Ph table of various aqueous salt solutions. Calculating the pH of some solutions
1. Calculation of pH in solutions of strong acids and bases.
Calculation of pH in solutions of strong monobasic acids and bases is carried out according to the formulas:
pH \u003d - lg C to and pH \u003d 14 + lg C o
Where C to, C o is the molar concentration of an acid or base, mol / l
2. Calculation of pH in solutions of weak acids and bases
The calculation of pH in solutions of weak monobasic acids and bases is carried out according to the formulas: pH \u003d 1/2 (pK to - lgC to) and pH \u003d 14 - 1/2 (pK O - lg C O)
3. Calculation of pH in solutions of hydrolyzable salts
There are 3 cases of hydrolysis of salts:
a) hydrolysis of the salt by the anion (the salt is formed by a weak acid and a strong base, for example CH 3 COO Na). The pH value is calculated by the formula: pH = 7 + 1/2 pK to + 1/2 lg C s
b) salt hydrolysis by cation (salt is formed by a weak base and a strong acid, for example NH 4 Cl). Calculation of pH in such a solution is carried out according to the formula: pH = 7 - 1/2 pK o - 1/2 lg C s
c) hydrolysis of the salt by cation and anion (the salt is formed by a weak acid and a weak base, for example CH 3 COO NH 4). In this case, the calculation of pH is carried out according to the formula:
pH \u003d 7 + 1/2 pK to - 1/2 pK o
If the salt is formed by a weak polybasic acid or a weak multiprotonic base, then in the formulas (7-9) listed above for calculating pH, the values of pK k and pK o according to the last stage of dissociation are substituted
4. Calculation of pH in solutions of various mixtures of acids and bases
When acid and base are poured, the pH of the resulting mixture depends on the amounts of acid and base taken and their strength.
4. Buffer systems
Buffer systems include mixtures of:
a) a weak acid and its salt, for example CH 3 COO H + CH 3 COO Na
b) a weak base and its salt, for example NH 4 OH + NH 4 Cl
c) a mixture of acid salts of different acidity, for example NaH 2 PO 4 + Na 2 HPO 4
d) a mixture of acid and medium salts, for example NaНCO 3 + Na 2 CO 3
e) a mixture of basic salts of different basicity, for example Al (OH) 2 Cl + Al (OH) Cl 2, etc.
Calculation of pH in buffer systems is carried out according to the formulas: pH = pK to - lg C to / C s and pH = 14 - pK o + lg C o / C s
Acid-base buffer solutions, Henderson-Hasselbach equation. General characteristics. Operating principle. Calculation of the pH of the buffer solution. buffer capacity.
buffer solutions - systems that maintain a certain value of a parameter (pH, system potential, etc.) when the composition of the system changes.
Acid-base called buffer solution , which maintains an approximately constant pH value when not too large amounts of a strong acid or strong base are added to it, as well as when diluted and concentrated. Acid-base buffer solutions contain weak acids and their conjugate bases. A strong acid, when added to a buffer solution, "turns" into a weak acid, and a strong base into a weak base. Formula for calculating the pH of a buffer solution: pH = pK about + lg C about /WITH with This equation Henderson-Hasselbach . It follows from this equation that the pH of a buffer solution depends on the ratio of the concentrations of a weak acid and its conjugate base. Since this ratio does not change when diluted, the pH of the solution remains constant. Dilution cannot be unlimited. With a very significant dilution, the pH of the solution will change, because, firstly, the concentrations of the components will become so small that it will no longer be possible to neglect the autoprotolysis of water, and secondly, the activity coefficients of uncharged and charged particles depend differently on the ionic strength of the solution.
The buffer solution maintains constant pH when only small amounts of a strong acid or strong base are added. The ability of a buffer solution to "resist" a change in pH depends on the ratio of the concentrations of a weak acid and its conjugate base, as well as on their total concentration - and is characterized by a buffer capacity.
Buffer capacity - the ratio of an infinitesimal increase in the concentration of a strong acid or strong base in a solution (without a change in volume) to the change in pH caused by this increase (p. 239, 7.79)
In a strongly acidic and strongly alkaline environment, the buffer capacity increases significantly. Solutions in which a sufficiently high concentration of a strong acid or strong base also have buffering properties.
Buffer capacity is maximum at pH=pKa. To maintain a certain pH value, a buffer solution should be used, in which the pKa value of the weak acid included in its composition is as close as possible to this pH. It makes sense to use a buffer solution to maintain the pH in the pKa + _ 1 range. This interval is called the working force of the buffer.
