Science, Tech, Math › Science What Are Acids and Bases? Share Flipboard Email Print Vikram Raghuvanshi / Getty Images Science Chemistry Basics Chemical Laws Molecules Periodic Table Projects & Experiments Scientific Method Biochemistry Physical Chemistry Medical Chemistry Chemistry In Everyday Life Famous Chemists Activities for Kids Abbreviations & Acronyms Biology Physics Geology Astronomy Weather & Climate By Anne Marie Helmenstine, Ph.D. Chemistry Expert Ph.D., Biomedical Sciences, University of Tennessee at Knoxville B.A., Physics and Mathematics, Hastings College Dr. Helmenstine holds a Ph.D. in biomedical sciences and is a science writer, educator, and consultant. She has taught science courses at the high school, college, and graduate levels. our editorial process Facebook Facebook Twitter Twitter Anne Marie Helmenstine, Ph.D. Updated May 05, 2019 There are several methods of defining acids and bases. While these definitions don't contradict each other, they do vary in how inclusive they are. The most common definitions of acids and bases are Arrhenius acids and bases, Brønsted-Lowry acids and bases, and Lewis acids and bases. Antoine Lavoisier, Humphry Davy, and Justus Liebig also made observations regarding acids and bases, but didn't formalize definitions. Svante Arrhenius Acids and Bases The Arrhenius theory of acids and bases dates back to 1884, building on his observation that salts, such as sodium chloride, dissociate into what he termed ions when placed into water. acids produce H+ ions in aqueous solutionsbases produce OH- ions in aqueous solutionswater required, so only allows for aqueous solutionsonly protic acids are allowed; required to produce hydrogen ionsonly hydroxide bases are allowed Johannes Nicolaus Brønsted - Thomas Martin Lowry Acids and Bases The Brønsted or Brønsted-Lowry theory describes acid-base reactions as an acid releasing a proton and a base accepting a proton. While the acid definition is pretty much the same as that proposed by Arrhenius (a hydrogen ion is a proton), the definition of what constitutes a base is much broader. acids are proton donorsbases are proton acceptorsaqueous solutions are permissiblebases besides hydroxides are permissibleonly protic acids are allowed Gilbert Newton Lewis Acids and Bases The Lewis theory of acids and bases is the least restrictive model. It doesn't deal with protons at all, but deals exclusively with electron pairs. acids are electron pair acceptorsbases are electron pair donorsleast restrictive of the acid-base definitions Properties of Acids and Bases Robert Boyle described the qualities of acids and bases in 1661. These characteristics may be used to easily distinguish between the two sets up chemicals without performing complicated tests: Acids taste sour (don't taste them!)—the word 'acid' comes from the Latin acere, which means 'sour'acids are corrosiveacids change litmus (a blue vegetable dye) from blue to redtheir aqueous (water) solutions conduct electric current (are electrolytes)react with bases to form salts and waterevolve hydrogen gas (H2) upon reaction with an active metal (such as alkali metals, alkaline earth metals, zinc, aluminum) Common Acids citric acid (from certain fruits and veggies, notably citrus fruits)ascorbic acid (vitamin C, as from certain fruits)vinegar (5% acetic acid)carbonic acid (for carbonation of soft drinks)lactic acid (in buttermilk) Bases taste bitter (don't taste them!)feel slippery or soapy (don't arbitrarily touch them!)bases don't change the color of litmus; they can turn red (acidified) litmus back to bluetheir aqueous (water) solutions conduct an electric current (are electrolytes)react with acids to form salts and water Common Bases detergentssoaplye (NaOH)household ammonia (aqueous) Strong and Weak Acids and Bases The strength of acids and bases depends on their ability to dissociate or break into their ions in water. A strong acid or strong base completely dissociates (e.g., HCl or NaOH), while a weak acid or weak base only partially dissociates (e.g., acetic acid). The acid dissociation constant and base dissociation constant indicates the relative strength of an acid or base. The acid dissociation constant Ka is the equilibrium constant of an acid-base dissociation: HA + H2O ⇆ A- + H3O+ where HA is the acid and A- is the conjugate base. Ka = [A-][H3O+] / [HA][H2O] This is used to calculate pKa, the logarithmic constant: pka = - log10 Ka The larger the pKa value, the smaller the dissociation of the acid and the weaker the acid. Strong acids have a pKa of less than -2.