Acid–base reaction

• 1. Introduction

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  An acid-base reaction is one of chemistry's most familiar events, yet no single rule has ever fully captured it. When hydrochloric acid meets sodium hydroxide in water, the two dissolve into a solution of sodium chloride and a little extra water. That tidy result looks like the whole story. It is not. The French chemist Antoine Lavoisier offered the first scientific concept of acids and bases around 1776, and ever since, chemists have kept rewriting what an acid and a base even are. Why would a settled idea need redefining again and again? How did a theory built on oxygen give way to one built on hydrogen, and then to ones that need no water at all? And how can the same substance be a strong acid in one liquid and a weak base in another? The answers live in a chain of competing definitions, each broader than the last.

• 2. Naming The Base

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  In 1754, Guillaume-Francois Rouelle introduced the word "base" into chemistry. He used it to mean a substance that reacts with an acid to give it solid form, as a salt. Bases, he noted, are mostly bitter in nature. That single naming choice gave the field a vocabulary it still uses, long before anyone understood what was happening at the level of ions. The word arrived decades ahead of the science meant to explain it.

• 3. The Oxygen Mistake

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  Lavoisier believed acids owed their character to oxygen. His knowledge of strong acids was mostly limited to oxoacids like nitric acid and sulfuric acid, which carry central atoms in high oxidation states surrounded by oxygen. Not knowing the true makeup of the hydrohalic acids HF, HCl, HBr, and HI, he defined acids by their oxygen content. He even coined the element's name from Greek words meaning "acid-former". That definition held for over 30 years. Sir Humphry Davy broke it in an 1810 article and later lectures, proving that hydrogen sulfide, hydrogen telluride, and the hydrohalic acids contain no oxygen. Davy stopped short of a replacement theory, concluding that "acidity does not depend upon any particular elementary substance, but upon peculiar arrangement of various substances". Jons Jacob Berzelius offered a notable patch, casting acids as oxides of nonmetals and bases as oxides of metals. The hydrogen-based future Davy had hinted at was still waiting for someone to build it.

• 4. Hydrogen Takes Over

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  In 1838, Justus von Liebig proposed that an acid is a hydrogen-containing compound whose hydrogen can be replaced by a metal. His extensive study of the chemical composition of organic acids carried the shift Davy had started from oxygen to hydrogen. Liebig's definition was completely empirical, yet it stayed in use for almost 50 years. The next leap came from Svante Arrhenius, whose 1884 work with Friedrich Wilhelm Ostwald established the presence of ions in aqueous solution. That work earned Arrhenius the Nobel Prize in Chemistry in 1903. He defined an acid as a substance that ionises in water to form hydrogen cations, raising the concentration of H+ ions in solution. A base, in turn, dissociates in water to form hydroxide ions. The protonation of water creates the hydronium ion, since a bare proton does not exist as a free species in aqueous solution. This is the species that pH indicators measure. Arrhenius bought clarity at a cost: his definitions hold only in water, leaving non-aqueous chemistry unexplained.

• 5. Salt And A Rising Loaf

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  The reaction of an acid with a base is called a neutralization reaction, and its products are a salt and water. Chemists traditionally write it as a double-replacement reaction, as when hydrochloric acid and sodium hydroxide yield sodium chloride and extra water. Arrhenius implied the "aq" modifier rather than writing it out, marking that HCl, NaOH, and NaCl, though able to exist as pure compounds, fully dissociate into aquated ions in water. Baking powder turns this chemistry into breakfast. It makes dough rise by creating millions of tiny carbon dioxide bubbles. It should not be confused with baking soda, which is sodium bicarbonate; baking powder mixes that bicarbonate with acidic salts. When water joins them, the bicarbonate and acid salts react to release gaseous carbon dioxide. A typical formulation by weight might call for 30% sodium bicarbonate, 5-12% monocalcium phosphate, and 21-26% sodium aluminium sulfate. A commercial blend might swap in sodium acid pyrophosphate instead, or use cream of tartar, a derivative of tartaric acid.

• 6. Donors And Acceptors

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  In 1923, Johannes Nicolaus Bronsted in Denmark and Martin Lowry in England independently arrived at the same idea. An acid donates a hydrogen cation, a proton, and a base accepts it. Removing a proton from an acid produces its conjugate base; adding one to a base produces its conjugate acid. Unlike earlier views, this model never mentions salt or solvent, and the concept of neutralization is absent. Acids and bases simply transfer a proton to form a new acid and a new base. Bronsted-Lowry behavior is formally independent of any solvent, which is its great expansion. Where Arrhenius needed alkalis dissolving in water, this model let pH be tested across gas, liquid, and solid. Water itself proves the point, being amphoteric: one water molecule can donate a proton while another accepts it. When acetic acid dissolves in liquid ammonia, it forms the acetate ion while ammonia becomes the ammonium ion, no water required. The same year, Gilbert N. Lewis was preparing a definition that would drop the hydrogen requirement entirely.

• 7. The Electron Pair

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  Gilbert N. Lewis devised his definition in 1923, the same year as Bronsted-Lowry, but did not elaborate it until 1938. He recast the relationship around electrons: a Lewis base donates an electron pair, and a Lewis acid receives one. Boron trifluoride is a typical Lewis acid, accepting a pair of electrons into a vacancy in its octet, while a fluoride ion with its full octet supplies the pair. All group 13 compounds of that form can behave as Lewis acids, while group 15 compounds such as amines and phosphines act as Lewis bases. Adducts join an acceptor and a donor through a dative covalent bond. Even carbon monoxide acts as a Lewis base when it bonds with boron trifluoride. Lewis had warned against the narrower view in 1938, writing, "To restrict the group of acids to those substances that contain hydrogen interferes as seriously with the systematic understanding of chemistry as would the restriction of the term oxidizing agent to substances containing oxygen." His framework treats reactions where a stronger base like ammonia replaces a weaker one like water, and it stays consistent with Bronsted-Lowry. Yet some chemists still felt there was something intrinsically acidic about hydrogen compounds that no electron-pair rule could capture.

• 8. Beyond Water And Hydrogen

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  Edward Curtis Franklin studied acid-base reactions in liquid ammonia in 1905, noting their likeness to the water-based Arrhenius theory. Working with liquid phosgene, Albert F.O. Germann formulated the solvent-based theory in 1925, extending the idea to aprotic solvents. He described solvonium ions, the positive species, and solvate ions, the negative ones; a solute that raises solvonium concentration is an acid, and one that raises solvate concentration is a base. Under this view, the same substance can flip roles, acting as a strong acid in water, a weak acid in acetic acid, and a weak base in fluorosulfonic acid. Critics called the theory too general to be useful. Other definitions pushed in other directions. Hermann Lux revived the oxygen theory in 1939, improved by Hakon Flood around 1947, defining an acid as an oxide ion acceptor and a base as an oxide ion donor; it still serves modern geochemistry and the electrochemistry of molten salts, and helps systematize the xenon oxides, fluorides, and oxofluorides. Mikhail Usanovich, publishing in 1938, went even broader than Lewis, casting acids as anything that accepts negative species or donates positive ones, which folded redox into acid-base chemistry. To rank these interactions, Ralph Pearson proposed the Hard and Soft Acids and Bases principle in 1963, made quantitative with Robert Parr in 1984, holding that hard-hard and soft-soft pairings are most stable. Russell S. Drago's ECW model went further still, assigning E and C parameters to many acids and bases and showing there is no single order of Lewis strengths at all.