protolytic balance. Protolytic theory of acids and bases
1. Reactions of protolysis (ionization).
These include the reactions of the interaction of an acid or base with water:
Set 1 main 2 set 2 main 1
Set 1 main.2 set 2 main. one
2. Autoprotolysis reactions associated with the transfer of a proton from one water molecule to another.
Hydrolysis reactions
CH 3 COONa + H 2 O ←→ CH 3 COOH + NaOH
CH 3 COO - + H 2 O ←→ CH 3 COOH + OH -
main 2 set 1 set 2 main 1
Acid-base reactions
NH 3 + HCl → NH 4 + + Cl -
main 2 set 1 set 2 main 1
From the point of view of analytics, the following types of reactions are distinguished:
1) with proton transfer - acid-base;
2) with electron transfer - OB reaction;
3) with the transfer of electron pairs with the formation of bonds by the donor-acceptor mechanism - complexation reactions.
2.2.2 Acidity and basicity constant. pH calculations
The ability of an acid to donate a proton, and a base to accept it (i.e., the strength of acids and bases) can be characterized by equilibrium constants,
HS - solvent
who are called acidity constants (K a ) and basicity (K b ).
Solvent activity - constant value (table data)
Positions of acid-base equilibria
and the values of the corresponding acidity and basicity constants depend on the nature of the solvent.
If the solvent is a stronger proton acceptor than water (for example, ammonia), then the strength of acids in it increases. So acids that are weak in aqueous solutions can be strong in ammonia.
The stronger the basic properties of the solvent, the more acids are leveled in it.
Similarly, the stronger the acidic properties of the solvent, the more bases are leveled in it.
When moving from a more to a less basic solvent, strong acids can be weak (eg, HCl and HClO 4 in water are strong acids, but become weak in glacial acetic acid).
pH calculation
Calculations of acid-base equilibria are used for:
1) finding the pH of the solution from known equilibrium concentrations;
2) determination of equilibrium concentrations by known pH value
pH is an important assessment for biological fluids.
It is typical for living organisms to maintain the acid-base state at a certain level. This finds expression in fairly constant pH values of biological media and the ability to restore normal pH values when exposed to protoliths.
The system that maintains protolytic homeostasis includes not only physiological mechanisms (pulmonary and renal compensation), but also physicochemical action, ion exchange, and diffusion.
In analytical chemistry, it is important to know the concentrations of all particles in a solution of an acid or base after equilibrium has been established, in particular the concentration of H + ions (pH).
- weak electrolyte
- strong electrolyte
Pure water
Pure water does not exist. Sea water contains almost all chemical elements.
Solutions of weak acids
Because 
, then
Solutions of weak bases
Solutions of strong acids
To take into account the influence of the electrostatic interaction of ions, the concept ionic strength of the solution. It depends on the concentration of the ion and its charge.
For strong electrolytes, the law of mass action is satisfied if activities are used. Activity takes into account the concentration of reagents, inter-ion interaction (ion-ion, ion-dipole, dipole-dipole, hydrogen bonds).
According to the theory of Debye and Hückel
- dependence of mobility coefficient on ionic strength
A depends on the dielectric constant of the solvent and the temperature of the system. At t=25°С A=0.512 and for a binary electrolyte
Solutions of strong bases
3.3Protolytic equilibrium in buffer solutions
In a broad sense, buffer systems are called systems that maintain a certain value of a parameter when the composition changes.
Buffer solutions can be acid-base - they maintain a constant pH value when acids or bases are introduced; redox - keep the potential of the system constant when oxidizing or reducing agents are introduced; metal buffer solutions are known.
The buffer solution is a conjugated pair; in particular, the acid-base buffer is a conjugated acid-base pair:
pH meter method
Measurements are carried out in dilute solutions, taking the activity coefficient equal to one.
If we do not take into account the reaction of water autoprotolysis, then the equation of ionic equilibria in an aqueous solution of a weak monobasic acid will have the following form:
HA + H 2 O = H 3 O + + A - x
The acidity constant is expressed as:
Moreover, [c] = 1 mol/l
If the acid is weak, then 
From here we get
Prepare solutions with different initial acid concentrations and measure their pH.
