Chemical reactions happen spontaneously



 

A chemical reaction is the process in which chemical compounds (the “starting materials” or “reactants”) are converted into other chemical compounds (the “products”) by the atoms of the reactants forming different bonds. In the process, energy is converted (endothermic reaction) or released (exothermic reaction), which is known as the enthalpy of reaction. The reaction is subject to the laws of energy conservation and entropy. The process ends when chemical equilibrium is reached or the reactants have been completely converted.

The course of chemical reactions is shown as a reaction scheme or reaction equation:

Basics

During the chemical reaction, chemical bonds are broken and new ones made. The specific properties of the raw materials disappear. The newly created substances have other specific properties, such as: B. color, smell, viscosity, density, fixed point, boiling point, optical activity, etc.

Each reaction has a different reaction speed.

Their mechanism proceeds according to principles such as substance association (synthesis, addition reaction), substance breakdown (analysis, elimination reaction) and substance rearrangement (exchange and substitution reactions: redox reaction, acid-base reaction (protolysis), complex formation reaction / ligand exchange, precipitation reaction / formation of precipitates, organic substitution (radical , electrophilic, nucleophilic).

Differentiation from physical processes

 

When a chemical reaction arises at least a newer Substance (product) - in physical processes, it is not the substance-specific properties that change, but only physical properties such as heat content, physical state and expansion. However, all chemical reactions are also accompanied by physical changes in the substances. The release or absorption of energy, changes in the physical state or the color can be observed.

(The distinction between chemical reaction and physical process cannot be clearly defined in exceptional cases; for example, the dissolution of sodium in liquefied ammonia gas is considered a chemical reaction, since the liquid turns blue. After the ammonia has evaporated, however, sodium remains and not - how in the case of a chemical reaction would be expected - some combination of sodium and nitrogen.)

Regardless of what has been said here, reactions can also be described using physical means.

Substance conversion

In a chemical reaction, chemical substances are converted into other substances. Depending on whether elements or compounds occur in the starting materials or products, a distinction is made between two basic types of reaction:

synthesis

Substance union - two elements come together to form a connection:

A + B → AB

Examples:

  • The Elements magnesium and Iodine form the connection Magnesium iodide
  • The elements hydrogen and fluorine react to hydrogen fluoride:

analysis

Material decomposition - A compound is broken down into its elements:

AB → A + B

Example: Water is broken down into oxygen and hydrogen by electrolysis or at 2000 ° C:

All other types of reactions (regrouping of substances) can be put together from analysis and synthesis:

Easy implementation

Here one element reacts with a connection, whereby another element and a new connection arise:

A + BC → AB + C

 

Example: Elemental chlorine releases elemental iodine from sodium iodide and sodium chloride (table salt) is formed:

The resulting iodine is shaken out with the solvent hexane (extraction) and turns it pink-violet (detection reaction).

Chlorinated water reacts similarly with sodium bromide solution; the added hexane turns orange instead of pinkish-violet.

A simple implementation can be imagined as a combination of the two partial reactions analysis and synthesis:

Analysis: BC → B + C Synthesis: A + C → AC Sum equation: A + BC → AC + B

Double implementation

Here two connections react with each other in such a way that two new connections are created:

AB + CD → AD + CB

Example: If a solution of magnesium iodide is mixed with a solution of lead chloride, yellow lead iodide precipitates and magnesium chloride remains in the solution.

A double conversion can be imagined as composed of two analysis and two synthesis reactions:

Analysis 1: AB → A + B Analysis 2: CD → D + C Synthesis 1: A + D → AD Synthesis 2: B + C → CB Sum equation: AB + CD → AD + CB

More options

Donor-acceptor principle

Acid-base reactions, redox reactions and complex formation reactions can be traced back to the donor-acceptor principle. In these cases the products are created through the exchange of elementary particles between the starting materials.

Acid-base reactions

These reactions are based on an exchange of protons between the starting materials. They can be viewed as a special case of double implementation (see above). An acid and a base are always used as starting materials. The acid as a proton donor releases at least one proton to the base as a proton acceptor.

Example: Hydrochloric acid is formed when hydrogen chloride gas is dissolved in water. The ampholyte water serves as the base.

Acid-base reactions can be broken down into two partial reactions, both of which are referred to as protolysis:

  1. Proton release:
  2. Proton uptake:

Example: Hydrogen chloride gas reacts with ammonia gas to form solid ammonium chloride (salmia).

Example: Ammonia reacts with the ampholyte water, which now functions as an acid, to form a basic solution.

