Potential of Hydrogen Content
The pH scale is logarithmic and inversely indicates the activity of hydrogen ions in the solution
![{\displaystyle {\ce {pH}}=-\log _{10}(a_{{\ce {H+}}})\thickapprox -\log _{10}([{\ce {H+}}]/{\ce {M}})}](https://wikimedia.org/api/rest_v1/media/math/render/svg/4bb26facbd1582452824ae7fa3636081a55503f0)
where [H+] is the equilibrium molar concentration of H+ (M = mol/L) in the solution.
In chemistry, hydronium (hydroxonium in traditional British English) is the cation [H3O]+, also written as H3O+, the type of oxonium ion produced by protonation of water. It is often viewed as the positive ion present when an Arrhenius acid is dissolved in water, as Arrhenius acid molecules in solution give up a proton (a positive hydrogen ion, H+) to the surrounding water molecules (H2O). In fact, acids must be surrounded by more than a single water molecule in order to ionize, yielding aqueous H+ and conjugate base.
The first modern definition of acids and bases in molecular terms was devised by Svante Arrhenius.[9][10] A hydrogen theory of acids, it followed from his 1884 work with Friedrich Wilhelm Ostwald in establishing the presence of ions in aqueous solution and led to Arrhenius receiving the Nobel Prize in Chemistry in 1903.
As defined by Arrhenius:
- An Arrhenius acid is a substance that ionises in water to form hydrogen ions (H+);[11] that is, an acid increases the concentration of H+ ions in an aqueous solution.
This causes the protonation of water, or the creation of the hydronium (H3O+) ion.[note 1] Thus, in modern times, the symbol H+ is interpreted as a shorthand for H3O+, because it is now known that a bare proton does not exist as a free species in aqueous solution.[14] This is the species which is measured by pH indicators to measure the acidity or basicity of a solution.
- An Arrhenius base is a substance that dissociates in water to form hydroxide (OH−) ions; that is, a base increases the concentration of OH− ions in an aqueous solution.
The Arrhenius definitions of acidity and alkalinity are restricted to aqueous solutions and are not valid for most non-aqueous solutions, and refer to the concentration of the solvent ions. Under this definition, pure H2SO4 and HCl dissolved in toluene are not acidic, and molten NaOH and solutions of calcium amide in liquid ammonia are not alkaline. This led to the development of the Brønsted–Lowry theory and subsequent Lewis theory to account for these non-aqueous exceptions.[15]
The reaction of an acid with a base is called a neutralization reaction. The products of this reaction are a salt and water.
In this traditional representation an acid–base neutralization reaction is formulated as a double-replacement reaction. For example, the reaction of hydrochloric acid (HCl) with sodium hydroxide (NaOH) solutions produces a solution of sodium chloride (NaCl) and some additional water molecules.
The modifier (aq) in this equation was implied by Arrhenius, rather than included explicitly. It indicates that the substances are dissolved in water. Though all three substances, HCl, NaOH and NaCl are capable of existing as pure compounds, in aqueous solutions they are fully dissociated into the aquated ions H+, Cl−, Na+ and OH−.
Three main structures for the aqueous proton have garnered experimental support: the Eigen cation, which is a tetrahydrate, H3O+(H2O)3, the Zundel cation, which is a symmetric dihydrate, H+(H2O)2, and the Stoyanov cation, an expanded Zundel cation, which is a hexahydrate: H+(H2O)2(H2O)4.[1][2] Spectroscopic evidence from well-defined IR spectra overwhelmingly supports the Stoyanov cation as the predominant form.[3][4][5][6][non-primary source needed] For this reason, it has been suggested that wherever possible, the symbol H+(aq) should be used instead of the hydronium ion.[2]
The hydron ion can incorporate an electron pair from a Lewis base into the molecule by adduction:
- [H]+
+ :L → [HL]+
Because of this capture of the Lewis base (L), the hydron ion has Lewis acidic character. In terms of Hard/Soft Acid Base (HSAB) theory, the bare hydron is an infinitely hard Lewis acid.