Build a graph of pH versus lg c HA. From the above equation it follows that the segment cut off by a straight line on the y-axis is equal to 1/2рK kis.
Determination of the acidity constant by the potentiometric method
For a monobasic acid
.
To determine, it is necessary to measure the concentration of hydronium ions in a solution with a known concentration of acid. As an indicator electrode, you can use a glass or quinhydrone electrode, such as Ag | AgCl | KCl || H 3 O + , sat. x.g |Pt
To obtain more accurate results, a weak acid solution is titrated with a NaOH solution, during the titration, the EMF value of the element is measured, and pH is calculated.
The following reactions take place in the system:
H 2 O + H 2 O \u003d H 3 O + + OH - x 1
HA + H 2 O \u003d H 3 O + + A - x 2
H 3 O + + NaOH \u003d 2 H 2 O + Na x 3
We can assume that x 1<< x 2 и x 1 << x 3 .
The balance equations have the form: 
.
As shown earlier
SECTION 3. KINETIC REGULARITIES OF SIMPLE REACTIONS
Chemical kinetics is a science that studies the course of a chemical reaction or physico-chemical processes over time, this is a branch of physical chemistry that studies the dependence of the rate of a chemical reaction on the concentration of reagents, temperature, properties of the medium, radiation and other factors.
Classification of chemical reactions
From the point of view of kinetics, there are several principles for classifying chemical reactions:
1) according to the state of aggregation of the participants in the reaction, all reactions are divided into homogeneous and heterogeneous.
Homogeneous reactions when all reactants are in the same phase. They are:
a) gas phase
b) liquid phase
c) solid phase
Heterogeneous reactions, when the participants in the reaction are in different phases; the reaction proceeds at the phase boundary
2) according to the specifics of the elementary act
a) catalytic
b) non-catalytic
c) photochemical
d) electrochemical
e) chain
3) by the number of stages
a) simple (stage 1)
b) complex
4) by reaction reversibility
a) reversible (bilateral)
b) irreversible
The reaction is considered irreversible if:
a) the reaction produces a gas
HCOOH → H 2 O + CO 2
b) a sparingly soluble compound is formed
AgNO 3 + KJ → AgJ↓ + KNO 3
c) a low-dissociation compound is formed
HNO 3 + NaOH → NaNO 3 + H 2 O
d) a lot of heat is released
3Fe 3 O 4 + 8Al → 4Al 2 O 3 + 9Fe + ∆H
3.2. Elementary chemical reactions
The rate of chemical reactions depends on the pathway of the reaction. This path can be represented as a sum of elementary chemical reactions.
An elementary reaction is a one-way process of converting one component into another. It is a set of the same type of elementary acts of chemical transformation. Most chemical reactions are not elementary; they include several elementary stages - complex reactions.
The reaction mechanism is a set of elementary steps.
A reactant is a participant in a chemical reaction.
d ρ n k is an infinitesimal change in the number of moles of the component k in elemental reaction ρ
If a d ρ nk > 0 – reaction product
d ρ n k< 0 – starting material
d ρ n k = 0 – indifferent substance
3.3. The rate of a chemical reaction
The rate of a chemical reaction is the number of elementary acts of a chemical transformation of the same type that occur per unit of time in a unit of volume or per unit of surface.
Consider the reaction:
t = 0 - initial numbers of moles
t ≠ 0 n A n B n C n D - current numbers of moles ξ =
(xi) ξ – reaction depth
Self-ionization of water
Water, even after repeated distillation, retains the ability to conduct electricity. This ability of water is due to its self-ionization.
$2H_2O ↔ H_3O^+ + OH^-$
The thermodynamic equilibrium constant has the form:
Picture 1.
where $a_X^(rel)=\frac(a_X^(equal))(a_X^0)$ is the relative activity of the particle $X$ in the equilibrium system;
$aX^(equal)$ - absolute activity of the particle $X$ in the equilibrium system;
$(a_x)^0$ - absolute activity $X$ in the thermodynamic state of the system.
