La ley de velocidad o ecuación de velocidad para una reacción química es una ecuación que vincula la velocidad de reacción inicial o directa con las concentraciones o presiones de los reactivos y parámetros constantes (normalmente coeficientes de velocidad y órdenes de reacción parciales). [1] Para muchas reacciones, la velocidad inicial viene dada por una ley de potencia como
donde [A] y [B] expresan la concentración de las especies A y B, generalmente en moles por litro ( molaridad , M). Los exponentes x y y son las parciales órdenes de reacción para A y B y el general orden de reacción es la suma de los exponentes. Suelen ser números enteros positivos, pero también pueden ser cero, fraccionarios o negativos. La constante k es la constante de velocidad de reacción o el coeficiente de velocidad de la reacción. Su valor puede depender de condiciones tales como temperatura, fuerza iónica, área de superficie de un adsorbente o irradiación de luz . Si la reacción se completa, la ecuación de velocidad para la velocidad de reacción se aplica durante el transcurso de la reacción.
Las reacciones elementales (de un solo paso) y los pasos de reacción tienen órdenes de reacción iguales a los coeficientes estequiométricos para cada reactivo. El orden general de la reacción, es decir, la suma de los coeficientes estequiométricos de los reactivos, es siempre igual a la molecularidad de la reacción elemental. Sin embargo, las reacciones complejas (de varios pasos) pueden tener o no órdenes de reacción iguales a sus coeficientes estequiométricos. Esto implica que el orden y la ecuación de velocidad de una reacción dada no pueden deducirse de manera confiable de la estequiometría y deben determinarse experimentalmente, ya que un mecanismo de reacción desconocido podría ser elemental o complejo. Cuando se ha determinado la ecuación de velocidad experimental, a menudo se utiliza para la deducción del mecanismo de reacción .
La ecuación de velocidad de una reacción con un supuesto mecanismo de varios pasos a menudo se puede derivar teóricamente utilizando supuestos de estado cuasi-estacionario de las reacciones elementales subyacentes, y compararse con la ecuación de velocidad experimental como una prueba del mecanismo supuesto. La ecuación puede implicar un orden fraccionario y puede depender de la concentración de una especie intermedia .
Una reacción también puede tener un orden de reacción indefinido con respecto a un reactivo si la velocidad no es simplemente proporcional a alguna potencia de la concentración de ese reactivo; por ejemplo, no se puede hablar del orden de reacción en la ecuación de velocidad para una reacción bimolecular entre moléculas adsorbidas :
Definición
Considere una reacción química típica en la que dos reactivos A y B se combinan para formar un producto C:
Esto también se puede escribir
Los prefactores -1, -2 y 3 (con signos negativos para los reactivos porque se consumen) se conocen como coeficientes estequiométricos . Una molécula de A se combina con dos de B para formar 3 de C, así que si usamos el símbolo [X] para el número de moles de la sustancia química X, [2]
Si la reacción tiene lugar en un sistema cerrado a temperatura y volumen constantes, sin acumulación de intermedios de reacción , la velocidad de reacción Se define como
donde ν i es el coeficiente estequiométrico de la sustancia química X i , con signo negativo para un reactivo. [3]
La velocidad de reacción inicial 0 =( t = 0) tiene cierta dependencia funcional de las concentraciones de los reactivos,
y esta dependencia se conoce como ecuación de tasa o ley de tasa . [4] Esta ley generalmente no se puede deducir de la ecuación química y debe determinarse mediante un experimento. [5]
Leyes de poder
Una forma común de la ecuación de tasa es una ley de potencia: [5]
La constante k se llama constante de velocidad . Los exponentes, que pueden ser fraccionarios, [5] se denominan órdenes parciales de reacción y su suma es el orden general de reacción. [6]
En una solución diluida, se encuentra empíricamente que una reacción elemental (una que tiene un solo paso con un solo estado de transición) obedece a la ley de acción de masas . Esto predice que la tasa depende solo de las concentraciones de los reactivos, elevadas a las potencias de sus coeficientes estequiométricos. [7]
Determinación del orden de reacción
Método de tasas iniciales
El logaritmo natural de la ecuación de tasa de ley de potencia es
This can be used to estimate the order of reaction of each reactant. For example, the initial rate can be measured in a series of experiments at different initial concentrations of reactant A with all other concentrations [B], [C], ... kept constant, so that
The slope of a graph of as a function of then corresponds to the order x with respect to reactant A.[8][9]
However, this method is not always reliable because
- measurement of the initial rate requires accurate determination of small changes in concentration in short times (compared to the reaction half-life) and is sensitive to errors, and
- the rate equation will not be completely determined if the rate also depends on substances not present at the beginning of the reaction, such as intermediates or products.
