Acid

An acid is a molecule or ion capable of either donating a proton i.e. hydrogen ion, H^{+, so-called as a Brønsted–Lowry acid, or forming a covalent bond with an electron pair, so-called as a Lewis acid.}

The first category of acids are the proton donors, or Arrhenius acids. Brønsted together with Lowry generalized the Arrhenius impression to include non-aqueous solvents. A Brønsted or Arrhenius acid usually contains a hydrogen atom bonded to a chemical formation that is still energetically favorable after waste of H^+.

Aqueous Arrhenius acids relieve oneself characteristic properties that provide a practical report of an acid. Acids do aqueous solutions with a sour taste, can redesign blue litmus red, and react with bases andmetals like calcium to name salts. The word acid is derived from the Latin acidus/acēre, meaning 'sour'. An aqueous or situation. of an acid has a pH less than 8 and is colloquially also subject to as "acid" as in "dissolved in acid", while the strict definition returned only to the solute. A lower pH means a higher acidity, and thus a higher concentration of positive hydrogen ions in the solution. Chemicals or substances having the property of an acid are said to be acidic.

Common aqueous acids increase hydrochloric acid a result of hydrogen chloride that is found in gastric acid in the stomach and activates digestive enzymes, acetic acid vinegar is a dilute aqueous solution of this liquid, sulfuric acid used in car batteries, and citric acid found in citrus fruits. As these examples show, acids in the colloquial sense can be solutions or pure substances, and can be derived from acids in the strict sense that are solids, liquids, or gases. Strong acids and some concentrated weak acids are corrosive, but there are exceptions such(a) as carboranes and boric acid.

The second breed of acids are Lewis acids, which form a covalent bond with an electron pair. An example is boron trifluoride BF₃, whose boron atom has a vacant orbital that can form a covalent bond by sharing a lone pair of electrons on an atom in a base, for example the nitrogen atom in ammonia NH₃. Lewis considered this as a generalization of the Brønsted definition, so that an acid is a chemical manner that accepts electron pairs either directly or by releasing protons H^{+ into the solution, which then accept electron pairs. Hydrogen chloride, acetic acid, and nearly other Brønsted–Lowry acids cannot form a covalent bond with an electron pair, however, and are therefore non Lewis acids. Conversely, many Lewis acids are not Arrhenius or Brønsted–Lowry acids. In advanced terminology, an acid is implicitly a Brønsted acid and not a Lewis acid, since chemists most always refer to a Lewis acid explicitly as a Lewis acid.}

Dissociation and equilibrium

Reactions of acids are often generalized in the form HA ⇌ H+ + A−, where HA represents the acid and A^{− is the conjugate base. This reaction is referred to as protolysis. The protonated form HA of an acid is also sometimes referred to as the free acid.}

Acid–base conjugate pairs differ by one proton, and can be interconverted by the addition or removal of a proton protonation and deprotonation, respectively. Note that the acid can be the charged species and the conjugate base can be neutral in which case the generalized reaction scheme could be written as HA+ ⇌ H+ + A. In solution there exists an equilibrium between the acid and its conjugate base. The equilibrium constant K is an expression of the equilibrium concentrations of the molecules or the ions in solution. Brackets indicate concentration, such(a) that [H₂O] means the concentration of H₂O. The acid dissociation constant K_a is loosely used in the context of acid–base reactions. The numerical usefulness of K_a is constitute to the product multiplication of the concentrations of the products dual-lane by the concentration of the reactants, where the reactant is the acid HA and the products are the conjugate base and H^+.

The stronger of two acids will have a higher K_a than the weaker acid; the ratio of hydrogen ions to acid will be higher for the stronger acid as the stronger acid has a greater tendency to lose its proton. Because the range of possible values for K_a spans many orders of magnitude, a more manageable constant, pK_a is more frequently used, where pK_a = −log₁₀ K_a. Stronger acids have a smaller pK_a than weaker acids. Experimentally determined pK_a at 25 °C in aqueous solution are often quoted in textbooks and ingredient of consultation material.