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In [[chemistry]], a '''base''' is most commonly thought of as a substance that can accept [[proton]]s. This refers to the [[Acid-base reaction theories#The protonic (Brønsted-Lowry) definition|Brønsted-Lowry theory]] of acids and bases. Alternate definitions of bases include [[electron pair]] donors (Lewis), as sources of hydroxide anions ([[Arrhenius]]) and can be (commonly) thought of as any chemical compound that, when dissolved in water, gives a solution with a pH higher than 7.0.
Examples of simple bases are [[sodium hydroxide]] and [[ammonia]].

Bases can be thought of as the chemical opposite of [[acid]]s. A reaction between an acid and base is called [[neutralization]]. Bases and [[acid]]s are seen as opposites because the effect of an acid is to increase the [[hydronium]] ion (H<sub>3</sub>O<sup>+</sup>) [[concentration]] in water, whereas bases reduce this concentration. Bases react with acids to produce [[water]] and [[salts]] (or their solutions). 

==Definitions==
{{main|acid-base reaction theories}}
A [[strong]] base is a base which [[hydrolysis|hydrolyzes]] completely, raising the [[pH]] of the solution towards 14. Strong bases, like strong acids, attack living tissue and cause serious burns. They react differently to skin than acids do, so while strong acids are corrosive, we say that strong bases are [[Causticity|caustic]]. [[Superbase]]s are a class of especially basic compounds and [[non-nucleophilic base]]s are a special class of strong bases with poor nucleophilicity. Bases may also be [[weak base]]s such as [[ammonia]], which is used for cleaning. '''Arrhenius bases''' are water-soluble and these solutions always have a [[pH]] greater than 7. An [[alkali]] is a special example of a base, where in an aqueous environment, [[hydroxide|hydroxide ions]] (also viewed as OH<sup>-</sup>) are donated. There are other more generalized and advanced [[Acid-base reaction theories|definitions of acids and bases]].

The notion of a base as a concept in chemistry was first introduced by the French chemist [[Guillaume François Rouelle]] in 1754. He noted that acids which in those days were mostly volatile liquids (like [[acetic acid]]) turned into solid salts only when combined with specific substances. These substances form a concrete ''base'' for the salt <ref>''The Origin of the Term Base''  William B. Jensen Journal of Chemical Education • 1130 Vol. 83 No. 8 August '''2006'''</ref> and hence the name.

== Properties ==
Some general properties of bases include:
* [[Bitter taste]] (opposed to [[sour|sour taste]] of [[acid]]s and [[sweetness]] of [[aldehydes]] and [[ketones]])
* Slimy or soapy feel on fingers, due to [[saponification]] of the lipids in human skin
* Concentrated or strong bases are [[Corrosive|caustic]] (corrosive) on organic matter and react violently with acidic substances
* Aqueous solutions or molten bases dissociate in ions and conduct electricity
* Reactions with indicators: bases turn red [[litmus paper]] blue and [[phenolphthalein]] red

== Bases and pH ==
The [[pH]] of (impure) water is a measure of its [[acidity]]. In pure water, about one in ten million molecules dissociate into [[hydronium ion]]s (H<sub>3</sub>O<sup>+</sup>) and [[hydroxide|hydroxide ions]] (OH<sup>&minus;</sup>), according to the following equation:

:2H<sub>2</sub>O(l) → H<sub>3</sub>O<sup>+</sup>(aq) + OH<sup>-</sup>(aq)

The [[concentration]], measured in [[molarity]] (''M'' or [[Mole (unit)|mole]]s per dm³), of the ions is indicated as [H<sub>3</sub>O<sup>+</sup>] and [OH<sup>&minus;</sup>]; their product is the [[dissociation constant]] of water with and has the value 10<sup>&minus;7</sup> ''M''. The pH is defined as &minus;log&nbsp;[H<sub>3</sub>O<sup>+</sup>]; thus, pure water has a pH of 7. (These numbers are correct at 23&nbsp;°C and slightly different at other temperatures.)

A base accepts (removes) [[hydronium ion]]s (H<sub>3</sub>O<sup>+</sup>) from the solution, or donates [[hydroxide|hydroxide ions]] (OH<sup>-</sup>) to the solution. Both actions will lower the concentration of hydronium ions, and thus raise pH. By contrast, an acid donates H<sub>3</sub>O<sup>+</sup> ions to the solution or accepts OH<sup>&minus;</sup>, thus lowering pH.

For example, if 1 mole of [[sodium hydroxide]] (40 [[g]]) is dissolved in water to make 1 litre of solution, the concentration of hydroxide ions becomes [OH<sup>&minus;</sup>]&nbsp;=&nbsp;1&nbsp;mol/L. Therefore [H<sup>+</sup>]&nbsp;=&nbsp;10<sup>&minus;14</sup>&nbsp;mol/L, and pH&nbsp;=&nbsp;&minus;log&nbsp;10<sup>&minus;14</sup>&nbsp;=&nbsp;14. Note that in this calculation, it is assumed that the [[Activity (chemistry)|activity]] is equivalent to the concentration, which is not realistic at concentrations over 0.1 mol dm<sup>-3</sup>.

The '''[[base dissociation constant]]''' or '''K<sub>b</sub>''' is a measure of basicity.  pKb is the negative log of Kb and related to the [[pKa]] by the simple relationship pK<sub>a</sub>&nbsp;+&nbsp;pK<sub>b</sub>&nbsp;=&nbsp;14.

[[Alkalinity]] is a measure of the ability of a solution to neutralize acids to the equivalence points of carbonates or bicarbonates.

