Dependence of speed on temperature. Dependence of the rate of a chemical reaction on temperature. van't Hoff's rule. Temperature coefficient of reaction rate. Activation energy, entropy of activation of a reaction. Arrhenius equation Effect of pressure changes on p
At As the temperature increases, 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 energy greater than Ea is equal to the shaded area.
From Figure 2 it can be seen that as the temperature increases, the energy distribution of particles changes so that the proportion of particles with higher energy increases. Thus, an important concept for a chemical reaction is 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 occurring 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-expopotential factor, depending on the nature of the reacting substances;
R- gas constant = 8.314 J/(mol? K);
Ea- activation energy.
From the Arrhenius equation it follows 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 reaction path. From the above figure it can be seen that for an exothermic reaction (which occurs with the release of heat), the loss of active molecules is replenished by the energy released during the reaction. In the case of an endothermic reaction, heat is required to maintain the required reaction rate.
| Exothermic reaction | Endothermic reaction |
Figure 10.3 Energy diagram of a chemical reaction
A - reactants, C - products.
2.4 Influence of foreign substances
Foreign substances, depending on the impact they have, can accelerate reactions - catalysts or slow them down - inhibitors.
Catalysts- these are substances that accelerate chemical reactions, but remain unchanged after the reaction.
Inhibitors - these are substances that slow down reactions. In practice, it is sometimes necessary to slow down reactions (corrosion of metals, etc.) this is achieved by introducing inhibitors into the reaction system. For example, sodium nitrite, potassium chromate and dichromate reduce the rate of corrosion of metals.
Promoters- substances that increase the activity of the catalyst. In this case, promoters themselves may not have catalytic properties.
Catalytic poisons- foreign impurities in the reaction mixture, leading to partial or complete loss of catalyst activity. Thus, traces of arsenic and phosphorus cause a rapid loss of activity by the V 2 O 5 catalyst during the contact method for producing H 2 SO 4.
3. Chemical equilibrium
In chemical reactions, the starting substances are not always completely converted into reaction products. This occurs because as reaction products accumulate, conditions may be created for the reverse reaction to occur. Most chemical reactions are reversible.
As an example, let us analyze the reversible reaction of ammonia synthesis from nitrogen and hydrogen, which is extremely important for industry:
direct reaction -2N 2 + 3H 2 →2NH 3 ,
reverse reaction - 2NH 3 →N 2 + 3H 2,
reversible reaction - 2N 2 + 3H 2« 2NH 3.
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 forward and reverse reactions are equal. Examples of application of the law of mass action (LMA).
Let us derive the equilibrium constant using the example of the ammonia synthesis reaction.
Kinetic equation of forward reaction
N 2 +3H 2 →2NH 3
has the form Vpr = Kpr 3.
Kinetic equation of the reverse reaction
2NH 3 →N 2 + 3H 2
has the form Vobr = Cobr 2.
In a state of chemical equilibrium, Vpr = Vbr.
Substituting the expressions for the rates of direct and reverse reactions into the condition of chemical equilibrium, we obtain the following equality Kpr 3 = Cobr 2.
After transformation we get
.
4. Le Chatelier's principle
If a system in a state of chemical equilibrium is subject to any external influence, then the equilibrium as a result of processes occurring in the system will shift in such a way that the effect will decrease.
4.1 Effect of changing concentrations on equilibrium
When the concentration of any of the substances participating in the reaction increases, 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
2N 2 + 3H 2« 2NH 3
add N 2 or H 2 , then, in accordance with Le Chatelier’s principle, to reduce the concentrations of these substances, the equilibrium should shift to the right, the yield 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 changes on equilibrium
The pressure in a closed reaction system is determined by 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 are involved, and their quantity in the forward and reverse reactions is different.
If the pressure is increased in a system that is 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, a reaction occurs preferentially, as a result of which the amount of gaseous products increases.
Example 1. Is it possible to increase the yield of products in a 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 in the reaction does not change: 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. For this reason, changes in pressure do not affect the equilibrium state.
Example 2. How will the equilibrium concentrations of reactants change with increasing pressure in the system? N 2 + 3H 2 « 2NH 3 ?
