Temperature coefficient of reaction rate. The dependence of the reaction rate on temperature. Van't Hoff rule. The temperature coefficient of the reaction rate and its features for biochemical processes
The rate of chemical reactions increases with increasing temperature. The increase in the reaction rate with temperature can be estimated using the van't Hoff rule. According to the rule, an increase in temperature by 10 degrees increases the rate constant of the reaction by 2-4 times:
This rule is not followed when high temperatures when the rate constant almost does not change with temperature.
Van't Hoff's rule allows you to quickly determine the expiration date of a drug. An increase in temperature increases the rate of decomposition of the drug. This shortens the time to determine the expiration date of the drug.
The method lies in the fact that the drug is kept at elevated temperature T certain time tT, find the amount of decomposed drug m and recalculate to a standard storage temperature of 298K. Considering the process of decomposition of the drug as a first-order reaction, the rate is expressed at the selected temperature T and T = 298K:
Considering the mass of the decomposed drug to be the same for standard and real storage conditions, the decomposition rates can be expressed by the equations:
Assuming T=298+10n, where n = 1,2,3…,
Get the final expression for the shelf life of the drug under standard conditions 298K:
Theory of active collisions. Activation energy. Arrhenius equation. Relationship between reaction rate and activation energy.
The theory of active collisions was formulated by S. Arrhenius in 1889. This theory is based on the idea that for a chemical reaction to occur, a collision between the molecules of the initial substances is necessary, and the number of collisions is determined by the intensity of the thermal motion of the molecules, i.e. temperature dependent. But not every collision of molecules leads to a chemical transformation: only active collision leads to it.
Active collisions are collisions that occur, for example, between molecules A and B with large stock energy. The minimum amount of energy that the molecules of the initial substances must have in order for their collision to be active is called the energy barrier of the reaction.
Activation energy is the excess energy that can be communicated or transferred to one mole of a substance.
The activation energy significantly affects the value of the reaction rate constant and its dependence on temperature: the larger Ea, the lower the rate constant and the more significantly the change in temperature affects it.
The reaction rate constant is related to the activation energy by a complex relationship described by the Arrhenius equation:
k=Ae–Ea/RT, where A is the pre-exponential factor; Ea is the activation energy, R is the universal gas constant equal to 8.31 j/mol; T is the absolute temperature;
e is the base of natural logarithms.
However, the observed reaction rate constants are generally much smaller than those calculated using the Arrhenius equation. Therefore, the equation for the reaction rate constant is modified as follows:
(minus before whole fraction)
The multiplier causes the temperature dependence of the rate constant to differ from the Arrhenius equation. Since the Arrhenius activation energy is calculated as the slope of the logarithmic dependence of the reaction rate on the reciprocal temperature, then doing the same with the equation , we get:
Features of heterogeneous reactions. The rate of heterogeneous reactions and factors determining it. Kinetic and diffusion regions of heterogeneous processes. Examples of heterogeneous reactions of interest to pharmacy.
HETEROGENEOUS REACTIONS, chem. reactions involving substances in decomp. phases and constituting together a heterogeneous system. Typical heterogeneous reactions: thermal. decomposition of salts to form gaseous and solid products (e.g. CaCO3 -> CaO + CO2), reduction of metal oxides with hydrogen or carbon (e.g. PbO + C -> Pb + CO), dissolution of metals in acids (e.g. Zn + + H2SO4 -> ZnSO4 + H2), interaction. solid reagents (A12O3 + NiO -> NiAl2O4). In a special class, heterogeneous catalytic reactions occurring on the catalyst surface are distinguished; while the reagents and products may not be in different phases. Direction, in the reaction N2 + + 3H2 -> 2NH3 occurring on the surface of an iron catalyst, the reactants and the reaction product are in the gas phase and form a homogeneous system.