19. Basic concepts related to complex compounds. Classification of complex compounds. Equilibrium constants used to characterize complex compounds: formation constants, dissociation constants (general, stepwise, thermodynamic, real and conditional concentration)
Most often, a complex is a particle formed as a result of the donor-acceptor interaction of a central atom (ion), called a complexing agent, and charged or neutral particles, called ligands. The complexing agent and ligands must exist independently in the environment where the complexation occurs.
A complex compound consists of inner and outer spheres. K3(Fe(CN)6)- K3-outer sphere, Fe-complexing agent, CN- ligand, complexing agent + ligand=inner sphere.
Dentality is the number of ligand donor centers participating in the donor-acceptor interaction during the formation of a complex particle. Ligands are monodentate (Cl-, H2O, NH3), bidentate (C2O4(2-), 1,10-phenanthroline) and polydentate.
The coordination number is the number of ligand donor centers with which a given central atom interacts. In the above example: 6-coordination number. (Ag (NH3) 2) + - coordination number 2, since ammonia is a monodentate ligand, and in (Ag (S2O3) 2) 3- - coordination number 4, since the thiosulfate ion is a bidentate ligand.
Classification.
1) Depending on their charge: anionic ((Fe(CN)6)3-), cationic ((Zn(NH3)4)2 +) and uncharged or non-electrolyte complexes (HgCl2).
2) Depending on the number of metal atoms: mononuclear and polynuclear complexes. A mononuclear complex contains one metal atom, while a polynuclear complex contains two or more. Polynuclear complex particles containing identical metal atoms are called homonuclear (Fe2(OH)2)4+ or Be3(OH)3)3+), and those containing atoms of different metals are called heteronuclear (Zr2Al(OH)5)6+).
3) Depending on the nature of the ligands: homogeneous ligand and mixed ligand (mixed ligand) complexes.
Chelates are cyclic complex compounds of metal ions with polydentate ligands (usually organic), in which the central atom is part of one or more cycles.
Constants. The strength of a complex ion is characterized by its dissociation constant, called the instability constant.
If reference data on stepwise instability constants are not available, the general instability constant of the complex ion is used:
The general instability constant is equal to the product of stepwise instability constants.
In analytical chemistry, instead of the instability constants, the stability constants of the complex ion have recently been used: 
The stability constant refers to the process of formation of a complex ion and is equal to the reciprocal of the instability constant: Kst = 1/Knest.
The stability constant characterizes the equilibrium of complex formation.
See page 313 for thermodynamic and concentration constants.
20. Influence of various factors on the process of complex formation and stability of complex compounds. Influence of the concentration of reacting substances on complexation. Calculation of the molar fractions of free metal ions and complexes in an equilibrium mixture.
1) The stability of complex compounds depends on the nature of the complexing agent and ligands. The pattern of changes in the stability of many metal complexes with various ligands can be explained with help. Theories of hard and soft acids and bases (HMCA): soft acids form more stable compounds with soft bases, and hard acids with hard ones. Ligands (l. bases), and Ag+ or Hg2+ (m. to-you) with S-sod. Ligands (m. basic). Complexes of metal cations with polydentate ligands are more stable than complexes with similar monodentate ligands.
2) ionic strength. With an increase in ionic strength and a decrease in the activity coefficients of ions, the stability of the complex decreases.
3) temperature. If, during the formation of the complex, delta H is greater than 0, then the stability of the complex increases with increasing temperature; if delta H is less than 0, then it decreases.
4) side districts. The effect of pH on the stability of complexes depends on the nature of the ligand and the central atom. If the complex contains a more or less strong base as a ligand, then with a decrease in pH, the protonation of such ligands occurs and the molar fraction of the ligand form involved in the formation of the complex decreases. The effect of pH will be the stronger, the greater the strength of the given base and the lower the stability of the complex.
5) concentration. As the ligand concentration increases, the content of complexes with a large coordination number increases and the concentration of free metal ions decreases. With an excess of metal ions in the solution, the monoligand complex will dominate.
Molar fraction of metal ions not bound into complexes
Molar fraction of complex particles
The calculation of the activity of hydrogen ions according to the equations of the law of mass action is greatly simplified if we take the negative logarithms of the quantities included in these equations. We introduce the following designations: a - activity, f - activity coefficient, C - concentration, K1 - first and K2 - second dissociation constant, Ksh - water dissociation constant. Let us denote the negative decimal logarithm of the activity of hydrogen ions as pH.