From the reaction equation it is clear that from 4 moles of the gas of the initial products, 2 moles of the gas of the reaction products are formed. Thus, with an increase in pressure, the equilibrium will shift from the forward reaction, since it leads to a decrease in pressure.
4.3 Effect of temperature changes on chemical equilibrium
Most chemical reactions occur with the release or absorption of heat. In the first case, the temperature of the mixture increases, in the second it decreases.
If a reaction mixture in a state of chemical equilibrium is heated, then, in accordance with Le Chatelier’s principle, a predominantly reaction should occur, as a result of which heat will be absorbed, i.e. endothermic reaction; When the mixture is cooled, a reaction should predominantly occur, as a result of which heat will be released, i.e. endothermic reaction.
If the temperature is increased in a system 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, they maintain a temperature of 500 0C, since at a lower temperature the rate of the direct reaction sharply decreases.
Chemical equilibrium is dynamic in nature: forward and reverse reactions do not stop at equilibrium.
The equilibrium constant depends on the temperature and the nature of the reactants. The greater the equilibrium constant, the more the equilibrium is shifted towards the formation of direct reaction products
Le Chatelier's principle is universal, since it is applicable not only to purely chemical processes, but also to physicochemical phenomena, such as crystallization, dissolution, boiling, and phase transformations in solids.
The dependence of the reaction rate on temperature is approximately determined by the empirical Van't Hoff rule: With every 10 degree change in temperature, the rate of most reactions changes by 2-4 times.
Mathematically, van't Hoff's rule is expressed as follows:
where v(T2) and v(T1) are reaction rates, respectively, at temperatures T2 and T1 (T2> T1);
γ-temperature coefficient of reaction rate.
The value of γ for an endothermic reaction is higher than for an exothermic one. For many reactions, γ lies in the range of 2-4.
The physical meaning of the value γ is that it shows how many times the reaction rate changes with a change in temperature for every 10 degrees.
Since the reaction rate and the rate constant of a chemical reaction are directly proportional, expression (3.6) is often written in the following form:
(3.7)
where k(T2), k(T1) are reaction rate constants, respectively
at temperatures T2 and T1;
γ is the temperature coefficient of the reaction rate.
Example 8. How many degrees must the temperature be increased to increase the reaction rate by 27 times? The reaction temperature coefficient is 3.
Solution. We use expression (3.6):
We get: 27 = , = 3, DT = 30.
Answer: 30 degrees.
The speed of the reaction and the time during which it occurs are inversely proportional: the larger v, the more
less than t. Mathematically this is expressed by the relation
Example 9. At a temperature of 293 K, the reaction proceeds in 2 minutes. How long will this reaction take to occur at a temperature of 273 K if γ = 2.
Solution. From equation (3.8) it follows:
.
We use equation (3.6), since 
We get:
min.
Answer: 8 min.
Van't Hoff's rule is applicable to a limited number of chemical reactions. The effect of temperature on the rate of processes is often determined using the Arrhenius equation.
Arrhenius equation . In 1889, the Swedish scientist S. Arre-1ius, based on experiments, derived an equation that is named after him
where k is the reaction rate constant;
k0 - pre-exponential factor;
e is the base of the natural logarithm;
Ea is a constant called activation energy, determined by the nature of the reagents:
R is the universal gas constant equal to 8.314 J/mol×K.
Ea values for chemical reactions range from 4 to 400 kJ/mol.
Many reactions are characterized by a certain energy barrier. To overcome it, activation energy is necessary - some excess energy (compared to the harmful energy of molecules at a given temperature), which molecules must have in order for their collision to be effective, that is, to lead to the formation of a new substance. As the temperature rises, the number of active molecules increases rapidly, which leads to a sharp increase in the reaction rate.
In general, if the reaction temperature changes from T1 to T2, equation (3.9) after logarithm takes the form:
. (3.10)
This equation allows you to calculate the activation energy of a reaction as the temperature changes from T1 to T2.