The relative activity of water at equilibrium is practically equal to unity, since the degree of reaction is very small (if theoretically non-ionized water is taken as the standard state.
The activity coefficients of $OH^-$ and $H_3O^+$ ions will be close to unity in pure water. The equilibrium of the reaction is strongly shifted to the left. The relative activities of $OH^-$ and $H_3O^+$ are practically equal to their molar concentrations. Where
$(K_a)^0 \sim K_(auto) = $
where $ and $ are molar concentrations;
$K_(auto)$ - water autopropolis constant equal to $1.00\cdot 10^(-14) \ mol^2/l^2$ at $25^\circ \ C.$
In pure water, the concentrations of $ and $ will be equal, so
$==\sqrt(10^(-14))=10^(-7)$ for $25^\circ \ C.$
For ease of calculation, the concentration is indicated as a negative logarithm, denoted as $pH$:
$pH=-lg$
$pH$ values for pure water are $7$, in acidic solutions $pH $7$.
Acid dissociation and acidity constant
For the acid $AH$, the dissociation can be expressed by the equation:
$AH + H_2O ↔ A^- + H_3O^+$
In a state of equilibrium, the relative density of water changes insignificantly when passing from one acid to another, and with infinite dilution it approaches zero. Therefore, the thermodynamic acidity constant $K_a^0$ ($AH$) is used.
The ratio of activity coefficients is the same for all acids and is equal to one if the processes proceed in dilute solutions.
Then, in a dilute aqueous solution, the acidity constant $Ka (AH$) is used as a measure of acid strength, which can be determined by the formula:
$Ka (AH)=\frac()()$
The formula displays the molar concentration of particles at a fixed temperature $(25^\circ \ C)$ in the equilibrium state.
The higher the acidity constant, the higher the degree of dissociation, the stronger the acid. For calculations and characteristics of acidity, the negative logarithm of the acidity constant $pKa$ is used.
$pKa (AH)= -lgKa (AH)$
The larger the value of the acidity constant, the weaker the acid.
The value of the acidity constant is equal to the $pH$ value of the solution at which the acid will be ionized by half:
$pKa (AH) = pH - lg \frac()()$
The value characterizing the acidity of water molecules in an aqueous solution is:
$Ka=\frac()()=\frac(Ka_(auto))()=\frac(10^(-14))(55.5)$
Thus, at a temperature of $25^\circ C$, $pKa (H_2O) = 15.7$. This value characterizes the acidity of water molecules in solution.
For the hydroxonium ion $pKa (H_3O^+) = pK_(auto) - pKa = 14-15.7 = -1.7.$
The $pKa$ values are tabular data. However, for acids with $pKa 0$ the table data will be inaccurate.
It is possible to determine the acidity constants in water by directly measuring the concentrations of $A^-$ and $AH$ only when acid dissociation occurs at least to some extent, even barely noticeable.
If the acid is very weak, which practically does not dissociate, then the concentration of $A^-$ cannot be accurately measured. If, on the contrary, the acid is so strong that it dissociates almost completely, then it is impossible to measure the concentration of $AH$. In this case, indirect methods for determining acidity will be used.
Base ionization constant
To express the dissociation constant of a base in water, we use the equation:
$B + H_2O ↔ BH^+ + OH^-$
The basicity constant is:
$Kb=\frac()([B])$
Recently, basicity constants are practically not used in calculations, since the acidity constant of the conjugate acid can be used to obtain all the necessary information about the base $BH^+.$
$BH^+ + H_2O ↔ B + H_3O^+$
$Ka (BH^+) = \frac([B])()$
The acidity constant of an acid will be a measure of strength:
- $AH$ or $BH^+$ as proton donors;
- $A^-$ or $B$ as proton acceptors;
- a strong acid $AH$ or $BH^+$ corresponds to a weak conjugate base $A^-$ or $B$, and then $pKa$ is small or negative;
- a strong base $A^-$ or $B$ corresponds to a weak acid $AH$ or $BH^+$ and the acidity constant will be positive
It is possible to directly measure the strength of acids or bases only in a narrow range of $pKa (BH^+).$ Outside the interval, basicity will be determined by indirect methods. $pka (BH^+)$ values outside the range $-2$ to $17$ will be inaccurate.
