more than 0.01 mm. A buffer solution of known pH is poured into the glass electrode and placed in the test solution.

A calomel electrode is used as a reference electrode. This electrode is a vessel, at the bottom of which is mercury, connected to a chain by a platinum wire. Above mercury there is a calomel paste with KCl crystals, saturated KCl solutions and calomel (Hg2Cl2) on top. The contact of the electrode with the test solution occurs through a thin asbestos fiber. Calomel reference electrode can be used for pH measurements at a temperature not exceeding 60 ° C; Do not measure the pH of solutions containing fluorides.

The pH meter instrument is always checked and adjusted using the buffer solution whose pH is close to the pH of the test solution. For example, to measure pH in the range from 2 to 6, Zerensen buffer solution with pH \u003d 3 or 4 is prepared or a standard buffer solution with pH \u003d 4.62 is used.

In laboratory practice, a pH meter LPU-01 is used to measure pH, which is designed to determine the pH of solutions in the range from -2 to 14 with a range of 4 pH units: -2-2; 2-4; 6-10; 10-14. The sensitivity of the device is 0.01 pH. Also use a special laboratory pH meter LPS-02; pH meter type PL-U1 and portable pH meter-millivoltmeter PPM-03M1.

The industrial converter of increased accuracy is a pH meter type pH-261, which is designed to measure the pH of solutions and pulps. In the field, a pH meter pH-47M is used to measure the pH of aqueous solutions; for pH measurements of salt soil extracts - pH meter PLP-64; for milk and dairy products, use a pH meter pH-222-2. Work on pH meters is carried out according to the instructions attached to each device.

The hydrogen indicator - pH - is a measure of the activity (in the case of dilute solutions it reflects the concentration) of hydrogen ions in the solution, quantitatively expressing its acidity, is calculated as the negative (taken with the opposite sign) decimal logarithm of the activity of hydrogen ions, expressed in moles per liter.

pН \u003d - log

The inverse pH value, a measure of the basicity of the solution, pOH, equal to the negative decimal logarithm of the concentration of OH ions in the solution, was somewhat less widespread:

rON \u003d - lg

To w \u003d · \u003d 10 –14 [mol 2 / l 2] (at 25 ° C)

pH + pOH \u003d 14

PH determination

Several methods are widely used to determine the pH of solutions.

1) The hydrogen index can be approximately estimated using indicators, accurately measured with a pH meter, or determined analytically by acid-base titration.

The acidity of the medium is important for many chemical processes, and the possibility of occurrence or the result of a particular reaction often depends on the pH of the medium. To maintain a certain pH value in the reaction system during laboratory tests or in production, buffer solutions are used that can maintain a practically constant pH value when diluted or when small amounts of acid or alkali are added to the solution.

The hydrogen pH is widely used to characterize the acid-base properties of various biological media (Table 2).

The device can be used in stationary and mobile laboratories, including field, as well as clinical diagnostic, forensic, research, production, including meat, dairy and baking industries.

Recently, pH meters are also widely used in aquarium farms, for controlling water quality in domestic conditions, agriculture (especially in hydroponics), and also for monitoring the diagnosis of health conditions.

Table 2. pH values \u200b\u200bfor some biological systems and other solutions

Buffer solutions

PH measurement

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