The rate of chemical reactions increases in the presence of a catalyst. The effect of a catalyst is that it forms unstable intermediate compounds (activated complexes) with reagents, the decomposition of which leads to the formation of reaction products. In this case, the activation energy decreases, and molecules whose energy was insufficient to carry out the reaction in the absence of a catalyst become active. As a result, the total number of active molecules increases and the reaction rate increases.
The change in reaction rate in the presence of a catalyst is expressed by the following equation:
, (3.11)
where vcat, and Ea(cat) are the speed and activation energy of a chemical reaction in the presence of a catalyst;
v and Ea are the speed and activation energy of a chemical reaction without a catalyst.
Example 10. The activation energy of a certain reaction in the absence of a catalyst is 75.24 kJ/mol, with a catalyst - 50.14 kJ/mol. How many times does the reaction rate increase in the presence of a catalyst if the reaction occurs at a temperature of 298 K? Solution. Let's use equation (3.11). Substituting data into the equation
Ticket#2
1) MAIN CLASSES OF INORGANIC COMPOUNDS: Bases, oxides, acids, salts.
2) Be – beryllium.
Chemical properties: beryllium is relatively little reactive at room temperature. In its compact form, it does not react with water and steam even at red heat temperatures 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, and 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 persistent 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 durable protective film of magnesium oxide MgO , which is destroyed when heated in air to approximately 600 °C, after which the metal burns with a blinding white flame to form magnesium oxide and nitride Mg3N2.
Chemical properties: A mixture of powdered magnesium with potassium permanganate KMnO4 - explosive
Hot magnesium reacts with water:
Mg (dec.) + H2O = MgO + H2;
Alkalis do not affect magnesium; it dissolves easily in acids, releasing 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 No. 3. Solubility- the ability of a substance to form homogeneous systems with other substances - solutions in which the substance is found in the form of individual atoms, ions, molecules or particles.
Saturated solution- a solution in which the solute under given conditions has reached its maximum concentration and no longer dissolves. The precipitate of this substance is in equilibrium with the substance in solution.
Unsaturated solution- a solution in which the concentration of the solute is less than in a saturated solution, and in which under the given conditions some more of it can be dissolved.
Supersaturated solutions- solutions characterized by the fact that the content of dissolved substance in them is greater than that corresponding to 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 above 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 gas above the solution,
c is the gas concentration in the solution in fractions of a mole,
k - Henry's coefficient.
Extraction(from Late Latin extractio - extraction), extraction, the process of separating a mixture of liquid or solid substances using selective solvents (extractants).
Ticket No. 4. 1)Mass fraction This is the ratio of the mass of the solute to the total mass of the solution. For a binary solution
ω(x) = m(x) / (m(x) + m(s)) = m(x) / m
where ω(x) is the mass fraction of dissolved substance X
m(x) - mass of dissolved substance X, g;
m(s) - mass of solvent S, g;
m = m(x) + m(s) - mass of 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 produced in the atmosphere when argon nuclei are bombarded by cosmic ray protons.
Receipt:
It consists of dissolving aluminum oxide Al2O3 in molten cryolite Na3AlF6, followed by electrolysis using consumable coke or graphite electrodes. This production method requires a lot of electricity, and therefore became popular only in the 20th century.
Aluminothermy- a method for producing metals, non-metals (as well as alloys) by reducing their oxides with metallic aluminum.
Ticket No. 5. SOLUTIONS OF NON-ELECTROLYTES, binary or multicomponent mol. systems, the composition of which can change continuously (at least within certain limits). Unlike solutions of electrolytes, solutions of nonelectrolytes (mol. solutions) contain no charged particles in any noticeable concentrations. solutions of non-electrolytes can be solid, liquid and gaseous.
Raoult's first law
Raoult's first law relates the saturated vapor pressure above 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, with the proportionality coefficient being equal to the saturated vapor pressure above the pure component.
Raoult's second law
The fact that the vapor pressure above the solution is different from the vapor pressure above the pure solvent significantly affects the processes of crystallization and boiling. 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 Raoult's second law.
Cryoscopy(from Greek kryos - cold and scopeo - I 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, mainly calcium and magnesium.