Integral method
The tentative rate equation determined by the method of initial rates is therefore normally verified by comparing the concentrations measured over a longer time (several half-lives) with the integrated form of the rate equation; this assumes that the reaction goes to completion.
For example, the integrated rate law for a first-order reaction is
- ,
where [A] is the concentration at time t and [A]0 is the initial concentration at zero time. The first-order rate law is confirmed if is in fact a linear function of time. In this case the rate constant is equal to the slope with sign reversed.[10][11]
Method of flooding
The partial order with respect to a given reactant can be evaluated by the method of flooding (or of isolation) of Ostwald. In this method, the concentration of one reactant is measured with all other reactants in large excess so that their concentration remains essentially constant. For a reaction a·A + b·B → c·C with rate law: , the partial order x with respect to A is determined using a large excess of B. In this case
with ,
and x may be determined by the integral method. The order y with respect to B under the same conditions (with B in excess) is determined by a series of similar experiments with a range of initial concentration [B]0 so that the variation of k' can be measured.[12]
Zero order
For zero-order reactions, the reaction rate is independent of the concentration of a reactant, so that changing its concentration has no effect on the speed of the reaction. Thus, the concentration changes linearly with time. This may occur when there is a bottleneck which limits the number of reactant molecules that can react at the same time, for example if the reaction requires contact with an enzyme or a catalytic surface.[13]
Many enzyme-catalyzed reactions are zero order, provided that the reactant concentration is much greater than the enzyme concentration which controls the rate, so that the enzyme is saturated. For example, the biological oxidation of ethanol to acetaldehyde by the enzyme liver alcohol dehydrogenase (LADH) is zero order in ethanol.[14]
Similarly reactions with heterogeneous catalysis can be zero order if the catalytic surface is saturated. For example, the decomposition of phosphine (PH3) on a hot tungsten surface at high pressure is zero order in phosphine which decomposes at a constant rate.[13]
In homogeneous catalysis zero order behavior can come about from reversible inhibition. For example, ring-opening metathesis polymerization using third-generation Grubbs catalyst exhibits zero order behavior in catalyst due to the reversible inhibition that is occur between the pyridine and the ruthenium center.[15]
First order
A first order reaction depends on the concentration of only one reactant (a unimolecular reaction). Other reactants can be present, but each will be zero order. The rate law for such a reaction is
The half-life is independent of the starting concentration and is given by .
Examples of such reactions are:
In organic chemistry, the class of SN1 (nucleophilic substitution unimolecular) reactions consists of first-order reactions. For example, in the reaction of aryldiazonium ions with nucleophiles in aqueous solution ArN2+ + X− → ArX + N2, the rate equation is v = k[ArN2+], where Ar indicates an aryl group.[16]
Second order
A reaction is said to be second order when the overall order is two. The rate of a second-order reaction may be proportional to one concentration squared , or (more commonly) to the product of two concentrations . As an example of the first type, the reaction NO2 + CO → NO + CO2 is second-order in the reactant NO2 and zero order in the reactant CO. The observed rate is given by , and is independent of the concentration of CO.[17]
For the rate proportional to a single concentration squared, the time dependence of the concentration is given by
The time dependence for a rate proportional to two unequal concentrations is
- ;
if the concentrations are equal, they satisfy the previous equation.
The second type includes nucleophillic addition-elimination reactions, such as the alkaline hydrolysis of ethyl acetate:[16]
- CH 3COOC 2H 5 + OH − → CH 3COO − + C 2H 5OH
This reaction is first-order in each reactant and second-order overall: 0 = k [CH3COOC2H5][OH−]
If the same hydrolysis reaction is catalyzed by imidazole, the rate equation becomes v = k[imidazole][CH3COOC2H5].[16] The rate is first-order in one reactant (ethyl acetate), and also first-order in imidazole which as a catalyst does not appear in the overall chemical equation.
