Chemical reaction
A chemical reaction is a process that results in the interconversion of chemical substances [1]. The substance or substances initially involved in a chemical reaction are called reactants. Chemical reactions are characterized by a chemical change and it yields one or more products which are different from the reactants. Classically, chemical reactions encompass changes that strictly involve the motion of electrons in the forming and breaking of chemical bonds, although the general concept of a chemical reaction, in particular the notion of a chemical equation, is applicable to transformations of elementary particles, as well as nuclear reactions.
Many different chemical reactions are used in combinations in chemical synthesis in order to get a desired product. In biochemistry, series of chemical reactions form metabolic pathways, since straight synthesis of a product would be energetically impossible in conditions within a cell. Chemical reactions are also divided into organic reactions and inorganic reactions.
Reaction types
Chemical reactions may be classified in different ways depending on the particular aspect considered for elaborating the division, or on the branch of Chemistry which the classification originates from. Some examples of widely used terms for describing common kinds of reactions are:
- Isomerisation, in which a chemical compound undergoes a structural rearrangement without any change in its net atomic composition; see stereoisomerism
- Direct combination or synthesis, in which two or more chemical elements or compounds unite to form a more complex product:
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- 2H2 (g) + O2 (g) → 2H2O (l)
- Chemical decomposition or analysis, in which a compound is decomposed into smaller compounds or elements:
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- 2H2O (l) → 2H2 (g) + O2(g)
- Single displacement or substitution, characterized by an element being displaced out of a compound by a more reactive element:
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- 2Na(cr) + 2HCl (aq) → 2NaCl (aq) + H2 (g)
- Double displacement or coupling substitution , in which two compounds in aqueous solution (usually ionic) exchange elements or ions to form different compounds:
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- NaCl (aq) + AgNO3 (aq) → NaNO3 (aq) + AgCl (s)
- Combustion, in which any combustible substance combines with an oxidizing element, usually oxygen, to generate heat and form oxidized products. The term combustion is used usually only for reactions that destroy complex molecules, i.e. a controlled oxidation of a single functional group is not combustion.
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- C10H8 (g) + 14O2 (g) → 10CO2 (g) + 4H2O (g) + heat
- CH2S + 6 F2 → CF4 + 2 HF + SF6 + heat
Some branches of chemistry include any minor changes in chemical conformation in the reaction types, while others consider these changes merely as physical properties of a compound.
The collision of more than two particles into the ordered structure necessary to perform chemical transformations is extremely unlikely; which is why ternary reactions in practice are not observed. A chemical reaction may require three or more reagents, but the process can generally be decomposed into a stepwise series or a set of stepwise reactions of the above.
The large diversity of chemical reactions makes it difficult to establish simple criteria for functional (as opposed to mechanistic) classification. However, some kinds of reactions have similarities which make it possible to define some larger groups. A few examples are:
- Organic reactions encompass several different kinds of reactions involving compounds which have carbon as the main element in their molecular structure. These reactions occur mostly according to, within, by, or via functional groups.
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- Petrochemical reactions are often distinguished from other organic reactions.
- Redox reactions involve augmenting or decreasing the electrons associated with a particular atom. according to its oxidation number.
- Combustion, in which a substance reacts with an oxidizing element, such as oxygen gas.
Reactions are also classified according to their mechanism:
- Reactions of ions, e.g. disproportionation of hypochlorite
- Reactions with reactive ionic intermediates, e.g. reactions of enolates
- Radical reactions, e.g. combustion at high temperature
- Reactions of carbenes
Thermochemistry
See main article: Thermochemistry.
Thermochemistry deciphers whether a specific chemical reaction can or cannot occur. Thermodynamics (or what is now known as equilibrium thermodynamics) understands the reaction in terms of the initial and final states of the reaction mixture.
Reactions very seldom occur directly. Usually, reactants must collide to form an activated complex. This complex has a higher internal energy than the original reactants combined, having gained some from the kinetic energy of the reactant substances’ collision. This energy allows for the rearrangement of bonds which constitutes the reaction. In some reactions, the reactants may pass through several reactive intermediates before becoming products.
Thermodynamics does not attempt to figure out the process by which a reaction occurs. This field of study is taken up by the field of chemical kinetics. Another question “How fast is the reaction?” is also left completely unanswered by it. Chemical kinetics attempts to put all these phenomena into perspective.
Chemical equilibrium
Every chemical reaction is, in theory, reversible. In a forward reaction the substances defined as reactants are converted to products. In a reverse reaction products are converted into reactants.