Ticket No. 6. ELECTROLYTE SOLUTIONS, contain noticeable concentrations of ions-cations and anions formed as a result of the electrolytic dissociation of molecules of the dissolved substance.
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 that are in dynamic equilibrium with undissociated molecules.
Electrolytic dissociation- the process of decomposition of an electrolyte into ions when it is dissolved in a polar solvent or during melting.
Ostwald's dilution law- a relationship expressing the dependence of the equivalent electrical conductivity of a dilute solution of a binary weak electrolyte on the concentration of the solution:
Group 4 P-elements– carbon, silicon, germanium, tin and lead.
Ticket No. 7. 1) Electrolytic dissociation- This is the decomposition of a substance into ions under the influence 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 = CaO + CO2^,
Ca(HCO3)2 = 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
Na2СО3 > 2Nа+ + СО3-2,
CO3-2 + H+-OH- - HCO3- + OH-.
Sodium bicarbonate (baking soda) – 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 baking soda interact with acids
Na2СО3 + 2НCl - 2NаСl + СО2^ + Н2О,
2Nа+ + СО3-2 + 2Н+ + 2Сl- - 2Nа+ + 2Сl- + СО2^ + Н2О,
CO3-2 + 2H+ - CO2^ + H2O;
NaHCO3 + CH3COON - CH3COONa + CO2^ + H2O,
Na+ + HCO3- + CH3COOH - CH3COO- + Na+ + CO2^ + H2O,
HCO3- + CH3COOH - CH3COO- + CO2^ + H2O.
Ticket No. 8. 1)_ion exchange in solutions:
Na2CO3 + H2SO4 → Na2SO4 + CO2 +H2O
2Na + CO3 + 2H + SO4 → 2Na + SO4 + CO2 + H2O
CO3 + 2H → CO2 + H2O
C gas evolution: Na2CO3 + 2HCl = CO2 + H2O + 2NaCl
2) Chemical properties of Nitrogen. Nitrogen interacts only with such active metals as lithium, calcium, magnesium 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 have been well studied.
Physical properties of Nitrogen. Nitrogen is slightly lighter than air; density 1.2506 kg/m3 (at 0°C and 101325 n/m2 or 760 mm Hg), melting point -209.86°C, boiling point -195.8°C. Nitrogen liquefies with difficulty: its critical temperature is quite 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. Nitrogen is soluble in some hydrocarbons better than in water.
Ticket No. 9. Hydrolysis (from the Greek hydro - water, lysis - decomposition) means the decomposition of a substance by water. Salt hydrolysis is the reversible reaction 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 NaHCО3 + NaOH Alkaline
Aluminum chloride Al3+ + HOH AlOH2+ + H+,
Al3+ + 3Cl– + H2O AlОH2+ + 2Cl– + H+ + Cl–,
AlCl3 + H2O AlOHCl2 + HCl Acid
Dependence of the rate of a chemical reaction on temperature.
Speed of heterogeneous reactions.
In heterogeneous systems, reactions occur at the interface. In this case, the concentration of the solid phase remains almost 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, the concentrations of solids are not indicated in the kinetic equation; 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 between chromium and aluminum is 1. Write the chemical and kinetic equations of the reaction.
The reaction between aluminum and chlorine is heterogeneous, the kinetic equation can be written
EXAMPLE 5. Kinetic equation of the reaction
looks like
Determine the dimension of the rate constant and calculate the rate of silver dissolution at a partial pressure of oxygen Pa and a potassium cyanide concentration of 0.055 mol/l.
The dimension of the constant is determined from the kinetic equation given in the problem statement:
Substituting the problem data into the kinetic equation, we find the rate of silver dissolution:
EXAMPLE 6. Kinetic equation of the reaction
looks like
How will the reaction rate change if the concentration of mercuric chloride (M) 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 by 2 times.
As the temperature increases, 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 reaction rate (), 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 logarithmic forms of these equations
The increase in reaction rate with increasing temperature explains activation theory. According to this theory, when particles of reacting substances collide, they must overcome repulsive forces, weaken or break old chemical bonds and form new ones. They must expend a certain energy for this, i.e. overcome some kind of energy barrier. A particle that has 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 particles is by increasing the temperature. As the temperature rises, the number of active particles in the system increases sharply and the reaction rate increases.