Another well-known class of second-order reactions are the SN2 (bimolecular nucleophilic substitution) reactions, such as the reaction of n-butyl bromide with sodium iodide in acetone:
- CH 3CH 2CH 2CH 2Br + NaI → CH 3CH 2CH 2CH 2I + NaBr↓
This same compound can be made to undergo a bimolecular (E2) elimination reaction, another common type of second-order reaction, if the sodium iodide and acetone are replaced with sodium tert-butoxide as the salt and tert-butanol as the solvent:
- CH 3CH 2CH 2CH 2Br + NaO t-Bu → CH 3CH 2CH=CH 2 + NaBr + HO t-Bu
Pseudo-first order
If the concentration of a reactant remains constant (because it is a catalyst, or because it is in great excess with respect to the other reactants), its concentration can be included in the rate constant, obtaining a pseudo–first-order (or occasionally pseudo–second-order) rate equation. For a typical second-order reaction with rate equation v = k[A][B], if the concentration of reactant B is constant then 0 = k [A][B] = k'[A], where the pseudo–first-order rate constant k' = k[B]. The second-order rate equation has been reduced to a pseudo–first-order rate equation, which makes the treatment to obtain an integrated rate equation much easier.
One way to obtain a pseudo-first order reaction is to use a large excess of one reactant (say, [B]≫[A]) so that, as the reaction progresses, only a small fraction of the reactant in excess (B) is consumed, and its concentration can be considered to stay constant. For example, the hydrolysis of esters by dilute mineral acids follows pseudo-first order kinetics where the concentration of water is present in large excess:
- CH 3COOCH 3 + H 2O → CH 3COOH + CH 3OH
The hydrolysis of sucrose (C12H22O11) in acid solution is often cited as a first-order reaction with rate r = k[C12H22O11]. The true rate equation is third-order, r = k[C12H22O11][H+][H2O]; however, the concentrations of both the catalyst H+ and the solvent H2O are normally constant, so that the reaction is pseudo–first-order.[18]
Summary for reaction orders 0, 1, 2, and n
Elementary reaction steps with order 3 (called ternary reactions) are rare and unlikely to occur. However, overall reactions composed of several elementary steps can, of course, be of any (including non-integer) order.
Zero order | First order | Second order | nth order (g = 1-n) | |
---|---|---|---|---|
Rate Law | [19] | |||
Integrated Rate Law | [19] | [Except first order] | ||
Units of Rate Constant (k) | ||||
Linear Plot to determine k | [A] vs. t | vs. t | vs. t | vs. t [Except first order] |
Half-life | [19] | [Limit is necessary for first order] |
Where M stands for concentration in molarity (mol · L−1), t for time, and k for the reaction rate constant. The half-life of a first order reaction is often expressed as t1/2 = 0.693/k (as ln(2)≈0.693).
Fractional order
In fractional order reactions, the order is a non-integer, which often indicates a chemical chain reaction or other complex reaction mechanism. For example, the pyrolysis of acetaldehyde (CH3CHO) into methane and carbon monoxide proceeds with an order of 1.5 with respect to acetaldehyde: r = k[CH3CHO]3/2.[20] The decomposition of phosgene (COCl2) to carbon monoxide and chlorine has order 1 with respect to phosgene itself and order 0.5 with respect to chlorine: v = k[COCl2] [Cl2]1/2.[21]
The order of a chain reaction can be rationalized using the steady state approximation for the concentration of reactive intermediates such as free radicals. For the pyrolysis of acetaldehyde, the Rice-Herzfeld mechanism is
- Initiation
- CH 3CHO → •CH 3 + •CHO
- Propagation
- •CH 3 + CH 3CHO → CH 3CO• + CH 4
- CH 3CO• → •CH 3 + CO
- Termination
- 2 •CH 3 → C 2H 6
where • denotes a free radical.[20][22] To simplify the theory, the reactions of the •CHO to form a second •CH3 are ignored.
In the steady state, the rates of formation and destruction of methyl radicals are equal, so that
- ,
so that the concentration of methyl radical satisfies
- .