Chemical equilibrium is the state in which the forward and reverse reaction rates are equal, thus preserving the amount of reactants and products. However, a reaction in equilibrium can be driven in the forward or reverse direction. This is done by changing the reaction conditions such as temperature or pressure. Le Chatelier’s principle can be used to predict whether products or reactants will be formed.
Although all reactions are reversible to some extent, some reactions can be classified as irreversible. An irreversible reaction is one that “goes to completion.” This phrase means that nearly all of the reactants are used to form products. These reactions are very difficult to reverse even under extreme conditions.
Exothermic reactions
A sketch of an exothermic reaction
According to energy balance criteria, that is, chemical reaction equilibria criteria, any closed system will tend to minimize its free energy. Without any outside influence, any reaction mixture, too, will try to do the same. For many cases, an analysis of the enthalpy of the system will give a decent account of the energetics of the reaction mixture. The enthalpy of a reaction is calculated using standard reaction enthalpies and the Hess’ law of constant heat summation. Many of these enthalpies may be found in beginners’ books on thermodynamics. For example, consider the combustion of methane in oxygen:
- CH4 + 2 O2 → CO2 + 2 H2O
By calculating the amounts of energy required to break all the bonds on the left (”before”) and right (”after”) sides of the equation using collected data, it is possible to calculate the energy difference between the reactants and the products. This is referred to as ΔH, where Δ (Delta) means difference, and H stands for enthalpy, a measure of energy which is equal to the heat transferred at constant pressure. ΔH is usually given in units of kilojoules (kJ) or in kilocalories (kcal).
If ΔH is negative for the reaction, then energy has been released often in the form of heat. This type of reaction is referred to as an exothermic reaction (literally, outside heat, or throwing off heat). An exothermic reaction is more favourable and thus more likely to occur. An example reaction is combustion, known from everyday experience, since burning gas in air produces heat.
Endothermic reactions
A reaction may have a positive ΔH. If a reaction has a positive ΔH, it consumes energy as the reaction moves towards completion. This type of reaction is called an endothermic reaction (literally, inside heat, or absorbing heat).
The above rule, “Exothermic reactions are favourable”, is usually true. However, there may be situations where exothermic reactions may not be favourable. This happens when the stability obtained due to loss of enthalpy is off set by a corresponding decrease in entropy (a measure of disorder). The exact rule is that a reaction is favourable when the Gibbs free energy of that reaction is negative where ΔG = ΔH − TΔS; ΔG being the change in Gibbs free energy, ΔH being the change in enthalpy, and ΔS is the change in entropy
A reaction is called spontaneous if its thermodynamically favoured, by that meaning that it causes a net increase on entropy. Spontaneous reactions (in opposition to non-spontaneous reactions) do not need external perturbations (such as energy supplement) to happen. In a system at chemical equilibrium, it is expected to have larger concentrations of the substances formed by the spontaneous direction of the process.
Thus, in a global isolated system (which it strictly isn’t, see entropy), spontaneous reactions may be understood to occur without human interference. Most spontaneus reactions in this system are exothermic (such as rusting) or metamorphism, thus increasing the global entropy, though photosynthesis is an important exception (in a global system).
Chemical kinetics
See main article: Chemical kinetics.
The rate of a chemical reaction is a measure of how the concentration of the involved substances changes with time. Analysis of reaction rates is important for several applications, such as in chemical engineering or in chemical equilibrium study. Rates of reaction depends basically on:
- Reactant concentrations, which usually make the reaction happen at a faster rate if raised,
- Surface Area, the amount of the substance being used,
- Pressure, By increasing the pressure, you squeeze the molecules together so you will increase the frequency of collisions between the molecules.
- Activation energy, which is defined as the amount of energy required to make the reaction start and carry on spontaneously. Higher activation energy implies that a reaction will be harder to start and, therefore, slower.
- Temperature, which hastens reactions if raised, because higher temperature means that the involved species will have more energy, thus making the reaction easier to happen,
- The presence or absence of a catalyst. Catalysts are substances which change the pathway (mechanism) of a reaction which in turn increases the speed of a reaction by lowering the activation energy needed for the reaction to take place. A catalyst is not destroyed or changed during a reaction, so it can be used again.
Reaction rates are related to the concentrations of substances involved in reactions, as quantified by the law of mass action. Reactions whose rates are independent of reactant concentrations are called zero-order reactions.
See also
- List of reactions
- List of publications in chemistry
External links
References
- ↑ IUPAC Gold Book Definition
cation is fairly common.