The values of pC, pf, pK for various concentrations, different ionic strengths and constants (pK) of some acids and bases are given in Table. 1-3. The value of the dissociation constant of water and its pKsh values at different temperatures are given in Table. 4.
Using table. 1-4, you can easily calculate the pH values of some solutions using the following equations (formulas 6-16):
Examples of calculations for these equations are given below. To simplify the calculations, you can avoid calculating the ionic strength of solutions to find pf, taking for monovalent electrolytes pf - 0.1, for bivalent pf - 0.2, for trivalent pf - 0.4; these values are close to those shown in Table. 3 at concentrations commonly used in laboratory practice (0.025-0.2 m). In other cases, you should use the data in Table. 3. With more approximate calculations, it is possible not to take into account the correction for pf at all.
When using equations (8)-(16), the correction for pf should be taken into account only if the value of the concentration dissociation constant is used in the calculation (in Table 2, these constants are not underlined).
The derivation of equations (6)-(16) is not given here; it can be found in physical chemistry textbooks.
Strong acids and bases and I - HCl, HNO3, HClO4, H2SO4, KOH, etc. Strong acids and bases dissociate almost completely. Therefore, the activity of hydrogen ions in solutions of strong acids according to (6) will be a = Cf, or in logarithmic form:
For strong bases, the activity of hydrogen ions in their solutions can be calculated from the equation:
the conclusion of which can be easily made taking into account the equalities:
Weak acids and bases - acetic acid, aqueous solutions of ammonia, aniline, etc. The negative logarithms of the constants of some of them, or their pK values, are given in Table. 2.
According to the law of mass action for weak acids, we have:
It should be noted that when deriving equations (8) and (9), it was assumed that the concentration of undissociated acid is equal to its total concentration, or [NA]=C. This approximate equality is valid with an accuracy of 1% only for acids whose dissociation constant K is equal to or less than 10v-3 or pK = 3. Thus, equations (8) and (9) should be used to calculate the pH of acids and bases for which pK > 3. When pK is less than 3, more complex calculations should be used, which are not presented here.
Polybasic weak acids having two or more dissociation constants - K1, K2, K3, etc., for example, carbonic, phosphoric, oxalic, etc. When calculating the pH value of their solutions, two cases should be borne in mind: 1) between the values of the first and second there is a significant difference between the constants, so that pK2-pK1 > 3, and 2) the values of the first and second constants are close to each other, so that pK2-pK1 In the second case, i.e., when pK2-pK1
Salt solutions. Depending on the composition of the salt, their solutions can be alkaline, neutral or acidic. Therefore, when determining the pH of a solution, four cases should be distinguished: 1) salts, the anion and cation of which belong to a strong base and a strong acid; 2) salts composed of a weak acid and a strong base; 3) salts composed of a strong acid and a weak base; 4) salts composed of a weak acid and a weak base.
The different reaction of aqueous solutions of salts is associated with different degrees of dissociation of strong and weak acids and bases. It was indicated above that strong electrolytes dissociate completely in aqueous solutions, and weak ones partially. Since salts dissociate completely in aqueous solutions, and weak acids and bases partially, if there are anions or cations of weak acids and bases in the solution, the latter are partially hydrolyzed, pass into relatively slightly dissociated compounds, and a free strong acid or alkali appears in the solution. For example, when the salt NH4Cl is dissolved in water (NH4 is a cation of a weak base, Cl is an anion of a strong acid), the NH4 ion is partially hydrolyzed and it passes into a relatively slightly dissociated compound NH4OH, and free HCl appears in the solution:
A solution of such a salt will be acidic. On the contrary, when a salt of a weak acid and a strong base is dissolved in water, alkalinization of the solution occurs for the same reasons:
Obviously, when calculating the pH value of salt solutions, one should take into account the degree of dissociation of the weak acids or bases that make up the salt, or the value of their dissociation constants.
1. Salts of a strong base and a strong acid: KCl, NaCl, KNO3, etc. Solutions of these salts should have a reaction close to neutral, since neither anion nor cation give poorly dissociated compounds with water. In practice, due to negligible buffering, the presence of contaminants, as well as dissolved carbon dioxide, solutions of these salts have pH values that differ from 7 by up to one and sometimes even more. Thorough purification of salts by recrystallization and removal of CO2 from their solutions bring the reaction of their solutions closer to neutral.