The features of heterogeneous reactions are due to the participation of condensed phases in them. This makes it difficult to mix and transport reactants and products; activation of reagent molecules on the interface is possible. The kinetics of any heterogeneous reaction is defined as the rate of the chemical itself. transformations and transfer processes (diffusion) necessary to replenish the consumption of reactants and remove reaction products from the reaction zone. In the absence of diffusion hindrances, the rate of a heterogeneous reaction is proportional to the size of the reaction zone; this is the name of the specific reaction rate calculated per unit surface (or volume) of the reaction. zones, does not change in time; for simple (single-step) reactions, it can be determined on the basis of the acting masses of the law. This law is not satisfied if the diffusion of substances proceeds more slowly than chemical. district; in this case, the observed rate of the heterogeneous reaction is described by the equations of diffusion kinetics.
The rate of a heterogeneous reaction is the amount of a substance that enters into a reaction or is formed during a reaction per unit time per unit area of the phase surface.
Factors affecting the rate of a chemical reaction:
The nature of the reactants
The concentration of reagents,
Temperature,
The presence of a catalyst.
Vheterog = Δp(S Δt), where Vheterog is the reaction rate in a heterogeneous system; n is the number of moles of any of the substances resulting from the reaction; V is the volume of the system; t - time; S is the surface area of the phase on which the reaction proceeds; Δ - increment sign (Δp = p2 - p1; Δt = t2 - t1).
As the temperature rises, the rate of a chemical process usually increases. In 1879, the Dutch scientist J. van't Hoff formulated an empirical rule: with an increase in temperature by 10 K, the rate of most chemical reactions increases by 2-4 times.
Mathematical notation of the rule I. van't Hoff:
γ 10 \u003d (k t + 10) / k t, where k t is the rate constant of the reaction at temperature T; k t+10 - reaction rate constant at temperature T+10; γ 10 - Van't Hoff temperature coefficient. Its value ranges from 2 to 4. For biochemical processesγ 10 varies from 7 to 10.
All biological processes proceed in a certain temperature range: 45-50°C. Optimum temperature is 36-40°C. In the body of warm-blooded animals, this temperature is maintained constant due to the thermoregulation of the corresponding biosystem. When studying biosystems, temperature coefficients γ 2 , γ 3 , γ 5 are used. For comparison, they are brought to γ 10 .
The dependence of the reaction rate on temperature, in accordance with the van't Hoff rule, can be represented by the equation:
V 2 /V 1 \u003d γ ((T 2 -T 1) / 10)
Activation energy. A significant increase in the reaction rate with increasing temperature cannot be explained only by an increase in the number of collisions between particles of reacting substances, since, in accordance with the kinetic theory of gases, the number of collisions increases slightly with increasing temperature. The increase in the reaction rate with increasing temperature is explained by the fact that a chemical reaction does not occur with any collision of particles of reacting substances, but only with a meeting of active particles that have the necessary excess energy at the moment of collision.
The energy required to turn inactive particles into active particles is called activation energy (Ea). Activation energy - excess, compared with the average value, the energy required for the entry of reactants into the reaction when they collide. The activation energy is measured in kilojoules per mole (kJ/mol). Usually E is from 40 to 200 kJ/mol.
Energy diagram of exothermic and endothermic reaction shown in fig. 2.3. For any chemical process, it is possible to distinguish the initial, intermediate and final states. At the top of the energy barrier, the reactants are in an intermediate state called the activated complex, or transition state. The difference between the energy of the activated complex and the initial energy of the reagents is Ea, and the difference between the energy of the reaction products and starting materials (reagents) is ΔH, thermal effect reactions. The activation energy, in contrast to ΔH, is always a positive value. For an exothermic reaction (Fig. 2.3, a), the products are located at a lower energy level than the reactants (Ea< ΔН).
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Ea is the main factor determining the reaction rate: if Ea > 120 kJ/mol (higher energy barrier, fewer active particles in the system), the reaction is slow; and vice versa, if Ea< 40 кДж/моль, реакция осуществляется с большой скоростью.
For reactions involving complex biomolecules, one should take into account the fact that in an activated complex formed during the collision of particles, the molecules must be oriented in space in a certain way, since only the reacting region of the molecule undergoes transformation, which is small in relation to its size.
If the rate constants k 1 and k 2 are known at temperatures T 1 and T 2 , the value of Ea can be calculated.