The reaction rate equals the rate of the propagation steps which form the main reaction products CH4 and CO:
in agreement with the experimental order of 3/2.[20][22]
Leyes complejas
Mixed order
More complex rate laws have been described as being mixed order if they approximate to the laws for more than one order at different concentrations of the chemical species involved. For example, a rate law of the form represents concurrent first order and second order reactions (or more often concurrent pseudo-first order and second order) reactions, and can be described as mixed first and second order.[23] For sufficiently large values of [A] such a reaction will approximate second order kinetics, but for smaller [A] the kinetics will approximate first order (or pseudo-first order). As the reaction progresses, the reaction can change from second order to first order as reactant is consumed.
Another type of mixed-order rate law has a denominator of two or more terms, often because the identity of the rate-determining step depends on the values of the concentrations. An example is the oxidation of an alcohol to a ketone by hexacyanoferrate (III) ion [Fe(CN)63−] with ruthenate (VI) ion (RuO42−) as catalyst.[24] For this reaction, the rate of disappearance of hexacyanoferrate (III) is
This is zero-order with respect to hexacyanoferrate (III) at the onset of the reaction (when its concentration is high and the ruthenium catalyst is quickly regenerated), but changes to first-order when its concentration decreases and the regeneration of catalyst becomes rate-determining.
Notable mechanisms with mixed-order rate laws with two-term denominators include:
- Michaelis-Menten kinetics for enzyme-catalysis: first-order in substrate (second-order overall) at low substrate concentrations, zero order in substrate (first-order overall) at higher substrate concentrations; and
- the Lindemann mechanism for unimolecular reactions: second-order at low pressures, first-order at high pressures.
Negative order
A reaction rate can have a negative partial order with respect to a substance. For example, the conversion of ozone (O3) to oxygen follows the rate equation in an excess of oxygen. This corresponds to second order in ozone and order (-1) with respect to oxygen.[25]
When a partial order is negative, the overall order is usually considered as undefined. In the above example for instance, the reaction is not described as first order even though the sum of the partial orders is , because the rate equation is more complex than that of a simple first-order reaction.
Reacciones opuestas
A pair of forward and reverse reactions may occur simultaneously with comparable speeds. For example, A and B react into products P and Q and vice versa (a, b, p, and q are the stoichiometric coefficients):
The reaction rate expression for the above reactions (assuming each one is elementary) can be written as:
where: k1 is the rate coefficient for the reaction that consumes A and B; k−1 is the rate coefficient for the backwards reaction, which consumes P and Q and produces A and B.
The constants k1 and k−1 are related to the equilibrium coefficient for the reaction (K) by the following relationship (set v=0 in balance):
Simple example
In a simple equilibrium between two species:
where the reaction starts with an initial concentration of reactant A, , and an initial concentration of 0 for product P at time t=0.
Then the constant K at equilibrium is expressed as:
Where and are the concentrations of A and P at equilibrium, respectively.
The concentration of A at time t, , is related to the concentration of P at time t, , by the equilibrium reaction equation:
The term is not present because, in this simple example, the initial concentration of P is 0.
This applies even when time t is at infinity; i.e., equilibrium has been reached:
then it follows, by the definition of K, that
and, therefore,
These equations allow us to uncouple the system of differential equations, and allow us to solve for the concentration of A alone.
The reaction equation was given previously as:
For this is simply
The derivative is negative because this is the rate of the reaction going from A to P, and therefore the concentration of A is decreasing. To simplify notation, let x be , the concentration of A at time t. Let be the concentration of A at equilibrium. Then:
Since:
The reaction rate becomes:
which results in:
A plot of the negative natural logarithm of the concentration of A in time minus the concentration at equilibrium versus time t gives a straight line with slope k1 + k−1. By measurement of [A]e and [P]e the values of K and the two reaction rate constants will be known.[26]
Generalization of simple example
If the concentration at the time t = 0 is different from above, the simplifications above are invalid, and a system of differential equations must be solved. However, this system can also be solved exactly to yield the following generalized expressions:
When the equilibrium constant is close to unity and the reaction rates very fast for instance in conformational analysis of molecules, other methods are required for the determination of rate constants for instance by complete lineshape analysis in NMR spectroscopy.
Reacciones consecutivas
If the rate constants for the following reaction are and ; , then the rate equation is:
- For reactant A:
- For reactant B:
- For product C:
With the individual concentrations scaled by the total population of reactants to become probabilities, linear systems of differential equations such as these can be formulated as a master equation. The differential equations can be solved analytically and the integrated rate equations are
The steady state approximation leads to very similar results in an easier way.