A special case of the acid-base reaction is neutralization, in which an acidic solution reacts with a basic solution to form a neutral solution. The actual neutralizing reaction consists in the fact that oxonium (hydronium) ions of the acid react as proton donors with hydroxide ions of the base as proton acceptors to form neutral water:

Displacement reactions: The stronger acid (or base) displaces the weaker acid (or base) or its gaseous anhydride from their salts. (Examples: sulfuric acid releases sulfur dioxide from sulphites and hydrogen chloride from chlorides, concentrated hydrochloric acid releases hydrogen sulphide from sulphides and hydrogen cyanide from cyanides; caustic soda releases ammonia from amines and ammonium compounds.

Redox reactions

These reactions are based on an exchange of electrons between the starting materials. They can be the basis of all four basic forms of chemical reactions. A reducing agent and an oxidizing agent are always used as starting materials. The reducing agent as electron donor releases at least one electron to the oxidizing agent as electron acceptor.

Example: Elemental copper is deposited on a zinc rod (reducing agent) that is immersed in a solution with copper (II) ions (oxidizing agent) (see cementation). The solution is enriched with zinc (II) ions.

Redox reactions can be divided into the partial reactions oxidation and reduction disassemble:

  1. Oxidation: (electronsdelivery)
  2. Reduction: (electronsadmission)

Burns are redox reactions that always have elemental oxygen as an oxidizing agent.

When extracting metals from ores, carbon or carbon monoxide is used as a reducing agent (see iron extraction).

Complex formation reactions

In a complex formation reaction, electrons are also exchanged - but here they are made available to a central atom (usually a cation as Lewis acid) by molecules or ions in the form of lone pairs of electrons (the ligand as Lewis base. This is how copper (II) reacts, for example -ions with ammonia solution:

Conservation laws in chemical reactions

  1. Preservation of the elements: Educts and products contain the same elements. No elements can arise or disappear. (This law of conservation was formulated by Daniel Sennert in 1618. It means a rejection of the alchemists' claim to be able to produce gold from non-gold-containing materials, provided that this project is understood physically.)
  2. The Law of Conservation of Mass: The sum of the mass of the starting materials is equal to the sum of the mass of the products. (1. Basic Law of Chemistry, Joachim Jungius, 1662 and Michail Wassiljewitsch Lomonossow, 1748).
    Example: 1.00 g iron and 0.75 g sulfur react completely to form 1.75 g iron sulfide (FeS).
  3. The Law of constant proportions states that chemical compounds contain the elements in an unchangeable mass ratio that is characteristic of the respective compound. (2nd Basic Law of Chemistry, Jeremias Benjamin Richter, 1792 and Joseph-Louis Proust, 1799).
    Example: When iron reacts with sulfur to form iron sulfide, only a mixture of iron and sulfur in a mass ratio of 1: 1.75 reacts completely. In the case of iron disulfide (iron pebbles), the mass ratio is 1: 0.875.
  4. The Law of Multiple Proportions means that in different compounds that contain the same elements, the mass ratios are themselves in the ratio of whole numbers. (3rd Basic Law of Chemistry, John Dalton, 1808)
    Example: The mass ratio of iron sulfide (1.75) is related to the mass ratio of iron disulfide (0.875) as 2: 1

Conclusions from this:

  1. The Principle of constant volume proportions means that in reactions in the gas phase, the volumes of the educts and the volumes of the products are in the ratio of whole numbers. (derived from the 2nd and 3rd Basic Laws of Chemistry and Avogadro's Law on constant molar volumes, Joseph Louis Gay-Lussac and Alexander von Humboldt, 1808)
    Example: In the synthesis of ammonia from the elements, the volumes of nitrogen, hydrogen and ammonia are in a ratio of 1: 3: 2.
  2. Educts and products contain the same number of element atoms (cf. 1. + 2. Basic Law of Chemistry). This is achieved in a reaction equation with given formulas of the substances by suitable choice of the coefficients.
  3. The sum of the electrical charges of the educts is equal to the sum of the electrical charges of the products (charge equalization).

Energetic considerations

 

Every chemical reaction takes place with the participation of energy, since the loosening and tying of chemical bonds is connected with energy conversions.

Involved forms of energy

Depending on the type of energy involved, different types of reactions can be distinguished:

  • At thermochemical reactions heat is absorbed from the environment (endothermic reactions) or given off to the environment (exothermic reactions). In principle, all chemical reactions can be activated by supplying thermal energy, just as fundamentally, with every energy conversion during a reaction, heat occurs as a "friction loss".
examples for exothermic Reactions: The synthesis of water in the oxyhydrogen reaction releases 242 kJ per mole of (redox reaction) energy under standard conditions - with modern explosives this can be considerably more.
examples for endothermic Reactions: Michael Faraday split water into its elements by guiding water vapor through red-hot iron pipes, Harold Urey created organic molecules from a reducing gas mixture by adding energy.

 

  • Photochemical reactions are either triggered by light (examples: photosynthesis, bromination of alkanes, curing of plastics by UV light in dental technology) or they take place under the appearance of light (example: luminol reaction)
  • Some reactions can also be triggered mechanically, such as the decomposition of TNT.