In biochemical processes, the activation energy is 2-3 times less than in inorganic ones. At the same time, the Ea of reactions involving foreign substances, xenobiotics, significantly exceeds the Ea of conventional biochemical processes. This fact is the natural bioprotection of the system from the influence of foreign substances, i.e. reactions natural for the body occur under favorable conditions with low Ea, and for foreign reactions, Ea is high. This is a gene barrier that characterizes one of the main features of the course of biochemical processes.
Factors affecting the course of the reaction
In the human body, thousands of enzymatic reactions take place in a living cell. However, in a multistage chain of processes, the difference between the rates of individual reactions is quite large. Thus, the synthesis of protein molecules in a cell is preceded by at least two more stages: the synthesis of transfer RNA and the synthesis of ribosomes. But the time during which the concentration of tRNA molecules doubles is 1.7 minutes, protein molecules - 17 minutes, and ribosomes - 170 minutes. The rate of the overall process of the slow (limiting) stage, in our example, the rate of ribosome synthesis. The presence of a limiting reaction provides high reliability and flexibility in controlling thousands of reactions occurring in the cell. It is enough to keep under observation and regulate only the slowest of them. This method of controlling the rate of multi-stage synthesis is called the minimum principle. It allows to significantly simplify and make more reliable the system of autoregulation in the cell.
Classifications of reactions used in kinetics: reactions, homogeneous, heterogeneous and microheterogeneous; simple and complex reactions (parallel, sequential, conjugated, chain). Molecularity of the elementary act of the reaction. Kinetic equations. Reaction order. Half life
Microheterogeneous reactions -
The molecularity of the reaction is determined by the number of molecules that enter into chemical interaction in the elementary act of the reaction. On this basis, the reactions are divided into monomolecular, bimolecular and trimolecular.
Then reactions of type A -> B will be monomolecular, for example:
a) C 16 H 34 (t ° C) -> C g H 18 + C 8 H 16 - hydrocarbon cracking reaction;
b) CaC0 3 (t ° C) -> CaO + C0 2 - thermal decomposition of calcium carbonate.
Reactions like A + B -> C or 2A -> C - are bimolecular, for example:
a) C + 0 2 -> C0 2; b) 2Н 2 0 2 -> 2Н 2 0 + 0 2 etc.
Trimolecular reactions are described general equations type:
a) A + B + C D; b) 2A + B D; c) 3A D.
For example: a) 2Н 2 + 0 2 2Н 2 0; b) 2NO + H 2 N 2 0 + H 2 0.
The reaction rate depending on the molecularity will be expressed by the equations: a) V = k C A - for a monomolecular reaction; b) V \u003d to C A C in or c) V \u003d to C 2 A - for a bimolecular reaction; d) V \u003d k C C in C e) V \u003d k C 2 A C in or e) V \u003d k C 3 A - for a trimolecular reaction.
Molecularity is the number of molecules that react in one elementary chemical act.
Often the molecularity of the reaction is difficult to establish, so more formal sign is the order of a chemical reaction.
Reaction order is equal to the sum indicators of degrees of concentration in the equation expressing the dependence of the reaction rate on the concentration of reactants (kinetic equation).
The order of the reaction most often does not coincide with the molecularity due to the fact that the reaction mechanism, i.e., the "elementary act" of the reaction (see the definition of the sign of molecularity), is difficult to establish.
Consider a number of examples illustrating this position.
1. The rate of dissolution of crystals is described by the equations of zero-order kinetics, despite the monomolecular nature of the reaction: AgCl (TB) -> Ag + + CI", V = k C (AgCl (TB p = k" C (AgCl (ra)) - p - density and is a constant value, i.e., the dissolution rate does not depend on the amount (concentration) of the dissolved substance.
2. The reaction of sucrose hydrolysis: CO + H 2 0 -> C 6 H 12 0 6 (glucose) + C 6 H 12 0 6 (fructose) is a bimolecular reaction, but its kinetics is described by a first-order kinetic equation: V \u003d k * C cax , since under experimental conditions, including in the body, the concentration of water is a constant value С(Н 2 0) - const.
