Reaction rate

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The reaction rate of or the rate of reaction is the rate at which the reactants are converted into products. For example, the erosion of iron oxidation under the Earth's atmosphere is a slow reaction that can last for years, but the burning of cellulose in fire is a reaction that occurs in a fraction of a second. For most reactions, the velocity decreases as the reaction progresses.

Chemical Kinetics is part of the physical chemistry that studies the reaction rate. The concept of chemical kinetics is applied in many disciplines, such as chemical engineering, enzymes and environmental engineering.

Video Reaction rate

Definisi formal

Consider typical chemical reactions: a A b Ã, B -> p Ã, P q Q

The small letters ( a , b , p , and q ) represent stoichiometric coefficients, while capital letters represent reactants ( A and B) and products (P and Q).

Menurut buku Emas IUPAC, definisi laju reaksi r untuk reaksi kimia yang terjadi dalam sistem tertutup dalam kondisi isokorik, tanpa penumpukan intermediet reaksi, didefinisikan sebagai:

${\ displaystyle r = - {\ frac {1} {a}} {\ frac {d [\ mathrm {A}]} {dt}} = - {\ frac {1} {b}} {\ frac {d [\ mathrm {B}]} {dt}} = {\ frac {1} {p}} {\ frac {d [\ mathrm {P}]} {dt }} = {\ frac {1} {q}} {\ frac {d [\ mathrm {Q}]} {dt}}}  ÂÂ$

where [X] denotes the concentration of the substance X (= A, B, P or Q). The reaction rate usually has units of mol/L/s

The reaction rate is always positive. A negative sign is present to show that the reactant concentration decreases.) The IUPAC recommends that the time unit should always be the second. The reaction rate differs from the rate of increase of product P concentration by a constant factor (as opposed to a stoichiometric number) and to reactant A with minus opposite of the stoichiometric number. The stoichiometric figures are entered so that the specified level is independent of the reactants or the selected product species for the measurement. For example, if a = 1 and b = 3 then B is consumed three times faster than A, but v = -d [ A < i>]/dt = - (1/3) d [ B ]/dt is uniquely defined. An additional advantage of this definition is that for elementary and irreversible reactions, r is equal to the probability product overcoming the activation energy of the transition state and how many times per second the transition state is approached by the reactant molecule. When defined, for elementary and irreversible reactions, r is the rate at which a successful chemical reaction leads to the product.

The above definition applies only to single reactions , in closed systems of constant volume , assumed to be explicitly stated in the definition. If water is added to a pot of salty water, the salt concentration decreases, although there is no chemical reaction.

Untuk sistem terbuka, keseimbangan massa penuh harus diperhitungkan: in-out generation - consumption = akumulasi

${\ displaystyle F _ {\ mathrm {A} 0} -F_ {mathrm {A}} \ int _ {0} ^ {V} v \, dV = { \ frac {dN _ {\ mathrm {A}}} {dt}}}  ÂÂ$ ,

where F _A0 is the inlet flow rate A in molecules per second, F _A > v is the instantaneous reaction rate A (in the number of molar concentrations) in a certain differential volume, integrated across the system volume V at any given moment. When applied to a closed system at a constant volume previously considered, this equation reduces to:

${\ displaystyle r = {\ frac {d [A]} {dt}}}$ ,

where the concentration [A] is related to the number of molecules N _A by [A] Ã, = N A / N ₀ V span>. Here N ₀ is the Avogadro constant.

Untuk reaksi tunggal dalam sistem tertutup dari berbagai volume yang disebut laju konversi dapat digunakan, untuk menghindari konsentrasi penanganan. Ini didefinisikan sebagai turunan dari tingkat reaksi terhadap waktu.

${\ displaystyle r = {\ frac {d \ xi} {dt}} = {\ frac {1} {\ nu _ {i}}} {\ frac {dn_ {i}} {dt}} = {\ frac {1} {\ nu _ {i}}} {\ frac {d (C_ {i} V)} {dt}} = {\ frac {1} {\ nu _ {i}}} \ kiri (V {\ frac {dC_ {i}} {dt}} C_ {i} {\ frac {dV} {dt}} \ right)}  ÂÂ$

Here ? The _i is the stoichiometric coefficient for the substance i , equals a , b , and q in the typical reaction above. Also V is the volume of the reaction and C _i is the concentration of the substance i .