2. Salts of a strong base and a weak acid - CH3COONa, etc. In this case, the value of the constant of a weak acid should be taken into account, and the pH value of the solutions can be calculated using equation (11):
3. Salts of a weak base and a strong acid NH4Cl, NH4NO3, (NH4)2SO4, etc. You can find the pH value of solutions of these salts using equation (12):
4. Salts of a weak base and a weak acid - CH3COONH4, NH4NO2, etc. The pH value of such salts does not depend on their concentration and can be found by equation (13):
Acid salts. The reaction of acid salts such as NaHCO3, KHSO3, potassium acid tartrate [K (C4H6O6)] and others in cases where their concentration in the solution exceeds at least 100 times the value of the first dissociation constant (i.e., at pK1-pK2), does not depend on concentration and can be calculated from equation (14):
In cases where the concentration of the acid salt exceeds the value of the first constant (or the second constant for disubstituted salts) by less than 100 times (or at pK1-рС
The calculation according to this equation is complicated by the fact that before it is carried out, it is necessary to calculate the value of the sum K1 + C, and then find it according to Table. 1 p(K1+C) value. The values of dissociation constants are given in table. 2. For example, for phosphoric acid K1=7.5 10v_3, or 0.0075. Then for 0.01 n. solution we get: K1+C = 0.0075+0.01 = 0.0175, a p(K1+C) = 1.8.
Consider the method of calculating pH in solutions of hydrolyzable salts. The equilibrium concentrations of participants in the hydrolysis process for each stage (water is not taken into account) are related to each other through the corresponding hydrolysis constant Kg, which is calculated according to the following rule:
the hydrolysis constant K g is equal to the quotient of the division of the ionic product of water K w \u003d 10 -14 by the dissociation constant K a of a weak acid (K b of a weak base), which were formed as a result of hydrolysis.
Example. Weak phosphoric acid dissociates in three steps:
H 3 PO 4 → H 2 PO 4 – → HPO 4 2– → PO 4 3–
dissociation constants Ka 1 Ka 2 Ka 3
in this case, three types of anions capable of hydrolysis are formed.
The PO 4 3– ion is hydrolyzed in three steps and each of them has its own hydrolysis constant: PO 4 3- → HPO 4 2- → H 2 PO 4 - → H 3 PO 4. Hydrolysis constant: K g1 \u003d K w / Ka 3 ; K g2 \u003d K w / Ka 2; K g3 \u003d K w / Ka 1.
As an example, consider the simplest case of calculating pH in a solution of an average salt, for example, sodium phosphate with a molar concentration of 0 mol/l.
Na 3 PO 4 → 3Na + + PO 4 3–
Let us denote the degree of hydrolysis of the PO 4 3– ion in the first step as h 1 (h 1<<1), тогда к моменту установления равновесия подверглось гидролизу с гидр. (PO 4 3–) = h 1 ·с 0 и
PO 4 3– + H 2 O ↔ HPO 4 2– + OH –
before hydrolysis with 0 mol/l ½ ê - ½ -
equilibrium = with 0 - with hydr. =½ ê = h 1 s 0 ½ = h 1 s 0
C 0 (1 - h 1)
The hydrolyzing ion PO 4 3– was formed by the third stage of the dissociation of phosphoric acid, therefore K g1 = ,
whence h 1 = , = h 1 С 0 = and pOH = - lg = =(), and pH = 14 - pOH.
Note that the use of the simplified formula (1<< h 1) возможно, если константа гидролиза К г < 10 –3 , концентрация иона с 0 >0.001 mol/l; otherwise, calculations should be carried out according to the general formula.
The calculation of pH in acid salt solutions is more complicated, since the anion can participate in two competing processes, hydrolysis and dissociation. However, one can easily determine the nature of the solution by comparing the equilibrium constants of these processes, and the one with the larger constant prevails.
Example- in a solution of sodium hydrogen phosphate Na 2 HPO 4 → 2 Na + + HPO 4 2–, and the HPO 4 2– ion can further
a) dissociate according to stage III HPO 4 2– ↔ PO 4 3– + H + ; K 3 (H 3 RO 4) \u003d 1.26 10 -12
b) hydrolyze HPO 4 2– + H 2 O ↔ H 2 PO 4 - + OH - K g \u003d K w / K 2 (H 3 RO 4) \u003d
10 -14 / 6.34 10 -8 \u003d 1.57 10 -7.
It can be seen that the HPO 4 2– hydrolysis process predominates and the solution of this salt is weakly alkaline.