Reacciones paralelas o competitivas
When a substance reacts simultaneously to give two different products, a parallel or competitive reaction is said to take place.
Two first order reactions
and , with constants and and rate equations ; and
The integrated rate equations are then ; and .
One important relationship in this case is
One first order and one second order reaction
This can be the case when studying a bimolecular reaction and a simultaneous hydrolysis (which can be treated as pseudo order one) takes place: the hydrolysis complicates the study of the reaction kinetics, because some reactant is being "spent" in a parallel reaction. For example, A reacts with R to give our product C, but meanwhile the hydrolysis reaction takes away an amount of A to give B, a byproduct: and . The rate equations are: and . Where is the pseudo first order constant.[27]
The integrated rate equation for the main product [C] is , which is equivalent to . Concentration of B is related to that of C through
The integrated equations were analytically obtained but during the process it was assumed that therefeore, previous equation for [C] can only be used for low concentrations of [C] compared to [A]0
Redes de reacción estequiométricas
The most general description of a chemical reaction network considers a number of distinct chemical species reacting via reactions.[28][29] The chemical equation of the -th reaction can then be written in the generic form
which is often written in the equivalent form
Here
- is the reaction index running from 1 to ,
- denotes the -th chemical species,
- is the rate constant of the -th reaction and
- and are the stoichiometric coefficients of reactants and products, respectively.
The rate of such reaction can be inferred by the law of mass action
which denotes the flux of molecules per unit time and unit volume. Here is the vector of concentrations. This definition includes the elementary reactions:
- zero order reactions
for which for all ,
- first order reactions
for which for a single ,
- second order reactions
for which for exactly two , i.e., a bimolecular reaction, or for a single , i.e., a dimerization reaction.
Each of which are discussed in detail below. One can define the stoichiometric matrix
denoting the net extent of molecules of in reaction . The reaction rate equations can then be written in the general form
This is the product of the stoichiometric matrix and the vector of reaction rate functions. Particular simple solutions exist in equilibrium, , for systems composed of merely reversible reactions. In this case the rate of the forward and backward reactions are equal, a principle called detailed balance. Detailed balance is a property of the stoichiometric matrix alone and does not depend on the particular form of the rate functions . All other cases where detailed balance is violated are commonly studied by flux balance analysis which has been developed to understand metabolic pathways.[30][31]
Dinámica general de conversión unimolecular
For a general unimolecular reaction involving interconversion of different species, whose concentrations at time are denoted by through , an analytic form for the time-evolution of the species can be found. Let the rate constant of conversion from species to species be denoted as , and construct a rate-constant matrix whose entries are the .
Also, let be the vector of concentrations as a function of time.
Let be the vector of ones.
Let be the × identity matrix.
Let be the function that takes a vector and constructs a diagonal matrix whose on-diagonal entries are those of the vector.
Let be the inverse Laplace transform from to .
Then the time-evolved state is given by
- ,
thus providing the relation between the initial conditions of the system and its state at time .
Ver también
- Michaelis–Menten kinetics
- Molecularity
- Petersen matrix
- Reaction-diffusion equation
- Reactions on surfaces: rate equations for reactions where at least one of the reactants adsorbs onto a surface
- Reaction progress kinetic analysis
- Reaction rate
- Reaction rate constant
- Steady state approximation
- Gillespie algorithm
- Balance equation
- Belousov-Zhabotinsky reaction
- Lotka–Volterra equations
Referencias
- ^ IUPAC Gold Book definition of rate law. See also: According to IUPAC Compendium of Chemical Terminology.
- ^ Atkins & de Paula 2006, p. 794
- ^ IUPAC, Compendium of Chemical Terminology, 2nd ed. (the "Gold Book") (1997). Online corrected version: (2006–) "Rate of reaction". ‹See Tfd›doi:10.1351/goldbook.R05156‹See Tfd›
- ^ Atkins & de Paula 2006, p. 795
- ^ a b c Atkins & de Paula 2006, p. 796
- ^ Connors 1990, p. 13
- ^ Connors 1990, p. 12
- ^ Atkins & de Paula 2006, pp. 797–8
- ^ Espenson 1987, pp. 5–8
- ^ Atkins & de Paula 2006, pp. 798–800
- ^ Espenson 1987, pp. 15–18
- ^ Espenson 1987, pp. 30–31
- ^ a b Atkins & de Paula 2006, p. 796
- ^ Tinoco Jr. & Wang 1995, p. 331
- ^ Walsh, Dylan J.; Lau, Sii Hong; Hyatt, Michael G.; Guironnet, Damien (2017-09-25). "Kinetic Study of Living Ring-Opening Metathesis Polymerization with Third-Generation Grubbs Catalysts". Journal of the American Chemical Society. 139 (39): 13644–13647. doi:10.1021/jacs.7b08010. ISSN 0002-7863. PMID 28944665.