When byproducts or intermediate reactions are formed, IUPAC recommends the use of the terms the rate of appearance and the removal rate for the product and the reactants, correctly.

The reaction rate can also be determined on a basis that is not a reactor volume. When the catalyst is used, the reaction rate can be expressed at the weight of the catalyst (mol ^-1 s ^-1 ) or the surface area (mol m ^-2 soup <-1 ). If the base is a specific catalyst site that can be calculated strictly with the specified method, the rate is given in units of s ^-1 and is called the rotational frequency.

Maps Reaction rate

Influencing factor

• Nature of reaction : Some reactions are naturally faster than others. The number of species that react, their physical state (the particles that make up the solid move much slower than the gas particles or those in solution), the complexity of the reaction and other factors can greatly affect the rate of reaction.
• Concentration : The reaction rate increases with concentration, as explained by the rate law and is explained by the collision theory. As the reactant concentration increases, the collision frequency increases.
• Pressure : The gas reaction rate increases with pressure, which, in fact, equals the increase in gas concentration. The reaction rate increases to where there is less mole gas and decreases in the opposite direction. For the condensation phase reaction, the pressure dependence is weak.
• Order : The reaction sequence controls how the concentration of the reactant (or pressure) affects the reaction rate.
• Temperature : Usually a reaction at higher temperatures gives more energy into the system and increases the reaction rate by causing more collisions between particles, as explained by the collision theory. However, the main reason that the temperature increases the reaction rate is that more colliding particles will have the necessary activation energy resulting in a more successful collision (when the bond is formed between the reactants). The effect of temperature is illustrated by the Arrhenius equation.

For example, coal burns in the fireplace in the presence of oxygen, but not when it is stored at room temperature. The reaction is spontaneous at low and high temperatures but at room temperature, the speed is so slow that it can be ignored. An increase in temperature, as created by a match, allows a reaction to start and then heats itself, because it is exothermic. That applies to many other fuels, such as methane, butane, and hydrogen.

The reaction rate may depend on temperature ( non-Arrhenius ) or decrease with increasing temperature ( anti-Arrhenius ). Reactions without an activation barrier (eg, some radical reactions), tend to have anti-Arrhenius temperature dependence: the rate constant decreases with increasing temperature.

• Solvents : Many reactions occur in the solution and the properties of the solvent affect the reaction rate. The ionic strength also has an effect on the reaction rate.
• Electromagnetic radiation and light intensity : Electromagnetic radiation is a form of energy. Thus, it can speed up the rate or even make a spontaneous reaction because it gives the reactant particles with more energy. This energy in one way or another is stored in the reacting particles (can damage bonds, promote molecules to electronically or vibrationally excited states...) creating intermediate species that react easily. As the intensity of light increases, the particles absorb more energy and therefore the reaction rate increases.

For example, when methane reacts with chlorine in the dark, the reaction rate is very slow. This can be accelerated when the mixture is placed under diffuse light. In bright sunlight, his reaction is explosive.

• Catalyst : The presence of a catalyst improves the reaction rate (in forward and backward reactions) by providing an alternative pathway with a lower activation energy.

For example, platinum catalyzes the combustion of hydrogen with oxygen at room temperature.

• Isotop : The kinetic isotope effect consists in different reaction rates for the same molecule if it has different isotopes, usually the hydrogen isotope, because of the relative mass difference between hydrogen and deuterium.
• Surface Area : In reaction to the surface, which occurs eg during heterogeneous catalysis, the reaction rate increases as the surface area. That's because more solid particles are exposed and can be exposed to reactant molecules.
• Stirring : Stirring can have a powerful effect on the reaction rate for heterogeneous reactions.
• Limit of diffusion: Some reactions are limited by diffusion.

All factors affecting the reaction rate, except for the concentration and reaction sequence, are taken into account in the reaction rate coefficients (coefficients in the rate equation).

src: chem.libretexts.org

Rate equation

Untuk reaksi kimia a A b  B -> p  P q Q, persamaan laju atau hukum laju adalah ekspresi matematis yang digunakan dalam kinetika kimia untuk menghubungkan laju reaksi dengan konsentrasi setiap reaktan. Seringkali jenisnya:

${\ displaystyle \, r = k (T) [\ mathrm {A}] ^ {n} [\ mathrm {B}] ^ {m}}  ÂÂ$

For gas phase reactions, the rate is often expressed by partial pressure.