LEARNING TASKS
1. Calculate the hydrolysis constant K g, the degree of hydrolysis h and the pH of the ammonium chloride solution with the salt concentration c(NH 4 Cl)=0.01 mol/dm 3 .
1) salt NH 4 Cl is formed by a strong acid HCl and a weak base NH 4 OH - hydrolysis by cation; salt hydrolysis is a reversible process.
NH 4 Cl + H 2 O \u003d NH 4 OH + HCl
NH 4 + + H 2 O Û NH 4 OH + H + - as a result of hydrolysis, H + ions are formed, i.e. the solution is acidic.
2) hydrolysis constant K g calculated by the formula:
,
where K w is the ionic product of water, K w\u003d 10 -14 (25 0 C); K b(NH 4 OH) - ionization constant of the base (reference value), K b(NH 4 OH)=1.74 10 -5 .
3) the degree of hydrolysis h of the salt is calculated by the formula:
where c o is the molar concentration of salt in the solution.
4) the concentration of H + ions is equal to the concentration of the hydrolyzed part of the salt and is determined by the formula:
Answer: the hydrolysis constant of the NH 4 Cl salt is 5.75 10 -10; the degree of hydrolysis was 2.4·10 -4 ; The pH of the solution is 5.62.
2. Determine the hydrolysis constant, the degree of hydrolysis and the pH of the potassium acetate solution, if the concentration with(CH 3 SOOK) \u003d 0.1 mol / dm 3, and K a(CH 3 COOH)=1.8·10 -5.
1) salt CH 3 COOK is formed by a weak acid CH 3 COOH and a strong base KOH - hydrolysis by anion, the medium is alkaline as a result of hydrolysis:
CH 3 COOK + H 2 O \u003d CH 3 COOH + KOH;
СH 3 COO - + H 2 O Û CH 3 COOH + OH - - OH - ions accumulate, the environment is alkaline.
2) hydrolysis constant K g calculated by the formula:
,
where K a is the ionization constant of the acid. 
2) the degree of hydrolysis h of the salt is calculated by the equation:
where с 0 is the salt concentration in the solution.
3) the concentration of OH - - ions is equal to the concentration of the hydrolyzed part of the salt.
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1) NaCl pH = 7
2)NH 4 Cl pH<7
3) CH 3 COONa pH > 7
| 1 | 2 | 3 | 4 | 5 | 6 | 7 | 8 | 9 | 10 | 11 | 12 | 13 | 14 |
| 4 | 4 | 5 | 1 | 2 | 3 | 1 | 3 | 4 | 4 | 2 | 2 | 1 | 2 |
| 15 | 16 | 17 | 18 | 19 | 20 | 21 | 22 | 23 | 24 | 25 | 26 | 27 | 28 |
| 4 | 2 | 4 | 4 | 1 | 2 | 3 | 2 | 4 | 4 | 1 | 2 | 4 | 3 |
| 29 | 30 | 31 | 32 | 33 | 34 | 35 | 36 | 37 | 38 | 39 | 40 | 41 | 42 |
| 2 | 1 | 2 | 2 | 3 | 4 | 1 | 1 | 1 | 1 | 1 | - | 4 | 4 |
| 43 | 44 | 45 | 46 | 47 | 48 | 49 | 50 | 51 | 52 | 53 | 54 | 55 | 56 |
| 1 | 4 | 1 | 3 | 4 | 3 | 4 | 4 | 3 | 4 | 3 | 4 | 2 | 2 |
| 57 | 58 | 59 | 60 | ||||||||||
| 1 | 2 | 2 | - |
BUFFER SOLUTIONS, BUFFER SYSTEMS OF THE ORGANISM
1. pH value of buffer solutions when adding small amounts of acids and bases:
1) remain constant, because the added hydrogen cations and hydroxide anions are bound by the proton acceptors and donors of the buffer system, respectively;
2) They remain approximately constant as long as the concentrations of the components of the buffer systems exceed the concentrations of the added ions;
3) Change, because the concentrations of acids and bases in the system change;
2. pH values of buffer solutions when diluted…
1) remain constant, since the ratio of the concentrations of the components of the buffer systems does not change;
2) remain approximately constant up to certain concentrations;
They change because the concentration of system components decreases.