Correlation between structure and strength of acids
The relative strength of acids can be predicted based on the nature of the central atom and the structure of the acid molecule.
The strength of the oxygen-free acid $HX$ and $H_2X$ (where $X$ is halogen) is the higher, the weaker the bond $X-H$, that is, the greater the radius of the $X$ atom.
In the series $HF - HCl - HBr - HI$ and $H_2S - H_2Se - H_2Te$, the strength of acids increases.
For oxygen-containing acids, the stronger the value of m in the compound $E(OH)nOm$, the stronger the acid.
To the equilibrium that is established in a solution of a weak electrolyte between molecules and ions, one can apply the laws of chemical equilibrium and write down the expression for the equilibrium constant. For example, for the electrolytic dissociation (protolysis) of acetic acid, proceeding under the action of water molecules,
CH 3 COOH + H 2 O ↔ H 3 O + + CH 3 COO -
the equilibrium constant has the form
There are two ways to write the value of the acidity and basicity constants. In the first method, the values of the constant and temperature are indicated on the same line after the reaction equation and a comma, for example,
HF + H 2 O ↔ H 3 O + + F - , K k \u003d 6.67 10 -4 mol l -1 (25 ° С).
In the second method, the value of the constant is first recorded, and then the acidic and basic forms of the electrolyte, the solvent (usually water) and the temperature are given in brackets:
K k \u003d 6.67 10 -4 (HF, F -, H 2 O, 25 ° C) mol l -1.
The acidity and basicity constants depend on the nature of the electrolyte, solvent, temperature, but do not depend on the concentration of the solution. They characterize the ability of a given acid or a given base to decompose into ions: the higher the value of the constant, the easier the electrolyte dissociates.
Polybasic acids, as well as bases of two- or more valent metals, dissociate in steps. Complex equilibria are established in solutions of these substances, in which ions of different charges participate. For example, the dissociation of carbonic acid occurs in two steps:
H 2 CO 3 + H 2 O ↔ H 3 O + + HCO 3 -;
HCO 3 - + H 2 O ↔ H 3 O - + CO 3 2–.
First balance - first step of protolysis- characterized by an acidity constant, denoted by K k1:
total equilibrium
H 2 CO 3 + 2H 2 O ↔ 2H 3 O + + CO 3 2 -
corresponds to the total acidity constant K to:
| K k = |
The values K k, K k1, and K k2 are related to each other by the relation:
K k \u003d K k1 K k2.
In the case of stepwise dissociation of substances, decomposition in the next step always occurs to a lesser extent than in the previous one (in the second it is less than in the first, etc.) In other words, the following inequalities are observed:
K k > K k2 > K k3 and K 01 > K 02 > K 03. . .
This is explained by the fact that the energy that must be expended to detach an ion is minimal when it is detached from a neutral molecule and becomes larger as it dissociates along each next step.
If we denote the concentration of an electrolyte decomposing into two ions through c in, and the degree of its dissociation in a given solution as α, then the concentration of each of the ions will be c in α, and the concentration of undissociated molecules c in (1 - α). Then the equation of the protolysis constant K k, ω (either the acidity constant or the basicity constant) takes the form:
This equation expresses the Ostwald dilution law. It makes it possible to calculate the degree of dissociation at various electrolyte concentrations if its dissociation constant is known. Using this equation, one can also calculate the dissociation constant of the electrolyte, knowing its degree of dissociation at a given concentration.
For solutions in which the dissociation of the electrolyte is very small, the Ostwald equation is simplified. Since in such cases α<<, то величиной α в знаменателе уравнения для К к,ω можно пренебречь. При этом уравнение принимает вид.