In this equation k ( T ) is the reaction rate coefficient or constant rate , although it is not a constant, all parameters affecting the reaction rate, except for concentrations, are explicitly taken into account. Of all the parameters affecting the reaction rate, temperature is usually the most important and noted by the Arrhenius equation.

The exponents n and m are called the reaction commands and depend on the reaction mechanism. For the elementary reaction (one step), the order of each reactant is the same as the stoichiometric coefficient. For complex reactions (multistep), however, these are often incorrect and the rate equations are determined by detailed mechanisms, as illustrated below for H ₂ and NO reactions.

For elementary reactions or reaction steps, the sequence and stoichiometric coefficients are both equal to the molecules or number of participating molecules. For unimolecular reactions or rate steps proportional to the concentration of reactant molecules, so the law of rate is the first order. For bimolecular or step reactions, the number of collisions is proportional to the product of two reactant concentrations, or second order. The thermolecular step is predicted to be third order, but also very slow because the simultaneous collision of three molecules is rare.

By using a mass balance for the system in which the reaction occurs, the expression for the rate of change of concentration can be reduced. For closed systems with constant volume, such expressions can look like

$Â\; ÂÂ\; Â\; Â\; ÂÂ\; Â\; Â\; Â\; Â{\displaystyle Â\; Â\; Â\; Â\; Â\; Â\; Â\; ÂÂ\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\frac{Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; ÂdÂ\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â[Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; ÂÂ\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\mathrm{P}Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; ÂÂ\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â]Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â}{Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; ÂdÂ\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; ÂtÂ\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â}Â\; Â\; Â\; Â\; Â\; Â\; Â\; ÂÂ\; Â\; Â\; Â\; Â\; Â\; Â\; Â=Â\; Â\; Â\; Â\; Â\; Â\; Â\; Âk()\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; ÂTÂ\; Â\; Â\; Â\; Â\; Â\; Â\; Â)Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â[Â\; Â\; Â\; Â\; Â\; Â\; Â\; ÂÂ\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\mathrm{A}Â\; Â\; Â\; Â\; Â\; Â\; Â\; ÂÂ\; Â\; Â\; Â\; Â\; Â\; Â\; Â{]}^{Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; ÂnÂ\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â}Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â[Â\; Â\; Â\; Â\; Â\; Â\; Â\; ÂÂ\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\mathrm{B}Â\; Â\; Â\; Â\; Â\; Â\; Â\; ÂÂ\; Â\; Â\; Â\; Â\; Â\; Â\; Â{]}^{Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; ÂmÂ\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â\; Â}Â\; Â\; Â\; Â\; Â\; Â}Â\; Â\; Â\; Â\{\backslash \; Displaystyle\; \{\backslash \; frac\; \{d\; [\backslash \; mathrm\; \{P\}]\}\; \{dt\}\}\; =\; k\; (T)\; [\backslash \; mathrm\; \{A\}]\; ^\; \{\; n\}\; [\backslash \; mathrm\; \{B\}]\; ^\; \{m\}\}\; Â\; Â$

Examples of complex reactions: Hydrogen and nitrate oxide reactions

For reaction

2a, H ₂ (g) 2Ã,Tidak (g) -> N ₂ > O (g)

persamaan laju yang diamati (atau ekspresi laju) adalah: ${\ displaystyle r = k [{\ ce {H2}}] [{\ ce {NO}}] ^ {2} \,}  ÂÂ$

As for many reactions, the experimental rate equation not only reflects the stoichiometric coefficients in the overall reaction: This is the third order in its entirety: the first sequence in H ₂ and second in No, although the stoichiometric coefficient of both reactants is equal to 2.

In chemical kinetics, the overall reaction rate is often described using a mechanism consisting of a number of basic steps. Not all of these steps affect the rate of reaction; usually the basic steps slowest to control the reaction rate. For this example, the possible mechanisms are:

1. 2Ã, (g)? N ₂ O ₂ (g) (fast equilibrium)
2 2 2 2 2 2 O (slow)
2. N ₂ OH ₂ -> N ₂ H ₂ O ( fast)

Reaksi 1 dan 3 sangat cepat dibandingkan dengan yang kedua, jadi reaksi lambat 2 adalah langkah penentuan laju. Ini adalah reaksi elementer bimolekuler yang kursnya diberikan oleh persamaan orde kedua:

${\ displaystyle r = k_ {2} [{\ ce {H2}}] [{\ ce {N2O2}}] \,}  ÂÂ$ ,

where k ₂ is the constant rate for the second step.

