Introduction:

Chemical kinetics, as well termed as reaction kinetics, is basically the study of rates of chemical processes. Chemical kinetics comprises investigations of how different experimental conditions can affect the speed of a chemical reaction and yield information regarding the reaction's method and transition states, and also the construction of mathematical models that can illustrate the characteristics of the chemical reaction. In the year 1864, Peter Waage and Cato Guldberg pioneered the growth of chemical kinetics by formulating the law of mass action that states that the speed of a chemical reaction is proportional to the quantity of the reacting substances.

The chemical kinetics mainly deals with the experimental determination of the reaction rates from which rate laws and rate constants are derived. Relatively simple rate laws exist for zero order reactions (for which the reaction rates are independent of concentration), first order reactions, and second order reactions, and can be derived for others. The elementary reactions follow the law of mass action; however the rate law of stepwise reactions has to be derived via joining the rate laws of the different elementary steps, and can become instead complex. In consecutive reactions, the rate-determining step frequently finds out the kinetics. In consecutive first order reactions, a steady state approximation can ease the rate law. The activation energy for a reaction is experimentally found out via the Arrhenius equation and the Eyring equation. The major factors which affect the reaction rate comprise: the physical state of the reactants, the concentrations of reactants, the temperature at which the reaction takes place, and whether or not any catalysts are present in the reaction.

Factors influencing reaction rate:

1) Nature of the reactants:

Based on what substances are reacting, the reaction rate differs. Acid or base reactions, the formation of salts, and ion exchange are fast reactions. Whenever covalent bond formation occurs between the molecules and whenever large molecules are made up, the reactions tend to be extremely slow. Nature and strength of bonds in reactant molecules highly affect the rate of its transformation into products.

2) Physical state:

The physical state (that is, solid, liquid and gas) of a reactant is as well a significant factor of the rate of change. Whenever reactants are in the similar phase, as in the aqueous solution, thermal motion brings them to contact. Though, whenever they are in different stages, the reaction is limited to the interface between the reactants. The reaction can take place only at their area of contact; in the case of liquid and gas, at the surface of the fluid. Vigorous shaking and stirring might be required to bring the reaction to completion. This signifies that the more finely divided a solid or liquid reactant the more its surface area per unit volume and the more contact it with the other reactant, therefore the faster the reaction. To make an analogy, for illustration, whenever one begins a fire, one utilizes wood chips and small branches - one doesn't begin with large logs right away. In the organic chemistry, on water reactions are the exception to the rule that homogeneous reactions occur faster than the heterogeneous reactions.

3) Concentration:

Reactions are due to the collisions of reactant species. The frequency by which the molecules or ions collide based on their concentrations. The more crowded the molecules are, the more possible they are to collide and react with one other. Therefore, an increase in the concentrations of the reactants will generally outcome in the corresponding increase in the reaction rate, whereas a decrease in the concentrations will generally encompass a reverse effect. For illustration combustion that take place in air (that is, 21% oxygen) will take place more rapidly in pure oxygen.

4) Temperature:

Temperature generally consists of a major effect on the rate of a chemical reaction. Molecules at a higher temperature encompass more thermal energy. However, collision frequency is higher at higher temperatures; this alone contributes merely a very small proportion to the increase in rate of reaction. Much more significant is the fact that the proportion of reactant molecules by adequate energy to react (that is, energy greater than activation energy: E > E_a) is significantly higher and is illustrated in detail by the Maxwell-Boltzmann distribution of molecular energies.

The 'rule of thumb' that the rate of chemical reactions doubles for each and every 10 °C temperature increase is a general misconception. This might have been generalized from the special case of biological systems, where 'α' (temperature coefficient) is frequently between 1.5 and 2.5.

Reaction kinetics can as well be studied by a temperature jump approach. This comprises by employing a sharp increase in temperature and examining the relaxation time of the return to equilibrium. A principally helpful form of temperature jump apparatus is a shock tube that can quickly jump a gas's temperature by more than 1000 degrees.

5) Catalysts:

A catalyst is a substance which accelerates the rate of a chemical reaction however remains chemically unchanged afterwards. The catalyst raises the rate of the reaction by providing a different reaction method to take place having lower activation energy. In autocatalysis, a reaction product is itself a catalyst for that reaction leading to the positive feedback. Proteins which act as catalysts in the biochemical reactions are termed as enzymes. Michaelis-Menten kinetics illustrates the rate of enzyme mediated reactions. A catalyst doesn't influence the position of the equilibria, as the catalyst speeds up the forward and backward reactions equally.

In some organic molecules, specific substituent's can encompass an influence on reaction rate in neighboring group participation.

6) Pressure:

Increasing the pressure in the gaseous reaction will raise the number of collisions between the reactants, increasing the rate of reaction. This is since the activity of a gas is directly proportional to the partial pressure of the gas. This is identical to the effect of raising the concentration of a solution.