The hydron plays a central role in Brønsted–Lowry acid–base theory: a species that behaves as a hydron donor in a reaction is known as the Brønsted acid, while the species accepting the hydron is known as the Brønsted base. In the generic acid–base reaction shown below, HA is the acid, while B (shown with a lone pair) is the base:
- HA + :B → [HB]+
+ :A–
The hydrated form of the hydrogen cation, the hydronium (hydroxonium) ion H
3O+
(aq), is a key object of Arrhenius' definition of acid. Other hydrated forms, the Zundel cation H
5O+
2, which is formed from a proton and two water molecules, and the Eigen cation H
9O+
4, which is formed from a hydronium ion and three water molecules, are theorized to play an important role in the diffusion of protons though an aqueous solution according to the Grotthuss mechanism. Although the ion H
3O+
(aq) is often shown in introductory textbooks to emphasize that the hydron is never present as an unsolvated species in aqueous solution, it is somewhat misleading, as it oversimplifies infamously complex speciation of the solvated proton in water; the notation H+
(aq) is often preferred, since it conveys aqueous solvation while remaining noncommittal with respect to the number of water molecules involved.
- Proton, having the symbol p or 1H+, is the +1 ion of protium, 1H.
- Deuteron, having the symbol 2H+ or D+, is the +1 ion of deuterium, 2H or D.
- Triton, having the symbol 3H+ or T+, is the +1 ion of tritium, 3H or T.
Other isotopes of hydrogen are too unstable to be relevant in chemistry.
The molar concentration of hydronium or H+ ions determines a solution's pH according to
- pH = -log([H3O+]/M)
where M = mol/L. The concentration of hydroxide ions analogously determines a solution's pOH. The molecules in pure water auto-dissociate into aqueous protons and hydroxide ions in the following equilibrium:
- H2O ⇌ OH−(aq) + H+(aq)
In pure water, there is an equal number of hydroxide and H+ ions, so it is a neutral solution. At 25 °C (77 °F), pure water has a pH of 7 and a pOH of 7 (this varies when the temperature changes: see self-ionization of water). A pH value less than 7 indicates an acidic solution, and a pH value more than 7 indicates a basic solution.[7]
At 25 °C (77°F), solutions with a pH less than 7 are acidic, and solutions with a pH greater than 7 are basic. Solutions with a pH of 7 at 25 °C are neutral (i.e. have the same concentration of H+ ions as OH− ions, i.e. the same as pure water). The neutral value of the pH depends on the temperature and is lower than 7 if the temperature increases above 25 °C. The pH range is commonly given as zero to 14, but a pH value can be less than 0 for very concentrated strong acids or greater than 14 for very concentrated strong bases.[2]
The pH scale is traceable to a set of standard solutions whose pH is established by international agreement.[3] Primary pH standard values are determined using a concentration cell with transference by measuring the potential difference between a hydrogen electrode and a standard electrode such as the silver chloride electrode. The pH of aqueous solutions can be measured with a glass electrode and a pH meter or a color-changing indicator. Measurements of pH are important in chemistry, agronomy, medicine, water treatment, and many other applications.
An acid is a molecule or ion capable of either donating a proton (i.e. hydrogen ion, H+), known as a Brønsted–Lowry acid, or forming a covalent bond with an electron pair, known as a Lewis acid.[1]
The first category of acids are the proton donors, or Brønsted–Lowry acids. In the special case of aqueous solutions, proton donors form the hydronium ion H3O+ and are known as Arrhenius acids. Brønsted and Lowry generalized the Arrhenius theory to include non-aqueous solvents. A Brønsted or Arrhenius acid usually contains a hydrogen atom bonded to a chemical structure that is still energetically favorable after loss of H+.
Aqueous Arrhenius acids have characteristic properties that provide a practical description of an acid.[2] Acids form aqueous solutions with a sour taste, can turn blue litmus red, and react with bases and certain metals (like calcium) to form salts. The word acid is derived from the Latin acidus, meaning 'sour'.[3] An aqueous solution of an acid has a pH less than 7 and is colloquially also referred to as "acid" (as in "dissolved in acid"), while the strict definition refers only to the solute.[1] 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 include hydrochloric acid (a solution 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[1] sense) that are solids, liquids, or gases. Strong acids and some concentrated weak acids are corrosive, but there are exceptions such as carboranes and boric acid.
In chemistry, there are three definitions in common use of the word "base": Arrhenius bases, Brønsted bases, and Lewis bases. All definitions agree that bases are substances that react with acids, as originally proposed by G.-F. Rouelle in the mid-18th century.
In 1884, Svante Arrhenius proposed that a base is a substance which dissociates in aqueous solution to form hydroxide ions OH−. These ions can react with hydrogen ions (H+ according to Arrhenius) from the dissociation of acids to form water in an acid–base reaction. A base was therefore a metal hydroxide such as NaOH or Ca(OH)2. Such aqueous hydroxide solutions were also described by certain characteristic properties. They are slippery to the touch, can taste bitter[1] and change the color of pH indicators (e.g., turn red litmus paper blue).
In water, by altering the autoionization equilibrium, bases yield solutions in which the hydrogen ion activity is lower than it is in pure water, i.e., the water has a pH higher than 7.0 at standard conditions. A soluble base is called an alkali if it contains and releases OH− ions quantitatively. Metal oxides, hydroxides, and especially alkoxides are basic, and conjugate bases of weak acids are weak bases.
Bases and acids are seen as chemical opposites because the effect of an acid is to increase the hydronium (H3O+) concentration in water, whereas bases reduce this concentration. A reaction between aqueous solutions of an acid and a base is called neutralization, producing a solution of water and a salt in which the salt separates into its component ions. If the aqueous solution is saturated with a given salt solute, any additional such salt precipitates out of the solution.
A strong base is a basic chemical compound that can remove a proton (H+) from (or deprotonate) a molecule of even a very weak acid (such as water) in an acid–base reaction. Common examples of strong bases include hydroxides of alkali metals and alkaline earth metals, like NaOH and Ca(OH)
2, respectively. Due to their low solubility, some bases, such as alkaline earth hydroxides, can be used when the solubility factor is not taken into account.[8]
One advantage of this low solubility is that "many antacids were suspensions of metal hydroxides such as aluminium hydroxide and magnesium hydroxide";[9] compounds with low solubility and the ability to stop an increase in the concentration of the hydroxide ion, preventing the harm of the tissues in the mouth, oesophagus, and stomach.[9] As the reaction continues and the salts dissolve, the stomach acid reacts with the hydroxide produced by the suspensions.[9]
Strong bases hydrolyze in water almost completely, resulting in the leveling effect."[7] In this process, the water molecule combines with a strong base, due to the water's amphoteric ability; and, a hydroxide ion is released.[7] Very strong bases can even deprotonate very weakly acidic C–H groups in the absence of water. Here is a list of several strong bases:
| Lithium hydroxide | LiOH |
| Sodium hydroxide | NaOH |
| Potassium hydroxide | KOH |
| Rubidium hydroxide | RbOH |
| Cesium hydroxide | CsOH |
| Magnesium hydroxide | Mg(OH) 2 |
| Calcium hydroxide | Ca(OH) 2 |
| Strontium hydroxide | Sr(OH) 2 |
| Barium hydroxide | Ba(OH) 2 |
| Tetramethylammonium hydroxide | N(CH 3) 4OH |
| Guanidine | HNC(NH 2) 2 |
The cations of these strong bases appear in the first and second groups of the periodic table (alkali and earth alkali metals). Tetraalkylated ammonium hydroxides are also strong bases since they dissociate completely in water. Guanidine is a special case of a species that is exceptionally stable when protonated, analogously to the reason that makes perchloric acid and sulfuric acid very strong acids.
In chemistry, an acid–base reaction is a chemical reaction that occurs between an acid and a base. It can be used to determine pH via titration. Several theoretical frameworks provide alternative conceptions of the reaction mechanisms and their application in solving related problems; these are called the acid–base theories, for example, Brønsted–Lowry acid–base theory.
Their importance becomes apparent in analyzing acid–base reactions for gaseous or liquid species, or when acid or base character may be somewhat less apparent. The first of these concepts was provided by the French chemist Antoine Lavoisier, around 1776.[1]
It is important to think of the acid–base reaction models as theories that complement each other.[2] For example, the current Lewis model has the broadest definition of what an acid and base are, with the Brønsted–Lowry theory being a subset of what acids and bases are, and the Arrhenius theory being the most restrictive.
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