However N ₂ O ₂ is an unstable intermediary whose concentration is determined by the fact that the first step is in equilibrium, so [N ₂ = K ₁ ² K ₁ is the equilibrium constant of the first step. The substitution of this equation in the previous equation leads to the expressed equation expressed in the original reactant

${\ displaystyle r = k_ {2} K_ {1} [{\ ce {H2}}] [{\ ce {NO}}] ^ {2 } \,} Â Â$

This corresponds to the form of the observed rate equation if it is assumed that k sub> 1 . In practice, the rate equations are used to suggest possible mechanisms that predict the rate equations according to the experiment.

The second molecule H ₂ does not appear in the rate equation because it reacts in step three, which is a quick step after the rate stepper, so it does not affect the overall reaction rate.

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Temperature dependency

Setiap koefisien laju reaksi k memiliki ketergantungan suhu, yang biasanya diberikan oleh persamaan Arrhenius:

${\ displaystyle k = Ae ^ {- {\ frac {E _ {\ mathrm {a}}} {RT}}}}  ÂÂ$

E _a is the activation energy and R is the gas constant. Because at a temperature of T the molecule has the energy given by the Boltzmann distribution, one can expect the number of collisions with energies greater than E _a to be proportional to e ^{- E _a RT} . A is a pre-exponential or frequency factor.

The values ​​for A and E _a depend on the reaction. There is also a more complex equation, which illustrates temperature dependence from other speed constants that do not follow this pattern.

Chemical reactions occur only when the particles react to collide. However, not all collisions are effective in causing a reaction. The product is only formed when the colliding particle has a certain minimum energy called threshold energy. As a rule of thumb, the reaction rate for multiple reactions doubled for each temperature increase of 10 degrees Celsius. For the given reaction, the ratio of constant rates at higher temperatures to the rate constants at lower temperatures is known as the temperature coefficient. ( Q ). Q ₁₀ is usually used as a ratio of the rate constant which is 10 Â ° C.

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Dependency on pressure

The pressure dependence of the constant rate for the viscous-phase reaction (ie, when the reactants and products are solids or liquids) are usually quite weak in the range of pressures normally encountered in industries that are neglected in practice.

The pressure dependence of the rate constants is associated with the volume of activation. For reactions that take place through the complex state of activation:

A B? | A? B | ^? -> P

volume aktivasi,? V ^? , adalah:

${\ displaystyle \ Delta V ^ {\ ddagger} = {\ bar {V}} _ {\ ddagger} - {\ bar {V}} _ {\ mathrm {A }} - {\ bar {V}} _ {\ mathrm {B}}}  ÂÂ$

where V? shows the partial molar volume of a species and? shows state-activated complexes.

Untuk reaksi di atas, seseorang dapat mengharapkan perubahan konstanta laju reaksi (berdasarkan pada fraksi mol atau pada konsentrasi-molar) dengan tekanan pada suhu konstan untuk menjadi:

${\ displaystyle \ left ({\ frac {\ partial \ ln k_ {x}} {\ partial P}} \ right) _ {T} = - {\ frac { \ Delta V ^ {\ ddagger}} {RT}}}  ÂÂ$

In practice, this problem can be complicated because the volume of partial molar and activation volume can be a function of pressure.

Reactions can increase or decrease their rates with pressure, depending on the value? V ^? . As an example of the possibility of a pressure effect, some organic reactions are shown to double the reaction rate when the pressure rises from the atmosphere (0.1 MPa) to 50 MPa (which gives V ^? Ã, = Ã, -0,025 L/mol).

src: chem.libretexts.org

See also

• Value of solution
• Dilution (equation)
• diffusion controlled reaction
• Still status approach
• The theory of collisions and transition states is a chemical theory that attempts to predict and explain reaction rates.
• Isothermal microcalorimetry

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Note

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External links

• Chemical Kinetics, reaction rate, and order (requires flash player)
• Reaction kinetics, important law-level example (lecture with audio).
• Rates of reaction
• Overview of Bimolecular Reactions (Reactions involving two reactants)

Source of the article : Wikipedia