Moreover, to this straightforward mass-action effect, the rate coefficients themselves can modify due to pressure. The rate coefficients and products of numerous high-temperature gas-phase reactions change whenever an inert gas is added to the mixture; variations on this effect are termed as fall-off and chemical activation. Such phenomena are due to exothermic and endothermic reactions taking place faster than heat transfer, causing the reacting molecules to encompass non-thermal energy distributions (that is, non-Boltzmann distribution). Increasing the pressure, increases the heat transfer rate between the reacting molecules and the rest of the system, decreasing this effect.

Condensed-phase rate coefficients can as well be influenced by (very high) pressure; this is a totally different effect than fall-off or chemical-activation. This is frequently studied by employing diamond anvils.

Kinetic reaction can as well be studied via a pressure jump approach. This comprises making fast changes in pressure and examining the relaxation time of the return to equilibrium.

Experimental methods:

Fast reactions:

For faster reactions, the time needed to mix the reactants and bring them to a particular temperature might be comparable or longer than the half-life of the reaction. Special methods to begin fast reactions devoid of slow mixing step comprise:

a) Stopped flow processes that can decrease the mixing time to the order of a millisecond.

b) Chemical relaxation processes like temperature jump and pressure jump in which a pre-mixed system primarily at equilibrium is perturbed via rapid heating or depressurization in such a way that it is no longer at equilibrium, and the relaxation back to equilibrium is examined. For illustration, this process has been employed to study the neutralization H₃O⁺ + OH^- by a half-life of 1 μs or less under ordinary conditions.

c) Flash photolysis, in which a laser pulse generates highly, excited species such as free radicals, whose reactions are then studied.

Equilibrium:

As chemical kinetics is concerned by the rate of a chemical reaction, thermodynamics finds out the extent to which reactions take place. In a reversible reaction, chemical equilibrium is reached whenever the rates of the forward and reverse reactions are equivalent (that is, the principle of detailed balance) and the concentrations of the reactants and products no longer change. This is illustrated by, for illustration, the Haber-Bosch method for combining nitrogen and hydrogen to generate ammonia. Chemical clock reactions like the Belousov-Zhabotinsky reaction illustrate that component concentrations can oscillate for a long time prior to at last attaining the equilibrium.

Chemical equilibrium: 

Whenever a chemical reaction in the forward direction is happening at the similar rate as the chemical reaction in the reverse direction; whenever the overall concentrations or partial pressures of chemicals in the reaction are not changing.

Equilibrium constant (K): 

A number which associates the forward and reverse rates of reaction. When K = 1, the forward and reverse rates are equivalent. When K > 1, then the forward reaction is faster than the reverse reaction and if K < 1, then the reverse reaction is faster.

Kinetics and the Equilibrium Constant:

Consider the given reaction:

N₂ (g) + 3 H₂ (g) ↔ 2 NH₃ (g)

The rate laws for the forward and reverse reactions are as follows:

Forward: R_f = k_f [N₂] [H₂]³

Reverse: R_r = k_r [NH₃]²

At equilibrium condition, R_f = R_r, that signifies:

By using the algebra, we can get the rate constants on one side of the equation and the concentration terms on the other:

k_f/k_r = ([NH₃]²)/([N₂][H₂])³

It will be noted that this is precisely similar by means of the equilibrium expression:

K_eq = ([NH₃])²/([N₂][H₂])³

Thus, K_eq = k_f/k_r

Free energy:

In common terms, the free energy change (ΔG) of a reaction finds out whether a chemical change will occur, however kinetics illustrates how fast the reaction is. A reaction can be exothermic and encompass a very positive entropy change however will not occur in practice if the reaction is too slow. If a reactant can generate two different products, the thermodynamically most stable one will in common form, apart from in special conditions whenever the reaction is stated to be under kinetic reaction control. The Curtin-Hammett principle applies whenever finding out the product ratio for two reactants interconverting quickly, each and every going to a different product. This is possible to make predictions regarding reaction rate constants for a reaction from free-energy relationships.

The kinetic isotope effect is the difference in the rate of a chemical reaction whenever an atom in one of the reactants is substituted by one of its isotopes.

Chemical kinetics gives information on residence time and heat transfer in the chemical reactor in chemical engineering and the molar mass distribution in the polymer chemistry.

Applications:

The mathematical models which illustrate chemical reaction kinetics give chemists and chemical engineers by tools to better comprehend and illustrate chemical methods like food decomposition, stratospheric ozone decomposition, microorganism growth and the complex chemistry of biological systems. Such models can as well be employed in the design or modification of chemical reactors to optimize product yield, more proficiently separate products, and remove environmentally injurious by-products. Whenever performing the catalytic cracking of heavy hydrocarbons into gasoline and light gas, for illustration, kinetic models can be employed to determine the pressure and temperature at which the highest outcome of heavy hydrocarbons into gasoline will take place. Kinetics is as well a fundamental feature of chemistry.

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