=== Common Bases ===

*[[Baking Soda]]
*[[Ammonia]]

== Neutralization of acids ==

When dissolved in water, the strong base [[sodium hydroxide]] decomposes into hydroxide and sodium ions:

:NaOH → Na<sup>+</sup> + OH<sup>-</sup>

and similarly, in water [[hydrogen chloride]] forms hydronium and chloride ions:

:HCl + H<sub>2</sub>O → H<sub>3</sub>O<sup>+</sup> + Cl<sup>-</sup>

When the two solutions are mixed, the H<sub>3</sub>O<sup>+</sup> and OH<sup>&minus;</sup> ions combine to form water molecules:

:H<sub>3</sub>O<sup>+</sup> + OH<sup>-</sup> → 2 H<sub>2</sub>O

If equal quantities of NaOH and HCl are dissolved, the base and the acid exactly neutralize, leaving only NaCl, effectively [[table salt]], in solution.

Weak bases, such as soda or egg white, should be used to neutralize any acid spills. Neutralizing acid spills with strong bases, such as [[sodium hydroxide]] or [[potassium hydroxide]] can cause a violent exothermic reaction, and the base itself can cause just as much damage as the original acid spill.

== Alkalinity of non-hydroxides == 
Bases are generally compounds that can neutralize an amount of acids. Both [[sodium carbonate]] and [[ammonia]] are bases, although neither of these substances contains OH<sup>&minus;</sup> groups. Both compounds accept H<sup>+</sup> when dissolved in water:

:Na<sub>2</sub>CO<sub>3</sub> + H<sub>2</sub>O → 2 Na<sup>+</sup> + HCO<sub>3</sub><sup>-</sup> + OH<sup>-</sup>
:NH<sub>3</sub> + H<sub>2</sub>O → NH<sub>4</sub><sup>+</sup> + OH<sup>-</sup>

From this, a [[pH]], or acidity, can be calculated for aqueous solutions of bases. Bases also directly act as electron-pair donors themselves:

:CO<sub>3</sub><sup>2-</sup> + H<sup>+</sup> → HCO<sub>3</sub><sup>-</sup>
:NH<sub>3</sub> + H<sup>+</sup> → NH<sub>4</sub><sup>+</sup>

[[Carbon]] can act as a base as well as [[nitrogen]] and [[oxygen]]. This occurs typically in compounds such as [[butyl lithium]], [[alkoxide]]s, and metal [[amide]]s such as [[sodium amide]]. Bases of carbon, nitrogen and oxygen without resonance stabilization are usually very strong, or [[superbase]]s, which cannot exist in a water solution due to the acidity of water. Resonance stabilization, however, enables weaker bases such as carboxylates; for example, [[sodium acetate]] is a [[weak base]].

==Strong bases==
A strong base is a basic chemical compound that is able to deprotonate very weak acids in an acid-base reaction. Compounds with a pKa of more than about 13 are called strong bases. Common examples of strong bases are the hydroxides of alkali metals and alkaline earth metals like NaOH and Ca(OH)<sub>2</sub>. Very strong bases are even able to deprotonate very weakly acidic C-H groups in the absence of water.
Hydroxide compounds in order of strongest to weakest:{{Fact|date=December 2007}}
* [[Potassium hydroxide]] (KOH)
* [[Barium hydroxide]] (Ba(OH)<sub>2</sub>)
* [[Caesium hydroxide]] (CsOH)
* [[Sodium hydroxide]] (NaOH)
* [[Strontium hydroxide]] (Sr(OH)<sub>2</sub>)
* [[Calcium hydroxide]] (Ca(OH)<sub>2</sub>)
* [[Lithium hydroxide]] (LiOH)
* [[Rubidium hydroxide]] (RbOH)

The cations of these strong bases appear in the 1st and 2nd groups of the periodic table (alkali and earth alkali metals).

Group 1 salts of carbanions, amides, and hydrides tend to be even stronger bases due the conjugate acids, which are stable hydrocarbons, amines, and water. Usually these bases are created by adding pure alkali metals such as sodium into the conjugate acid. They are called [[superbase]]s and it is not possible to keep them in water solution, due to the fact they are stronger bases than the hydroxide ion and as such it will deprotonate the conjugate acid water.  For example the ethoxide ion (conjugate base of ethanol) in the presence of water will undergo this reaction.

CH<sub>3</sub>CH<sub>2</sub>O<sup>-</sup> + H<sub>2</sub>O  -->   CH<sub>3</sub>CH<sub>2</sub>OH + OH<sup>-</sup>
 

* [[Butyl lithium]] (n-BuLi)
* [[Lithium diisopropylamide]] (LDA) (C<sub>6</sub>H<sub>14</sub>LiN)
* [[Sodium amide]] (NaNH<sub>2</sub>)
* [[Sodium hydride]] (NaH)

== Bases as heterogeneous catalysts ==

Basic substances can be used as [[solubility|insoluble]] heterogeneous [[catalyst]]s for [[chemical reaction]]s. Examples are metal oxides such as [[magnesium oxide]], [[calcium oxide]], and [[barium oxide]] as well as [[potassium fluoride]] on [[alumina]] and some [[zeolite]]s. A great deal of [[transition metal]]s make good catalysts, many of which form basic substances. Basic catalysts have been used for [[hydrogenation]]s, the migration of [[double bond]]s, in the [[Meerwein-Ponndorf-Verley reduction]], the [[Michael reaction]], and many other reactions.

== See also ==
* [[Acid-base reaction]]s
* [[Acid]]s

==External links==
* Acid-Base equilibrium diagrams, pH calculation and titration curves simulation and analysis - freeware [http://www2.iq.usp.br/docente/gutz/Curtipot_.html Link]

==References==
{{reflist|2}}

[[Category:Chemical compounds]]
[[Category:Bases]]

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