The law of mass action establishes the relationship between the masses of reacting substances in chemical reactions at equilibrium. The law of mass action was formulated in 1864-1867. K. Guldberg and P. Waage. According to this law, the rate at which substances react with each other depends on their concentration. The law of mass action is used in various calculations of chemical processes. It makes it possible to solve the question in which direction the spontaneous course of the reaction under consideration is possible at a given ratio of the concentrations of the reacting substances, what yield of the desired product can be obtained.
Question 18. Van't Hoff's rule.
Van't Hoff's rule is 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). Van't Hoff, based on many experiments, formulated the following rule: With every 10 degree increase in temperature, the rate constant of a homogeneous elementary reaction increases two to four times. The equation that describes this rule is:
V = V0 * Y(T2 − T1) / 10
where V is the reaction rate at a given temperature (T2), V0 is the reaction rate at temperature T1, Y is the temperature coefficient of the reaction (if it is equal to 2, for example, then the reaction rate will increase 2 times when the temperature increases by 10 degrees).
It should be remembered that Van't Hoff's rule has a limited scope of applicability. Many reactions do not obey it, for example, reactions occurring at high temperatures, very fast and very slow reactions. Van't Hoff's rule also does not apply to reactions involving bulky molecules, such as proteins in biological systems. The temperature dependence of the reaction rate is more correctly described by the Arrhenius equation.
V = V0 * Y(T2 − T1) / 10
Question 19. Activation energy.
Activation energy in chemistry and biology, the minimum amount of energy that must be supplied to the system (in chemistry expressed in joules per mole) for a reaction to occur. The term was introduced by Svante August Arrhenius in 1889. A typical designation for the reaction energy is Ea.
Activation energy in physics is the minimum amount of energy that electrons of a donor impurity must receive in order to enter the conduction band.
In the chemical model known as the Theory of Active Collisions (TAC), there are three conditions necessary for a reaction to occur:
Molecules must collide. This is an important condition, but it is not sufficient, since a collision does not necessarily cause a reaction.
Molecules must have the necessary energy (activation energy). During a chemical reaction, the interacting molecules must pass through an intermediate state, which may have more energy. That is, the molecules must overcome an energy barrier; if this does not happen, the reaction will not begin.
The molecules must be correctly oriented relative to each other.
At low (for a certain reaction) temperature, most molecules have energy less than the activation energy and are unable to overcome the energy barrier. However, in a substance there will always be individual molecules whose energy is significantly higher than the average. Even at low temperatures, most reactions continue to occur. Increasing the temperature allows you to increase the proportion of molecules with sufficient energy to overcome the energy barrier. This increases the reaction speed.
Mathematical description
The Arrhenius equation establishes the relationship between activation energy and reaction rate:
k is the reaction rate constant, A is the frequency factor for the reaction, R is the universal gas constant, T is the temperature in kelvins.
As the temperature rises, the probability of overcoming the energy barrier increases. General rule of thumb: a 10K increase in temperature doubles the reaction rate
Transition state
The relationship between the activation energy (Ea) and the enthalpy (entropy) of the reaction (ΔH) in the presence and absence of a catalyst. The highest point of energy represents an energy barrier. In the presence of a catalyst, less energy is required to start a reaction.
A transition state is a state of a system in which the destruction and creation of a connection are balanced. The system is in a transition state for a short time (10-15 s). The energy that must be expended to bring the system into a transition state is called activation energy. In multistep reactions that include several transition states, the activation energy corresponds to the highest energy value. After overcoming the transition state, the molecules scatter again with the destruction of old bonds and the formation of new ones or with the transformation of the original bonds. Both options are possible, since they occur with the release of energy (this is clearly visible in the figure, since both positions are energetically lower than the activation energy). There are substances that can reduce the activation energy for a given reaction. Such substances are called catalysts. Biologists call such substances enzymes. Interestingly, catalysts thus speed up the reaction without participating in it themselves.
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