- ^ a b c Connors 1990
- ^ Whitten K. W., Galley K. D. and Davis R. E. General Chemistry (4th edition, Saunders 1992), pp. 638–9 ISBN 0-03-072373-6
- ^ Tinoco Jr. & Wang 1995, pp. 328–9
- ^ a b c NDRL Radiation Chemistry Data Center. See also: Capellos, Christos; Bielski, Benon H. (1972). Kinetic systems: mathematical description of chemical kinetics in solution. New York: Wiley-Interscience. ISBN 978-0471134503. OCLC 247275.
- ^ a b c Atkins & de Paula 2006, p. 830
- ^ Laidler 1987, p. 301
- ^ a b Laidler 1987, pp. 310–311
- ^ Espenson 1987, pp. 34,60
- ^ Mucientes, Antonio E.; de la Peña, María A. (November 2006). "Ruthenium(VI)-Catalyzed Oxidation of Alcohols by Hexacyanoferrate(III): An Example of Mixed Order". Journal of Chemical Education. 83 (11): 1643. doi:10.1021/ed083p1643. ISSN 0021-9584.
- ^ Laidler 1987, p. 305
- ^ Rushton, Gregory T.; Burns, William G.; Lavin, Judi M.; Chong, Yong S.; Pellechia, Perry; Shimizu, Ken D. (September 2007). "Determination of the Rotational Barrier for Kinetically Stable Conformational Isomers via NMR and 2D TLC". Journal of Chemical Education. 84 (9): 1499. doi:10.1021/ed084p1499. ISSN 0021-9584.
- ^ Manso, José A.; Pérez-Prior, M. Teresa; García-Santos, M. del Pilar; Calle, Emilio; Casado, Julio (2005). "A Kinetic Approach to the Alkylating Potential of Carcinogenic Lactones". Chemical Research in Toxicology. 18 (7): 1161–1166. CiteSeerX 10.1.1.632.3473. doi:10.1021/tx050031d. PMID 16022509.
- ^ Heinrich, Reinhart; Schuster, Stefan (2012). The Regulation of Cellular Systems. Springer Science & Business Media. ISBN 9781461311614.
- ^ Chen, Luonan; Wang, Ruiqi; Li, Chunguang; Aihara, Kazuyuki (2010). Modeling Biomolecular Networks in Cells. doi:10.1007/978-1-84996-214-8. ISBN 978-1-84996-213-1.
- ^ Szallasi, Z. and Stelling, J. and Periwal, V. (2006) System modeling in cell biology: from concepts to nuts and bolts. MIT Press Cambridge.
- ^ Iglesias, Pablo A.; Ingalls, Brian P. (2010). Control theory and systems biology. MIT Press. ISBN 9780262013345.
Books cited
- Atkins, Peter; de Paula, Julio (2006). "The rates of chemical reactions". Atkins' Physical chemistry (8th ed.). W.H. Freeman. pp. 791–823. ISBN 0-7167-8759-8.
- Connors, Kenneth Antonio (1990). Chemical kinetics : the study of reaction rates in solution. John Wiley & Sons. ISBN 9781560810063.
- Espenson, James H. (1987). Chemical kinetics and reaction mechanisms (2nd ed.). McGraw Hill. ISBN 9780071139496.
- Laidler, Keith James (1987). Chemical kinetics (3rd ed.). Harper & Row. ISBN 9780060438623.
- Tinoco Jr., Ignacio; Wang, James C. (1995). Physical chemistry : principles and applications in biological sciences (3rd ed.). Prentice Hall. ISBN 9780131865457.
enlaces externos
- Chemical kinetics, reaction rate, and order (needs flash player)
- Reaction kinetics, examples of important rate laws (lecture with audio).
- Rates of Reaction