3. Which of the following conjugated acid-base pairs have buffer properties: a) HCOO - /HCOOH; b) CH 3 COO - /CH 3 COOH; c) Cl - /HCl; d) HCO - 3 /CO 2; e) H 2 PO - 4;
1) all;
2) a, b, d, e;
3) b, d, e;
4. From the listed conjugate acid-base pairs, select systems that have buffer properties: a) H 3 PO 4 /H 2 PO - 4; b) H 2 PO 4 /H 2 PO 2- 4; c) HPO 2- 4 /PO 3- 4 ; d) HNO 3 /NO - 3; e) HCOOH/HCOO - .
1) all;
2) b, e;
3) a, b, c, e;
B, c, d.
5. Which of the acid-base pairs have buffer properties; a) Hb-/HHb; b) HbO 2 /HhbO 2;; c) HSO - 4 / H 2 SO 4; d) NH + 4 /NH 4 OH; e) NO - 3 / HNO 3?
1) all;
2) a, b, c, d;
3) a, b, c;
6. Which of the acid-base pairs have buffer properties: a) Cl - / HCl; b) NO - 3 / HNO 3; c) HSO - 4 / H 2 SO 4; d)CH 3 COO - /CH 3 COOH; e) NH + 4 / NH 4 OH?
1) all;
2) a, b, c;
3) d, e;
C, d, d.
7. Which of the conjugated acid-base pairs have buffer properties: a) HCOO - / HCOOH; b) HPO 2- 4 / H 2 PO 4; c) H 3 PO - 4; d) HCO - 4 / CO 2?
1) all;
2) a, b;
3) b, c, d;
8. Which of the buffer systems contain only salts in their composition: a) CO 2- 3 / HCO - 3; b) HCO - 3 /CO 2; c) HPO 2- 4 / H 2 PO - 4; d) H 2 PO - 4 /H 3 PO 4; e) HCOO - /HCOOH; e) PO 3- 4 /HPO 2- 4 .
1) a, c, d;
2) a, c, e;
3) a, b, c, d, f;
A, b, c.e.
9. Buffer solutions include mixtures: a) NaH 2 PO 2 + Na 2 HPO 4; b) H 3 PO 4 + NaH 2 PO 4; c) Ns 2 CO 3 + NaHCO 3; d) Na 2 HPO 4 + Na 3 PO 4.
1) All;
2) a, b;
3) c, d;
A B C.
10. When adding HCl in the buffer system HPO 2- 4 /H 2 PO - 4:
11. When NaOH is added to the HPO 2-4 / H 2 PO - 4 buffer system:
1) active concentration (HPO 2-4) increases, (H 2 PO - 4) - decreases.
2) active concentration (HPO 2-4) decreases, (H 2 PO - 4)) increases.
The activity of the components does not change.
12. When NaOH is added to the NH + 4 /NH 3 H 2 O buffer system:
The activity of the components does not change.
13. When adding HCl in the buffer system NH + 4 /NH 3 H 2 O:
1) the active concentration (NH + 4) increases, (NH 3 H 2 O) - decreases.
2) the active concentration (NH + 4) decreases, (NH 3 H 2 O)) increases.
The activity of the components does not change.
14. When NaOH is added to the CH 3 COO - / CH 3 COOH buffer system:
The activity of the components does not change.
15. When adding HCl in the buffer system CH 3 COO - / CH 3 COOH:
1) active concentration (CH 3 COOH) increases, (CH 3 COO -) - decreases.
2) active concentration (CH 3 COOH) decreases, (CH 3 COO -)) increases.
The activity of the components does not change.
16. The maximum buffer capacity of the system is at:
1) pH= p K a;
2) pH> p K a;
3) pH<p K a;
These parameters are not related to each other.
17. The acid-base conjugate pair has the maximum buffer capacity at physiological pH:
1) H 3 PO 4 / H 2 PO - 4 ( pK a(H 3 PO 4) = 2.1;
2) H 3 PO - 4 / H 2 PO 2- 4 ( pK a(H 2 PO - 4) = 6.8;
3) HPO 2- 4 /PO 3- 4 ( pK a(HPO - 4) = 12.3;
18. At the same concentrations of components, the buffer capacity:
1) maximum, because pH = p K a;
2) maximum, because pH> p K a;
3) minimum, because pH = p K a;
19. Buffer capacity when diluting solutions:
1) decreases due to a decrease in the concentration of all components of the system;
2) increases, since the degree of dissociation of electrolytes increases;
3) does not change, because the ratio of component concentrations remains constant;
Practically does not change, because the number of system components remains unchanged.
20. Buffer systems maintain balance in the body:
1) acid-base;
2) redox;
3) heterogeneous;
Ligand-exchange.
21. Acidosis is:
