Chemical reaction

• 1. Introduction

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  A chemical reaction begins when atoms rearrange themselves. One set of chemical substances transforms into another, and an energy change rides along with it. New products appear, often with properties nothing like the materials that made them. Hazardous sodium metal meets toxic chlorine gas, and the result is common table salt. That single fact captures something strange about this subject. The dangerous becomes the everyday, all through the shifting of electrons.

  Classically, these changes touch only the positions of electrons. Bonds break and form between atoms, while the nuclei stay untouched and the elements present do not change. A chemical equation can often describe the whole event, with an arrow read as the word yields. But the questions pile up fast. Why do some reactions release heat while others swallow it? Why does raising the temperature usually make a reaction race faster? What lets a single substance speed a reaction without ever being consumed? And how did humanity learn to coax these transformations, from alchemists chasing gold to factories making acid by the ton? This is the story of how matter changes, and of the rules that govern when, why, and how fast.

• 2. From Alchemy To Industry

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  Combustion in fire, fermentation, and the reduction of ores to metals were known since antiquity. The Greek philosopher Empedocles offered an early theory with his Four-Element Theory, holding that any substance is composed of fire, water, air, and earth. In the Middle Ages, alchemists took up the study of chemical transformations. They tried, in particular, to convert lead into gold, using reactions of lead and lead-copper alloys with sulfur.

  The works attributed to Jābir ibn Ḥayyān, dated to roughly 850 to 950, describe the synthesis of ammonium chloride from organic substances. Later alchemists, starting around 1300, produced mineral acids such as sulfuric and nitric acids. They heated sulfate and nitrate minerals, including copper sulfate, alum, and saltpeter. In the 17th century, Johann Rudolph Glauber made hydrochloric acid and sodium sulfate by reacting sulfuric acid with sodium chloride.

  The lead chamber process, introduced in 1746, and the Leblanc process brought large-scale production of sulfuric acid and sodium carbonate. With them, chemical reactions entered industry. Sulfuric acid technology improved further into the contact process in the 1880s. The Haber process followed, developed in 1909 to 1910 for the synthesis of ammonia, carrying these old transformations to the scale of the modern factory.

• 3. Theories Of Transformation

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  From the 16th century onward, Jan Baptist van Helmont, Robert Boyle, and Isaac Newton tried to build theories for the chemical changes they observed in experiments. In 1667, Johann Joachim Becher proposed the phlogiston theory. It claimed a fire-like element called phlogiston lived inside combustible bodies and escaped during combustion. The idea held until 1785, when Antoine Lavoisier proved it false. He found that combustion is a reaction with oxygen from the air.

  Joseph Louis Gay-Lussac recognized in 1808 that gases always react in fixed relationships with one another. That insight joined the atomic theory of John Dalton, and from it Joseph Proust developed the law of definite proportions. These ideas later gave rise to stoichiometry and to chemical equations themselves.

  Compounds drawn from living organisms were long thought too complex to make synthetically. The concept of vitalism held that organic matter carried a vital force setting it apart from inorganic materials. Friedrich Wöhler ended that divide in 1828 by synthesizing urea from inorganic precursors. Alexander William Williamson contributed his synthesis of ethers, and Christopher Kelk Ingold established the mechanisms of substitution reactions, among many other discoveries.

• 4. Heat, Order, And Direction

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  The laws of thermodynamics decide whether a reaction will proceed. A reaction can run by itself if it is exergonic, releasing free energy. That free energy change combines two quantities, enthalpy and entropy, linked through temperature. Exothermic reactions like combustion, precipitation, and crystallization release energy and form ordered solids from disordered gases or liquids. Endothermic reactions instead pull heat from their surroundings, often by raising entropy through gaseous or dissolved products.

  Temperature can flip the very sign of a reaction's enthalpy. In the carbon monoxide reduction of molybdenum dioxide, the reaction to form carbon dioxide and molybdenum is endothermic at low temperatures. Its enthalpy reaches zero at 1855 and becomes exothermic above that point. Temperature can also reverse a reaction's preferred direction. The water gas shift reaction is favored at low temperatures, while its reverse is favored at high ones, with the crossover at 1100.

  Most chemical reactions are reversible and run in both directions at once. The forward and reverse reactions compete at rates that shift over time. The reverse rate climbs until it matches the forward rate, reaching chemical equilibrium. At equilibrium, the Gibbs free energy of reaction must be zero. Le Chatelier's principle explains pressure effects: an increase in pressure from shrinking volume pushes the reaction toward the side with fewer moles of gas.

• 5. The Pace Of Change

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  Reaction kinetics studies the speed at which reactions take place. Raising reactant concentrations usually quickens a reaction through more collisions per unit of time. Some reactions, called zero order, ignore concentration entirely because catalytic sites are limited. Larger surface area between reactants, especially solids in heterogeneous systems, drives rates higher. Increased pressure shrinks the space between molecules and raises collision frequency.

  Activation energy is the amount of energy needed to make a reaction start and carry on spontaneously. Higher activation energy means reactants need more energy to begin. Rising temperature hastens reactions by adding energy to the molecules, creating more collisions. A catalyst lowers the activation energy by changing a reaction's pathway, and it emerges unchanged, ready to be used again. For some reactions, ultraviolet light is needed to break bonds and start the process, especially where radicals are involved.

  The rate of a first-order reaction depends only on concentration and the substance's properties, and can be described by a characteristic half-life. The temperature dependence of the rate constant usually follows the Arrhenius equation, built around the activation energy and the Boltzmann constant. Collision theory is one of the simplest models. More realistic ones include transition state theory, the potential energy surface, Marcus theory, and the Rice-Ramsperger-Kassel-Marcus theory.

• 6. The Family Of Reactions

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  Four basic types organize the simplest reactions. In a synthesis reaction, two or more simple substances combine into something more complex, as when iron and sulfur form iron(II) sulfide. A decomposition reaction does the reverse, breaking a complex substance into simpler parts, as in the electrolysis of water into oxygen and hydrogen gas. A single displacement reaction sends one element to trade places with another in a compound, as when magnesium replaces hydrogen in water to make magnesium hydroxide and hydrogen gas. In a double displacement reaction, the anions and cations of two compounds switch partners, as when lead(II) nitrate reacts with potassium iodide to yield lead(II) iodide and potassium nitrate.

  Redox reactions turn on the transfer of electrons from a reducing agent to an oxidizing agent. The reducing agent is oxidized and the oxidizing agent is reduced. When sodium metal reacts with chlorine gas, sodium moves from oxidation state 0 to +1, losing one electron, while chlorine moves from 0 to -1, gaining one. Elements with low electronegativities, like most metals, donate electrons easily and act as reducing agents. Electrolytic electrochemical reactions are an important class, used to produce elements such as chlorine and aluminium, and their reverse powers batteries.

  Combustion reactions pit an element or compound against an oxidant, usually oxygen, often releasing heat or light. They frequently involve a hydrocarbon. The combustion of 1 mole, or 114 grams, of octane in oxygen releases 5500 kilojoules. Carbon, magnesium, or sulfur reacting with oxygen also counts as combustion.

• 7. Bonds Made And Broken

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  Organic chemistry adds reactions built around covalent bonds between carbon atoms or between carbon and heteroatoms like oxygen, nitrogen, and halogens. Many of these are name reactions, called after their discoverers. One of the most industrially important is the cracking of heavy hydrocarbons at oil refineries into smaller molecules, the process used to manufacture gasoline.

  In a substitution reaction, one functional group replaces another, and the substituting species may be nucleophilic, electrophilic, or radical. Nucleophilic substitution runs through two mechanisms, SN1 and SN2, where S stands for substitution, N for nucleophilic, and the number marks the kinetic order. The SN1 reaction proceeds in two steps through a carbocation and yields geometric isomers. The SN2 mechanism produces a Walden inversion, reversing the existing stereochemistry. Electrophilic substitution takes place almost exclusively in aromatic hydrocarbons, where the electrophile's attack abolishes the aromatic system before a proton splits off and restores it.

  Addition and elimination change the number of substituents on a carbon atom and form or cleave multiple bonds. The E1 mechanism ejects the leaving group first to form a carbocation, while the E2 mechanism has the base attack and the leaving group depart at once. Markovnikov's rule predicts where a carbocation forms in electrophilic addition. The Michael reaction, a conjugate addition, ranks among the most useful methods for the mild formation of carbon-carbon bonds. The Diels-Alder reaction, a 4+2 cycloaddition between a conjugated diene and a substituted alkene, builds a substituted cyclohexene system, with its outcomes governed by the Woodward-Hoffmann rules.

• 8. Reactions Of Life And Industry

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  Enzymes, complex proteins usually specialized to catalyze a single specific reaction, control most biochemical reactions. The reaction happens in the active site, a small cleft or pocket lined by amino acid residues, while the rest of the enzyme provides stabilization. The biochemical reactions in living organisms are known collectively as metabolism. Anabolism builds large molecules like proteins and carbohydrates from smaller units through DNA and enzyme-controlled processes. Glucose and oxygen are important energy sources, and all organisms use their energy to make adenosine triphosphate, the molecule that powers other reactions.

  Photochemistry runs many of these processes. Plants use solar energy in photosynthesis to convert carbon dioxide and water into glucose, releasing oxygen as a side-product. Humans rely on photochemistry to form vitamin D, and vision begins with a photochemical reaction of rhodopsin. In fireflies, an enzyme in the abdomen catalyzes a reaction that produces bioluminescence. Ozone formation and other atmospheric chemistry occur high in the Earth atmosphere.

  Chemical reactions sit at the heart of chemical engineering, synthesizing new compounds from petroleum, mineral ores, and oxygen in air. The goal is efficiency: maximizing yield while minimizing reagents, energy, and waste. The thermite reaction holds a narrower niche, generating light and heat for pyrotechnics and welding. Though less controllable than oxy-fuel, arc, or flash welding, it needs far less equipment and still mends rails in remote areas. Watching these reactions unfold depends on their speed. Slow ones can be tracked by pH, color, and emission spectra, while the fastest are caught by femtosecond lasers, freezing short-lived transition states down to a few femtoseconds.