29. Weak electrolytes. Acidity and basicity constant. Oswald's law of dilution.
Weak electrolytes are chemical compounds whose molecules, even in highly dilute solutions, are slightly dissociated into ions that are in dynamic equilibrium with undissociated molecules. Weak electrolytes include most organic acids and many organic bases in aqueous and non-aqueous solutions.
Weak electrolytes are:
almost all organic acids and water;
some inorganic acids: HF, HClO, HClO 2 , HNO 2 , HCN, H 2 S, HBrO, H 3 PO 4 , H 2 CO 3 , H 2 SiO 3 , H 2 SO 3 and others;
some sparingly soluble metal hydroxides: Fe (OH) 3, Zn (OH) 2, etc.
Acid dissociation constant (Ka) - the equilibrium constant of the reaction of dissociation of an acid into a hydrogen ion and an anion of an acid residue. For polybasic acids, the dissociation of which takes place in several stages, they operate with separate constants for different stages of dissociation, denoting them as K a1, K a2, etc.
An example of the calculation of Diabasic acid:
More often, instead of the dissociation constant K itself, the pK value is used, which is defined as the negative decimal logarithm of the constant itself:
A base is a chemical compound capable of forming a covalent bond with a proton (Brønsted base) or with a vacant orbital of another chemical compound (Lewis base). In a narrow sense, bases are understood as basic hydroxides - complex substances, during the dissociation of which in aqueous solutions only one type of anion is split off - hydroxide ions OH-.
The Bronsted-Lowry theory makes it possible to quantify the strength of bases, that is, their ability to split off a proton from acids. This is usually done using the basicity constant Kb - the equilibrium constant of the reaction of a base with a reference acid, for which water is chosen. The higher the basicity constant, the higher the strength of the base and the greater its ability to split off a proton. Often, the basicity constant is expressed as an index of the basicity constant pKb. For example, for ammonia as a Bronsted base, one can write:
The Ostwald dilution law is a relation expressing the dependence of the equivalent electrical conductivity of a dilute solution of a binary weak electrolyte on the concentration of the solution: 
Here K is the dissociation constant of the electrolyte, c is the concentration, λ and λ∞ are the values of the equivalent electrical conductivity, respectively, at concentration c and at infinite dilution. The ratio is a consequence of the law of mass action and equality where α is the degree of dissociation.
30. Water is a weak electrolyte. Ionic product of water. Ph. POh
The ionic product of water is the product of the concentrations of hydrogen ions H+ and hydroxyl ions OH− in water or in aqueous solutions, the constant of water autoprotolysis.
Water, although a weak electrolyte, dissociates to a small extent:
The equilibrium of this reaction is strongly shifted to the left. The dissociation constant of water can be calculated by the formula:
Hydronium ion concentration (protons);
Concentration of hydroxide ions;
The concentration of water (in molecular form) in water;
The concentration of water in water, given its low degree of dissociation, is practically constant and is (1000 g/l)/(18 g/mol) = 55.56 mol/l.
At 25 °C, the dissociation constant of water is 1.8 10−16 mol/l. Equation (1) can be rewritten as:
Let us denote the product K· \u003d K в \u003d 1.8 10 -16 mol / l 55.56 mol / l \u003d 10 -14 mol² / l² \u003d (at 25 ° C).
The constant K in, equal to the product of the concentrations of protons and hydroxide ions, is called the ionic product of water. It is constant not only for pure water, but also for dilute aqueous solutions of substances. With an increase in temperature, the dissociation of water increases, therefore, Kv also increases, with a decrease in temperature, vice versa.
Hydrogen index, pH - a measure of the activity of hydrogen ions in a solution, and quantitatively expressing its acidity, is calculated as a negative (taken with the opposite sign) decimal logarithm of the activity of hydrogen ions, expressed in moles per liter:
The reciprocal pH value has become somewhat less widespread - an indicator of the basicity of the solution, pOH, equal to the negative decimal logarithm of the concentration in the solution of OH ions -:
Connecting equation:
