At With an increase in temperature, the rate of most chemical reactions increases significantly, and for homogeneous reactions, when heated for every ten degrees, the reaction rate increases by 2-4 times.
The total number of particles in the system (N) is equal to the area under the curve. The total number of particles with energies greater than Ea is equal to the shaded area.
Figure 2 shows that as the temperature increases, the energy distribution of particles changes in such a way that the proportion of particles with higher energy increases. Thus, an important concept for a chemical reaction is the activation energy.
Activation energy is the energy that particles must have in order for their interaction to lead to a chemical reaction. The activation energy is expressed in kJ/mol. For reactions proceeding at a noticeable rate, the activation energy does not exceed 50 kJ/mol (for ion exchange reactions, Ea » 0); if Ea > 100 kJ/mol, then the reaction rate is immeasurably low.
In 1889, S. Arrhenius gave an equation for the dependence of the rate constant of a chemical reaction on temperature:
k = Ae - Ea/RT
where, A - pre-exponential factor depending on the nature of the reacting substances;
R- gas constant \u003d 8.314 J / (mol? K);
Ea- activation energy.
It follows from the Arrhenius equation that the higher the activation energy, the more it is necessary to increase the temperature to maintain the required reaction rate.
Figure 3 shows the dependence of the change in the potential energy of the reacting system on the path of the reaction. From the above figure it can be seen that for an exothermic reaction (going with the release of heat), the loss of active molecules is replenished due to the energy released during the reaction. In the case of an endothermic reaction, heat is required to maintain the desired reaction rate.
| exothermic reaction | Endothermic reaction |
Figure 10.3 Energy diagram of a chemical reaction
A - reactants, C - products.
2.4 Influence of foreign matter
Foreign substances, depending on the impact, can accelerate reactions - catalysts or slow down - inhibitors.
Catalysts- These are substances that accelerate chemical reactions, but after the reaction they themselves remain unchanged.
Inhibitors - these are substances that slow down the reaction. In practice, sometimes it is necessary to slow down the reactions (corrosion of metals, etc.) this is achieved by introducing inhibitors into the reaction system. For example, sodium nitrite, chromate and potassium dichromate reduce the corrosion rate of metals.
promoters- substances that increase the activity of the catalyst. In this case, the promoters themselves may not have catalytic properties.
Catalytic poisons- foreign impurities in the reaction mixture, leading to partial or complete loss of catalyst activity. So, traces of arsenic, phosphorus cause a rapid loss of activity of the catalyst V 2 O 5 in the contact method for obtaining H 2 SO 4 .
3. Chemical equilibrium
In chemical reactions, the starting materials are not always completely converted into reaction products. This is because as the reaction products accumulate, conditions can be created for the reverse reaction to occur. Most chemical reactions are reversible.
As an example, let us analyze the reversible reaction of the synthesis of ammonia from nitrogen and hydrogen, which is extremely important for industry:
direct reaction -2N2 + 3H2 →2NH 3 ,
reverse reaction - 2NH 3 →N 2 + 3H2,
reversible reaction - 2N 2 + 3H 2« 2NH3.
The forward and reverse reactions are separate reactions with their corresponding kinetic equations, pre-exposure factors, activation energies, etc.
An important quantitative characteristic of reversible reactions is the equilibrium constant, which is determined when the system reaches chemical equilibrium - a state in which the rates of direct and reverse reactions are equal. Examples of the application of the law of mass action (p.m.m.).
Let us derive the equilibrium constant using the example of the ammonia synthesis reaction.
Kinetic equation of direct reaction
N 2 + 3H 2 →2NH 3
has the form Vpr \u003d Kpr 3.
Kinetic equation of the reverse reaction
2NH3 →N 2 + 3H2
has the form Vrev \u003d Cobr 2.
In a state of chemical equilibrium, Vpr = Var.
Substituting in the condition of chemical equilibrium the expressions for the rates of direct and reverse reactions, we obtain the following equality Kpr 3 = Cobr 2.
After transformation we get
.
4. Le Chatelier's principle
If any external influence is exerted on a system that is in a state of chemical equilibrium, then the equilibrium will shift as a result of the processes occurring in the system in such a way that the impact will decrease.
4.1 Effect of concentration changes on equilibrium
With an increase in the concentration of any of the substances participating in the reaction, the equilibrium shifts towards the consumption of this substance, and when it decreases, towards the formation of this substance.
Example 1 If in an equilibrium system
2N2 + 3H2« 2NH3
add N 2 or H 2, then, in accordance with the Le Chatelier principle, to reduce the concentrations of these substances, the equilibrium should shift to the right, the output of NH 3 will increase. As the concentration of NH 3 increases, the equilibrium will correspondingly shift to the left.
4.2 Effect of pressure change on equilibrium
The pressure in a closed reaction system is due to the presence of gaseous substances in it: the more of them, the greater the pressure. Therefore, a change in external pressure will affect the equilibrium only in cases where gaseous substances participate in it, and their number in the forward and reverse reactions is different.
If the pressure is increased in a system in a state of chemical equilibrium, then a reaction will predominantly occur, as a result of which the amount of gaseous substances decreases; when the pressure decreases, the reaction predominantly proceeds, as a result of which the amount of gaseous products increases.
Example 1 Is it possible to increase the yield of products in the reaction by changing the pressure CO 2 (g) + H 2 (g)« CO(g) + H 2 O(g).
Solution: The reaction mixture includes gaseous reagents, but their quantity does not change in the reaction: from one mole of CO 2 (g) and one mole of H2 (g), one mole of CO (g) and H 2 O (g) are obtained each. For this reason, a change in pressure does not affect the state of equilibrium.
Example 2 How will the equilibrium concentrations of reagents change with increasing pressure in the system N 2 + 3H 2 "2NH 3?
It can be seen from the reaction equation that from 4 mol of the gas of the initial products, 2 mol of the gas of the reaction products is formed. Thus, with an increase in pressure, the equilibrium will shift to a direct reaction, since it leads to a decrease in pressure.
4.3 Effect of temperature change on chemical equilibrium
Most chemical reactions proceed with the release or absorption of heat. In the first case, the temperature of the mixture increases, in the second it decreases.
If the reaction mixture, which is in a state of chemical equilibrium, is heated, then, in accordance with the Le Chatelier principle, the reaction should proceed predominantly, as a result of which heat will be absorbed, i.e. endothermic reaction; when the mixture is cooled, a predominantly reaction should proceed, as a result of which heat will be released, i.e. endothermic reaction.
If the temperature is increased in a system that is in a state of chemical equilibrium, then the equilibrium shifts towards an endothermic reaction, and when the temperature decreases, towards an exothermic reaction.
Example: 2N 2 + 3H 2« 2NH3,H0 = - 92 kJ
The reaction is exothermic, therefore, as the temperature increases, the equilibrium shifts to the left, and as the temperature decreases, it shifts to the right.
It follows from this that to increase the yield of ammonia, the temperature must be lowered. In practice, a temperature of 500 0C is maintained, since at a lower temperature the rate of the direct reaction sharply decreases.
Chemical equilibrium has a dynamic character: forward and reverse reactions do not stop at equilibrium.
The equilibrium constant depends on the temperature and the nature of the reactants. The larger the equilibrium constant, the more the equilibrium is shifted towards the formation of direct reaction products
Le Chatelier's principle is universal, as it is applicable not only to purely chemical processes, but also to physical and chemical phenomena, such as crystallization, dissolution, boiling, and phase transformations in solids.
An increase in temperature speeds up all chemical reactions. Initially, van't Hoff experimentally found that when increase in temperature for every 10 degrees, the speed increases by 2 ¸ 4 times ( Van't Hoff's rule ). This corresponds to the power-law dependence of velocity on temperature:
where T > T 0, g - van't Hoff temperature coefficient.
However, this equation is not theoretically justified. ; experimental data are better described by an exponential function (Arrhenius equation):
,
where A is a pre-exponential factor independent of T, E a is the activation energy of a chemical reaction (kJ/mol), R is the universal gas constant.
The Arrhenius equation is usually written for the rate constant:
.
This equation is theoretically substantiated by the methods of statistical physics. Qualitatively, this justification is as follows: since reactions proceed as a result of random collisions of molecules, these collisions are characterized by an almost continuous set of energies from the smallest to the very largest. Obviously, a reaction will only occur when the molecules collide with enough energy to break (or significantly stretch) some chemical bonds. For each system, there is an energy threshold E a, starting from which the energy is sufficient for the reaction to proceed - this mechanism corresponds to curve 1 in Figure 5.1. Since collisions occur with a frequency that depends on temperature according to an exponential law, formulas 5.9 and 5.10 are obtained. Then the pre-exponential factors A and k 0 represent some characteristic of the total number of collisions, and the term is the fraction of successful collisions.
The analysis of experimental data is carried out using the logarithmic form of the Arrhenius equation:
.
The graph is built in the so-called Arrhenius coordinates
(ln k - ), fig. 7.2; from the graph find k o and E a.
In the presence of experimental data for two temperatures k o and E a, it is easy to theoretically find:
; 
;
The rate of a chemical reaction largely depends on the activation energy. For the vast majority of reactions, it lies in the range from 50 to 250 kJ/mol. Reactions for which
E a > 150 kJ/mol, practically do not leak at room temperature.
Example 1 The complex irreversible reaction 2N 2 O 5 \u003d 4NO 2 + O 2 is a first-order reaction. How will its speed change when the pressure is increased by 5 times?
Solution. The kinetic equation of this reaction in general form: V = k · a . Since the reaction is complex, it is possible that a ¹ 2. By condition, the order of the reaction
a = 1. For gas reactions, the role of concentration is played by pressure. That's why
V = kP, and if Р 1 = 5Р, then V 1 /V = 5, i.e. speed increases five times.
Find the rate constant, the orders of the reactants and write down the kinetic equation.
Solution. The general kinetic equation for the rate of this reaction is:
V = k a b .
The data in the table make it possible to find the reaction orders for NO (a) and H 2 (b) by lowering the reaction order, i.e. analyzing experiments in which one of the reagents has a constant concentration. So, = 0.01 in the first and second columns, while changing.
. (private order in H 2).
For the second and third columns, on the contrary, it is the same, but - are different, therefore:
(private order for NO).
Since a and b coincide with stoichiometric coefficients, the reaction can be simple. The rate constant can be found from each column's data:
Thus, the kinetic equation is: V = 2.5. 10 3 2 .
The total (general) order of this reaction (a + b) is 3.
Example 3 The reaction rate A + 3B = AB 3 is determined by the kinetic equation V = k[A]·[B]. Determine the general order of the reaction. Is this reaction simple or complex? How many times will the reaction rate increase when the concentration is increased by 3 times?
Solution. The reaction order is determined by the sum of the exponents of the reactants in the kinetic equation. For this reaction, the overall order is two (1 + 1).
If this reaction were simple, then according to the law of mass action
V = k[A] 1 . [B] 3 and the total order would be (1+ 3) = 4, i.e. the exponents in the kinetic equation do not coincide with the stoichiometric coefficients, therefore, the reaction is complex and takes place in several stages.
With an increase in the concentrations of reagents by 3 times: V 1 = k·3[A]·3[B] = 3 2 V, that is, the speed will increase by 3 2 = 9 times.
Example 4 Determine the activation energy of the reaction and its temperature coefficient, if at 398 and 600 0 C the rate constants are, respectively, 2.1×10 -4 and 6.25×10 -1 .
Solution. E a for two values can be calculated using the formula 5.12 :
192633 J/mol.
The temperature coefficient is found from expression (5.8), because Vµk:
.
Catalysis
One of the most common methods in chemical practice for accelerating chemical reactions is catalysis. A catalyst is a substance that repeatedly participates in the intermediate stages of a reaction, but leaves it chemically unchanged.
For example, for the reaction A 2 + B 2 \u003d 2AB
the participation of catalyst K can be expressed by the equation
A 2 + K + B 2 ® A 2 .... K + B 2 ® A 2 ... K ... B 2 ® 2AB + K.
These equations can be represented by potential energy curves (Fig. 5.2.).
Rice. 5.2. Energy scheme of the reaction
with and without catalyst
Figure 5.2 shows that:
1) the catalyst reduces the activation energy by changing the reaction mechanism - it proceeds through new stages, each of which is characterized by a low activation energy;
2) the catalyst does not change the DH of the reaction (as well as DG, DU, and DS);
3) if the catalyzed reaction is reversible, the catalyst does not affect the equilibrium, does not change the equilibrium constant and the equilibrium concentrations of the system components. It speeds up both the forward and reverse reactions equally, thereby speeding up the time to reach equilibrium.
Obviously, in the presence of a catalyst, the activation energy of the reaction decreases by the value DE k. Since in the expression for the reaction rate constant (Equation 5.10) the activation energy is included in the negative exponent, even a small decrease in E a causes a very large increase in the reaction rate: 
.
The effect of the catalyst on the decrease in Еа can be shown by the example of the decomposition reaction of hydrogen iodide:
2HI \u003d H 2 + I 2.
Thus, for the reaction under consideration, the decrease in energy
activation by 63 kJ, i.e. 1.5 times, corresponds to an increase in the reaction rate at 500 K by more than 10 6 times.
It should be noted that the pre-exponential factor of the catalytic reaction k 0 1 is not equal to k 0 and is usually much less, however, the corresponding decrease in the rate does not compensate for its increase due to Еа.
Example 5 The activation energy of a certain reaction in the absence of a catalyst is 75.24 kJ / mol, and with a catalyst - 50.14 kJ / mol. How many times does the reaction rate increase in the presence of a catalyst if the reaction proceeds at 25 0 C, and the pre-exponential factor in the presence of a catalyst decreases by 10 times.
Solution. Let us denote the activation energy of the reaction without a catalyst as E a, and in the presence of a catalyst - through Ea 1 ; the corresponding reaction rate constants will be denoted by k and k 1 . Using the Arrhenius equation (5.9) (see section 5.3) and assuming k 0 1 /k 0 = 10, we find:
From here 
We finally find:
Thus, a decrease in the activation energy by the catalyst by 25.1 kJ led to an increase in the reaction rate by a factor of 2500, despite a 10-fold decrease in the pre-exponential factor.
Catalytic reactions are classified by the type of catalysts and by the type of reactions. So, for example, according to the state of aggregation of catalysts and reagents, catalysis is divided into homogeneous(catalyst and reactant form one phase) and heterogeneous(the catalyst and the reagents are in different phases, there is a phase boundary between the catalyst and the reagents).
An example of homogeneous catalysis would be the oxidation of CO to CO 2 with oxygen in the presence of NO 2 (catalyst). The mechanism of catalysis can be represented by the following reactions:
CO (g) + NO 2 (g) ® CO 2 (g) + NO (g),
2NO (g) + O 2 (g) ® 2NO 2 (g);
and the catalyst (NO 2) again participates in the first reaction.
Similarly, the oxidation of SO 2 to SO 3 can be catalyzed; a similar reaction is used in the production of sulfuric acid by the "nitrous" process.
An example of heterogeneous catalysis is the production of SO 3 from SO 2 in the presence of Pt or V 2 O 5:
SO 2 (g) + O 2 (g) ® SO 3 (g).
This reaction is also used in the production of sulfuric acid (the "contact" method).
The heterogeneous catalyst (iron) is also used in the production of ammonia from nitrogen and hydrogen and in many other processes.
The efficiency of heterogeneous catalysts is usually much greater than that of homogeneous ones. The rate of catalytic reactions in the case of a homogeneous catalyst depends on its concentration, and in the case of a heterogeneous one, on its specific surface area (that is, dispersion) - the larger it is, the greater the rate. The latter is due to the fact that the catalytic reaction takes place on the surface of the catalyst and includes the stages of adsorption (sticking) of reactant molecules on the surface; after the completion of the reaction, its products are desorbed. To increase the surface area of the catalysts, they are crushed or obtained by special methods, in which very fine powders are formed.
The examples given are also examples redox catalysis. In this case, transition metals or their compounds (Mn 3+ , Pt, Au, Ag, Fe, Ni, Fe 2 O 3, etc.) usually act as catalysts.
In acid-base catalysis the role of the catalyst is performed by H + , OH - and other similar particles - carriers of acidity and basicity. So the hydrolysis reaction
CH 3 COOCH 3 + H 2 O CH 3 COOH + CH 3 OH
accelerates by about 300 times with the addition of any of the strong acids: HCl, HBr or HNO 3 .
Catalysis is of great importance in biological systems. In this case, the catalyst is called enzyme. The efficiency of many enzymes is much greater than conventional catalysts. For example, for the reaction of nitrogen binding to ammonia
N 2 + 3H 2 \u003d 2NH 3
In industry, a heterogeneous catalyst is used in the form of sponge iron with the addition of metal oxides and sulfates.
In this case, the reaction is carried out at T » 700 K and P » 30 MPa. The same synthesis takes place in the nodules of leguminous plants under the action of enzymes at ordinary T and P.
Catalytic systems are not indifferent to impurities and additives. Some of them increase the efficiency of catalysis, such as in the above example of catalysis of the synthesis of ammonia by iron. These catalyst additives are called promoters(potassium and aluminum oxides in iron). Some impurities, on the contrary, suppress the catalytic reaction ("poison" the catalyst), this catalytic poisons. For example, the synthesis of SO 3 on a Pt catalyst is very sensitive to impurities containing sulfide sulfur; sulfur poisons the surface of the platinum catalyst. Conversely, the catalyst based on V 2 O 5 is insensitive to such impurities; the honor of developing a catalyst based on vanadium oxide belongs to the Russian scientist G.K. Boreskov.
where g - ttemperature coefficient, taking values from 2 to 4.
The explanation of the dependence of the reaction rate on temperature was given by S. Arrhenius. Not every collision of reactant molecules leads to a reaction, but only the strongest collisions. Only molecules with an excess of kinetic energy are capable of a chemical reaction.
S. Arrhenius calculated the proportion of active (i.e. leading to a reaction) collisions of reacting particles a, depending on temperature: - a = exp(-E/RT). and brought Arrhenius equation for the reaction rate constant:
k = koe-E/RT
where ko and E d depend on the nature of the reagents. E is the energy that must be given to molecules in order for them to interact, called activation energy.
Van't Hoff's rule- an empirical rule that allows, as a first approximation, to estimate the effect of temperature on the rate of a chemical reaction in a small temperature range (usually from 0 ° C to 100 ° C). J. H. van't Hoff, on the basis of many experiments, formulated the following rule:
Activation energy in chemistry and biology, the minimum amount of energy required to impart to a system (in chemistry expressed in joules per mole) for a reaction to occur. The term was introduced by Svante August Arrhenius in. Typical reaction energy notation Ea.
The activation entropy is considered as the difference between the entropy of the transition state and the ground state of the reactants. It is determined mainly by the loss of the translational and rotational degrees of freedom of the particles during the formation of the activated complex. Significant changes (vibrational degrees of freedom) can also occur if the activated complex is somewhat more densely packed than the reactants.
The activation entropy of such a transition is positive.
The entropy of activation depends on many factors. When, in a bimolecular reaction, two initial particles join together to form a transition state, the translational and rotational entropy of the two particles is reduced to values corresponding to a single particle; a slight increase in vibrational entropy is not enough to compensate for this effect.
Activation entropies, in fact, vary more with structure than enthalpies. The activation entropies are in good agreement in most cases with the Price and Hammett rule. This series also has the particular significance that the increase and entropy of the silap can probably be accurately calculated from the known absolute entropies of the corresponding hydrocarbons.
Ticket number 2
1) MAIN CLASSES OF INORGANIC COMPOUNDS: Bases, oxides, acids, salts.
2) Be - beryllium.
Chemical properties: beryllium is relatively unreactive at room temperature. In compact form, it does not react with water and water vapor even at red heat and is not oxidized by air up to 600 °C. When ignited, beryllium powder burns with a bright flame, producing oxide and nitride. Halogens react with beryllium at temperatures above 600 °C, while chalcogens require even higher temperatures.
Physical properties: Beryllium is a relatively hard, but brittle, silvery-white metal. It has a high modulus of elasticity - 300 GPa (for steels - 200-210 GPa). In air, it is actively covered with a resistant oxide film.
Magnesium (Mg). Physical properties: Magnesium is a silver-white metal with a hexagonal lattice, space group P 63 / mmc, lattice parameters a = 0.32029 nm, c = 0.52000 nm, Z = 2. Under normal conditions, the surface of magnesium is covered with a strong protective film of magnesium oxide MgO , which is destroyed when heated in air to about 600 ° C, after which the metal burns with a dazzling white flame to form magnesium oxide and nitride Mg3N2.
Chemical properties: Mixture of powdered magnesium with potassium permanganate KMnO4 - explosive
Hot magnesium reacts with water:
Mg (decay) + H2O = MgO + H2;
Alkalis do not act on magnesium, it dissolves easily in acids with the release of hydrogen:
Mg + 2HCl = MgCl2 + H2;
When heated in air, magnesium burns to form an oxide; a small amount of nitride can also form with nitrogen:
2Mg + O2 = 2MgO;
3Mg + N2 = Mg3N2
Ticket number 3. Solubility- the ability of a substance to form homogeneous systems with other substances - solutions in which the substance is in the form of individual atoms, ions, molecules or particles.
saturated solution- a solution in which the solute has reached its maximum concentration under given conditions and is no longer soluble. The precipitate of a given substance is in equilibrium with the substance in solution.
unsaturated solution- a solution in which the concentration of a solute is less than in a saturated solution, and in which, under given conditions, some more of it can be dissolved.
Supersaturated solutions- solutions characterized by the fact that the content of a dissolved substance in them is greater than its normal solubility under given conditions.
Henry's law- the law according to which, at a constant temperature, the solubility of a gas in a given liquid is directly proportional to the pressure of this gas over the solution. The law is suitable only for ideal solutions and low pressures.
Henry's law is usually written as follows:
Where p is the partial pressure of the gas above the solution,
c is the gas concentration in the solution in fractions of a mole,
k is the Henry coefficient.
Extraction(from late Latin extractio - extraction), extraction, the process of separating a mixture of liquid or solid substances using selective (selective) solvents (extractants).
Ticket number 4. 1)Mass fraction is the ratio of the mass of the solute to the total mass of the solution. For binary solution
ω(x) = m(x) / (m(x) + m(s)) = m(x) / m
where ω(x) - mass fraction of the dissolved substance X
m(x) - mass of dissolved substance X, g;
m(s) is the mass of the solvent S, g;
m \u003d m (x) + m (s) - mass of the solution, g.
2)Aluminum- an element of the main subgroup of the third group of the third period of the periodic system of chemical elements of D. I. Mendeleev, with atomic number 13.
Finding in nature:
Natural aluminum consists almost entirely of a single stable isotope, 27Al, with traces of 26Al, a radioactive isotope with a half-life of 720,000 years, formed in the atmosphere when argon nuclei are bombarded by cosmic ray protons.
Receipt:
It consists in the dissolution of aluminum oxide Al2O3 in a melt of Na3AlF6 cryolite, followed by electrolysis using consumable coke oven or graphite electrodes. This method of obtaining requires large amounts of electricity, and therefore was in demand only in the 20th century.
Aluminothermy- a method for obtaining metals, non-metals (as well as alloys) by reducing their oxides with metallic aluminum.
Ticket number 5. SOLUTIONS OF NON-ELECTROLYTES, binary or multicomponent pier. systems, the composition of which can change continuously (at least within certain limits). Unlike electrolyte solutions, there are no charged particles in any noticeable concentrations in non-electrolyte solutions (mol. solutions). solutions of non-electrolytes can be solid, liquid and gaseous.
Raoult's first law
Raoult's first law relates the saturation vapor pressure over a solution to its composition; it is formulated as follows:
The partial pressure of the saturated vapor of a solution component is directly proportional to its mole fraction in the solution, and the coefficient of proportionality is equal to the saturated vapor pressure over the pure component.
Raoult's second law
The fact that the vapor pressure of a solution differs from the vapor pressure of a pure solvent significantly affects the crystallization and boiling processes. From Raoult's first law, two consequences are derived regarding the decrease in the freezing point and the increase in the boiling point of solutions, which, in their combined form, are known as the second Raoult's law.
Cryoscopy(from the Greek kryos - cold and scopeo - look) - measurement of the decrease in the freezing point of a solution compared to a pure solvent.
Van't Hoff's rule - For every 10 degrees increase in temperature, the rate constant of a homogeneous elementary reaction increases two to four times
Hardness of water- a set of chemical and physical properties of water associated with the content of dissolved salts of alkaline earth metals in it, mainly calcium and magnesium.
Ticket number 6. SOLUTIONS OF ELECTROLYTES, contain appreciable concentrations of ions-cations and anions formed as a result of the electrolytic dissociation of the molecules of the dissolved matter.
Strong electrolytes- chemical compounds whose molecules in dilute solutions are almost completely dissociated into ions.
Weak electrolytes- chemical compounds, the molecules of which, even in highly dilute solutions, are not completely dissociated into ions, which are in dynamic equilibrium with undissociated molecules.
electrolytic dissociation- the process of decomposition of the electrolyte into ions when it is dissolved in a polar solvent or when melted.
Ostwald dilution law- ratio expressing the dependence of the equivalent electrical conductivity of a dilute solution of a binary weak electrolyte on the concentration of the solution:
P-elements 4 groups- carbon, silicon, germanium, tin and lead.
Ticket number 7. 1) Electrolytic dissociation- this is the disintegration of a substance into ions under the action of polar solvent molecules.
pH = -lg.
buffer solutions- These are solutions when acids or alkalis are added to which their pH changes slightly.
Carbonic acid forms:
1) medium salts (carbonates),
2) acidic (hydrocarbonates).
Carbonates and hydrocarbonates are thermally unstable:
CaCO3 \u003d CaO + CO2 ^,
Ca (HCO3) 2 \u003d CaCO3v + CO2 ^ + H2O.
Sodium carbonate (soda ash) is one of the main products of the chemical industry. In aqueous solution, it hydrolyzes according to the reaction
Na2CO3 > 2Na+ + CO3-2,
CO3-2 + H + -OH- - HCO3- + OH-.
Sodium bicarbonate (baking soda) is widely used in the food industry. Due to hydrolysis, the solution also has an alkaline environment.
NaHCO3 > Na+ + HCO3-, HCO3- + H-OH - H2CO3 + OH-.
Soda ash and drinking soda interact with acids
Na2CO3 + 2HCl - 2NaCl + CO2 ^ + H2O,
2Na+ + CO3-2 + 2H+ + 2Cl- - 2Na+ + 2Cl- + CO2^ + H2O,
CO3-2 + 2H+ - CO2^ + H2O;
NaHCO3 + CH3COOH - CH3COOHa + CO2^ + H2O,
Na+ + HCO3- + CH3COOH - CH3COO- + Na+ + CO2^ + H2O,
HCO3- + CH3COOH - CH3COO- + CO2^ + H2O.
Ticket number 8. 1)_ion-exchange in solutions:
Na2CO3 + H2SO4 → Na2SO4 + CO2 +H2O
2Na + CO3 + 2H + SO4 → 2Na + SO4 + CO2 + H2O
CO3 + 2H → CO2 + H2O
With gas evolution: Na2CO3 + 2HCl = CO2 + H2O + 2NaCl
2) Chemical properties of nitrogen. Only with such active metals as lithium, calcium, magnesium, nitrogen interacts when heated to relatively low temperatures. Nitrogen reacts with most other elements at high temperatures and in the presence of catalysts. Nitrogen compounds with oxygen N2O, NO, N2O3, NO2 and N2O5 are well studied.
Physical properties of nitrogen. Nitrogen is slightly lighter than air; density 1.2506 kg/m3 (at 0°С and 101325 n/m2 or 760 mm Hg), mp -209.86°С, tbp -195.8°С. Nitrogen liquefies with difficulty: its critical temperature is rather low (-147.1°C) and its critical pressure is high, 3.39 MN/m2 (34.6 kgf/cm2); the density of liquid nitrogen is 808 kg/m3. Nitrogen is less soluble in water than oxygen: at 0°C, 23.3 g of nitrogen dissolves in 1 m3 of H2O. Better than water, nitrogen is soluble in some hydrocarbons.
Ticket number 9. Hydrolysis (from Greek hydro - water, lysis - decomposition) means the decomposition of a substance by water. Salt hydrolysis is the reversible interaction of salt with water, leading to the formation of a weak electrolyte.
Water, although to a small extent, dissociates:
H 2 O H + + OH -.
Sodium chloride H2O H+ + OH–,
Na+ + Cl– + H2O Na+ + Cl– + H+ + OH–,
NaCl + H2O (no reaction) Neutral
Sodium carbonate + HOH + OH–,
2Na+ + + H2O + OH–,
Na2CO3 + H2O NaHCO3 + NaOH Alkaline
Aluminum chloride Al3+ + HOH AlOH2+ + H+,
Al3+ + 3Cl– + H2O AlОH2+ + 2Cl– + H+ + Cl–,
AlCl3 + H2O AlOHCl2 + HCl acidic
The dependence of the rate of a chemical reaction on temperature.
The rate of heterogeneous reactions.
In heterogeneous systems, reactions proceed at the interface. In this case, the concentration of the solid phase remains practically constant and does not affect the reaction rate. The rate of a heterogeneous reaction will depend only on the concentration of the substance in the liquid or gaseous phase. Therefore, in the kinetic equation, the concentrations of solids are not indicated, their values are included in the values of the constants. For example, for a heterogeneous reaction
the kinetic equation can be written
EXAMPLE 4. The kinetic order of the reaction of the interaction of chromium with aluminum is 1. Write the chemical and kinetic equations of the reaction.
The reaction of interaction of aluminum with chlorine is heterogeneous, the kinetic equation can be written
EXAMPLE 5 Kinetic Reaction Equation
has the form
Determine the dimension of the rate constant and calculate the rate of dissolution of silver at a partial pressure of oxygen Pa and a concentration of potassium cyanide of 0.055 mol/l.
The dimension of the constant is determined from the kinetic equation given in the condition of the problem:
Substituting these problems into the kinetic equation, we find the rate of silver dissolution:
EXAMPLE 6 Kinetic Reaction Equation
has the form
How will the reaction rate change if the concentration of mercury chloride (P) is halved, and the concentration of oxalate – ions to double?
After changing the concentration of the starting substances, the reaction rate is expressed by the kinetic equation
Comparing and, we find that the reaction rate increased in 2 times.
As the temperature rises, the rate of a chemical reaction increases markedly.
The quantitative dependence of the reaction rate on temperature is determined by the van't Hoff rule.
To characterize the dependence of the rate of a chemical reaction (rate constant) on temperature, the temperature coefficient of the rate, reaction (), also called the Van't Hoff coefficient, is used. The temperature coefficient of the reaction rate shows how many times the reaction rate will increase with an increase in the temperature of the reactants by 10 degrees.
Mathematically, the dependence of the reaction rate on temperature is expressed by the relation
Where – temperature coefficient of speed;
– T;
T;
–– reaction rate constant at temperature T+ 10;
–– reaction rate at temperature T+ 10.
For calculations, it is more convenient to use the equations
as well as the logarithmic forms of these equations
The increase in the reaction rate with increasing temperature explains activation theory. According to this theory, the particles of the reacting substances in the collision must overcome the repulsive forces, weaken or break old chemical bonds and form new ones. For this, they must spend a certain amount of energy, i.e. overcome some energy barrier. A particle with excess energy sufficient to overcome the energy barrier is called active particles.
Under normal conditions, there are few active particles in the system, and the reaction proceeds at a slower rate. But inactive particles can become active if you give them additional energy. One way to activate the particles is by increasing the temperature. As the temperature rises, the number of active particles in the system sharply increases and the reaction rate increases.