Response path and energy balance

The thermodynamics of a reaction describe the course of a reaction from an energetic point of view.

A chemical system strives to adopt a state that is as low in energy as possible (enthalpy minimum) and the highest possible degree of disorder (entropy maximum).

Course of an exothermic reaction

The starting materials are initially in a metastable state. By briefly supplying a certain amount of energy, the activation energy (activation enthalpy), the system is lifted into the unstable state. Activation sets the reaction in motion and runs independently without any additional energy supply. In the overall balance, the chemical system gives off energy to the environment; it is referred to as the enthalpy of reaction. The products are now in a stable condition. (For the stability of systems see also system properties)

  Legend:
H: enthalpy
ΔHA.: Activation enthalpy
ΔHR.: Enthalpy of reaction
I: Initial state of the starting materials: metastable
II: transition state of the activated complex: unstable
III: Final state of the products: stable

Example: Coal burns with the oxygen in the air, generating heat (exothermic) to form carbon dioxide.

The enthalpy of reaction (enthalpy difference) ΔH of this reaction is negative. (See alsoStandard enthalpy of formation)

If the activation energy is very low, the reaction can be set in motion without additional external energy input. The necessary activation energy is withdrawn from the environment. The reaction takes place spontaneously.

Course of an endothermic reaction

The starting materials are initially in a stable state. By continuously supplying a certain amount of energy, the sum of activation energy and reaction energy, the system is lifted into an unstable state. If the energy supply is interrupted, the reaction also stops. In the overall balance, the chemical system absorbs energy from the environment; it is referred to as reaction energy. The products are now in a metastable state.

  Legend:
H: enthalpy
ΔHA.: Activation enthalpy
ΔHR.: Reaction energy
I: Initial state of the starting materials: stable
II: transition state of the activated complex: unstable
III: Final state of the products: metastable

An example of an endothermic reaction is the reaction of copper and sulfur to form copper sulfide. Other examples are the reaction of iron and sulfur to form monosulfide, etc. Metals and sulfur always react to form metal sulfides.

catalysis

By using catalysts, the activation enthalpy can be reduced in both endothermic and exothermic reactions. In autocatalysis, the resulting products act as catalysts for their formation.

Example: Autocatalytic formation of silver in the developer bath, see photography.):

Endergonic and exergonic reactions

If the change in entropy is included in a reaction, the Gibbs-Helmholtz equation is required for energetic considerations:

T = temperature in Kelvin
ΔG = change in free enthalpy
ΔS = change in entropy (at 298 K) from ΔS = Σ {S (products)} - Σ {S (starting materials)}
ΔH = enthalpy change (at 298 K) from ΔH = Σ {H (products)} - Σ {H (starting materials)}

In endergonic reactions, ΔG is positive, in exergonic reactions, ΔG is negative.

A chemical reaction only takes place automatically if the free enthalpy of reaction ΔG is negative.

Possible ways of interpreting the free enthalpy:

  • It is a measure of the stability of a chemical system.
  • It is a measure of the “voluntariness” or “driving force” of a reaction; with their help it can be decided whether a reaction (after possible activation) can take place without additional energy supply.
  • It is the maximum usable energy for living beings.

Statistical considerations

Some chemical reactions take place very slowly or not at all, although they are possible from a thermodynamic point of view. However, these reactions can also be accelerated by suitable reaction conditions (see reaction rate (chemistry)):

  • Increase in the degree of division and mixing and thus increase the reactive surface area in solid-state reactions. Example: Burning a lump of coal takes longer than burning the same amount of powdered coal blown into a stream of air.
  • Increase in the concentration of the reactants. Example: Carbon burns faster in pure oxygen than in the same amount of air.
  • Increase in temperature. (See RGT rule)

The kinetic gas theory of physics makes an important contribution to the modeling of the conditions of a chemical reaction at the level of the smallest particles.

  1. Educt particles are only converted into product particles if they collide with sufficiently high energy. With large molecules, the collision must also take place in the right place. (So ​​a substrate molecule has to hit the substrate binding site of its enzyme exactly. At another point of the enzyme, no reaction would be possible, even with a sufficiently high energy.)
  2. Not all particles have the same energy, so that in a reaction mixture there will also be collisions between particles with insufficient energy. There is no reaction, just an elastic shock.
  3. Whether and how quickly a reaction takes place depends on the probability with which reactant particles that have the necessary minimum energy or more will meet.

Ways to increase the likelihood of a "successful" collision:

1. Temperature increase: The higher the temperature, the more particles have the required minimum energy or more.

2. Increase in concentration: The higher the concentration of the reactants, the more likely it is that particles with the appropriate energy will meet.

Since the concentration of the starting materials decreases in the course of the reaction and the concentration of the products increases, the rate of the reaction decreases over time. The optimization of these reaction conditions is particularly important on an industrial and industrial scale.

For a detailed representation of the laws see kinetics, reaction kinetics, enzyme kinetics

Reversible reactions - equilibrium reactions

There are chemical reactions that do not take place completely. Even if the starting materials are mixed in a stoichiometric ratio, starting materials are still present at the end of the reaction.

Example: A stoichiometric mixture of light yellow iron (III) chloride and colorless potassium rhodanide solution reacts to form deep red iron rhodanide ("theater blood"). If ferric chloride is added after a while, the color will deepen. This means that unused potassium rhodanide was still available. The color can also be deepened by adding potassium rhodanide.

Hydrogen chloride reacts almost completely with water in an acid-base reaction. After the reaction has ended, there are no more hydrogen chloride molecules (H-Cl), only chloride (Cl-) and oxonium ions (H3O+) as products. Hydrogen chloride dissociates practically completely in water, its degree of dissociation is 100%. (This mixture is called hydrochloric acid designated.)

Acetic acid, on the other hand, does not completely dissociate in water, at the end of the reaction there are still 99.96% undissociated ethanoic acid molecules, the rest have reacted with water to form acetate ions.

Observations of this kind have led to the concept of reversible (reversible) chemical reactions and chemical equilibrium:

In principle, all chemical reactions are reversible, the products can react back to form starting materials. The reaction from the starting materials to the products will Forward reactionfrom the products back to the starting materials Reverse reaction called.

In the reaction equation, the reversibility of a reaction is shown by the "chemical" double arrow:

The definition of the forward reaction corresponds to the reading direction from left to right. The reverse reaction then corresponds to the reverse reading direction.

Examples:

  • Acetic acid reacts with water in a reversible reaction to form acetate and oxonium ions.
  • Reaction of ferric chloride with potassium rhodanite to form ferric rhodanite.

If the experimental conditions are not changed for reversible reactions, this will occur after some time chemical equilibrium a. The concentrations of the starting materials and products no longer change, although the back and forth reactions still continue. These unchangeable equilibrium concentrations are related to one another which is characteristic of the reaction and the reaction conditions.

For more information see under chemical equilibrium, equilibrium constant, acid constant, mass action constant, law of mass action, steady state).

In the case of reactions that appear to be complete, such as hydrogen chloride in water, the equilibrium is practically on the side of the products. The forward reaction is favored, the reverse reaction practically does not take place.

For the reaction of acetic acid with water, the equilibrium is on the side of the educts. Here the reverse reaction has a stronger influence on the equilibrium position than the forward reaction.

For educts that do not react with one another, the equilibrium is practically completely on the side of the educts.

Chemical reaction types and types

There are different types of chemical reactions:

At Additions two molecules come together to form a single one. (Opposite: decomposition)

  • A + B → AB (corresponding synthesis);
  • A + BC → ABC; AB + CD → ABCD; A + A → A2 (Dimerization)

At Condensation When two molecules come together to form a single one, they repel small molecules such as water or ammonia. (Opposite: hydrolysis)

At Polycondensations and PolyIn addition, many small molecules (monomers) react to form giant molecules (polymers). So arise z. B. from amino acids proteins.

  • According to reactions typical of the substance class or bond:
    • Production of ions and ion compounds
    • Reduction of ores to metals
    • Combustion reactions (are always redox reactions)
    • Decomposition reactions (analyzes on the whole)
    • Precipitation reactions
    • Representation reactions (syntheses in the broader sense)
    • Polymerization, polycondensation, polyaddition (giant molecules are formed)
  • According to the physical state of the substances involved
    • Gas reactions
    • Reactions in solution
    • Reactions in the melt
    • Solid reactions
    • Surface reactions
  • According to the type of reacting particles or according to reaction mechanisms

Further topics

  • Reaction scheme for symbolic representation of the reaction
  • transition state
  • Kolmogorov equation for the mathematical description of chemical reactions with diffusion and saturation terms

See also

literature

  • Eckhard Ignatowitz, Gerhard Haering: Chemistry for School and Work, p. 22ff and p. 41ff; Europe teaching aids, Haan-Gruiten; ISBN 3-8085-7054-8
  • Michael Wächter: Substances, particles, reactions. Verlag Handwerk und Technik, Hamburg 2000, pp. 154-169 ISBN 3-582-01235-2

Also teach and describe chemical reactions and their implementation, e.g. in the laboratory / in the form of detection reactions:

  • Gerhart Jander: Introduction to the inorganic-chemical internship. S. Hirzel Verlag, Stuttgart 1990 (in 13th edition), ISBN 3-7776-0477-1
  • Bertram Schmidkonz: Inorganic Analysis Internship. Harri Deutsch publishing house, Frankfurt 2002, ISBN 3-8171-1671-3

Categories: Chemical Reaction | Detection reaction