3. 
The decomposition reaction of hydrogen peroxide, which proceeds with the participation of catalysts, both inorganic ions Fe 3+, Cu 2+ of metallic platinum, and biological enzymes, such as catalase, has general form:
2H 2 0 2 -\u003e 2H 2 0 + O e, i.e., is bimolecular.
Dependence of reaction rate on concentration. Kinetic equations of reactions of the first, second and zero orders. Experimental methods for determining the rate and rate constant of reactions.
The dependence of the reaction rate on temperature. Van't Hoff rule. Temperature coefficient reaction rate and its features for biochemical processes.
γ is the temperature coefficient of the reaction rate.
physical meaning The value of γ lies in the fact that it shows how many times the reaction rate changes with a change in temperature for every 10 degrees.
15. The concept of the theory of active collisions. Energy profile of the reaction; activation energy; Arrhenius equation. The role of the steric factor. The concept of the theory of the transition state.
The relationship of the rate constant, activation energy and temperature is described by the Arrhenius equation: k T \u003d k 0 *Ae ~ E / RT, where k t and k 0 are the rate constants at temperature T and T e e is the base of the natural logarithm, A is the steric factor.
The steric factor A determines the probability of a collision of two reacting particles in active center molecules. This factor is especially important for biochemical reactions with biopolymers. At acid-base reactions The H + ion must react with the terminal carboxyl group - COO. "However, not every collision of the H + ion with a protein molecule will lead to this reaction. Only those collisions that are directly carried out at certain points of the macromolecules, called active centers, will be effective.
It follows from the Arrhenius equation that the higher the rate constant, the lower the value of the activation energy E and the higher the temperature T of the process.
Temperature and reaction rate
At a fixed temperature, a reaction is possible if the interacting molecules have a certain amount of energy. Arrhenius called this excess energy activation energy , and the molecules themselves activated.
According to Arrhenius, the rate constant k and activation energy E a are related by a relation called the Arrhenius equation:
Here A is the pre-exponential factor, R is the universal gas constant, T is the absolute temperature.
Thus, at a constant temperature, the reaction rate determines E a. The more E a, the smaller the number of active molecules and the slower the reaction proceeds. When decreasing E a speed increases and E a= 0 the reaction proceeds instantaneously.
Value E a characterizes the nature of the reacting substances and is determined experimentally from the dependence k = f(T). Writing equation (5.3) in logarithmic form and solving it for constants at two temperatures, we find E a:
γ is the temperature coefficient of the chemical reaction rate. The van't Hoff rule has limited application, since the value of γ depends on temperature, and outside the region E a= 50–100 kJ ∙ mol–1 this rule is not fulfilled at all.
On fig. 5.4 it can be seen that the energy spent on the transfer of the initial products to the active state (A * - activated complex) is then fully or partially re-emitted during the transition to the final products. The difference between the energies of the initial and final products determines Δ H reaction that does not depend on the activation energy.
Thus, on the way from the initial state to the final state, the system must overcome the energy barrier. Only active molecules possessing at the moment of collision the necessary energy excess equal to E a, can overcome this barrier and enter into a chemical interaction. As the temperature rises, the proportion of active molecules in the reaction medium increases.
Preexponential multiplierA characterizes total number collisions. For reactions with simple molecules A close to theoretical collision magnitude Z, i.e. A = Z calculated from the kinetic theory of gases. For complex molecules A ≠ Z, so it is necessary to introduce the steric factor P:
Here Z is the number of all collisions, P is the fraction of collisions favorable in spatial relation(takes values from 0 to ), is the fraction of active, i.e., energetically favorable collisions.
The dimension of the rate constant is obtained from the relation
Analyzing expression (5.3), we come to the conclusion that there are two fundamental possibilities for accelerating the reaction:
a) an increase in temperature,
b) decrease in activation energy.
Tasks and tests on the topic "Chemical kinetics. Temperature and reaction rate"
- The rate of a chemical reaction. Catalysts - Classification of chemical reactions and patterns of their course Grade 8–9
Lessons: 5 Assignments: 8 Quizzes: 1