Base (chemistry)

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Acids and bases:

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In chemistry, a base is most commonly thought of as a substance that can accept protons. This refers to the Brønsted-Lowry theory of acids and bases. Alternate definitions of bases include electron pair donors (Lewis), and as sources of hydroxide anions (Arrhenius). Examples of simple bases are sodium hydroxide and ammonia.

Bases can be thought of as the chemical opposite of acids. A reaction between an acid and base is called neutralization. Bases and acids are seen as opposites because the effect of an acid is to increase the hydronium ion (H3O+) concentration in water, whereas bases reduce this concentration. Bases react with acids to produce water and salts (or their solutions). Some general properties of bases include:

Contents

  • 1 Definitions
  • 2 Bases and pH
  • 3 Neutralization of acids
  • 4 Alkalinity of non-hydroxides
  • 5 Strong bases
  • 6 Bases as heterogeneous catalysts
  • 7 See also
  • 8 External links
  • 9 References

[edit] Definitions

Main article: acid-base reaction theories

A strong base is a base which 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 caustic. Superbases are a class of especially basic compounds and non-nucleophilic bases are a special class of strong bases with poor nucleophilicity. Bases may also be weak bases 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 ions(also viewed as OH-) are donated. There are other more generalized and advanced 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 [1] and hence the name.

[edit] 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 ions (H3O+) and hydroxide ions (OH−), according to the following equation:

2H2O(l) → H3O+(aq) + OH-(aq)

The concentration, measured in molarity (M or moles per dm³), of the ions is indicated as [H3O+] and [OH−]; their product is the dissociation constant of water with and has the value 10−7 M. The pH is defined as −log [H3O+]; thus, pure water has a pH of 7. (These numbers are correct at 23 °C and slightly different at other temperatures.)

A base accepts (removes) hydronium ions (H3O+) from the solution, or donates hydroxide ions (OH-) to the solution. Both actions will lower the concentration of hydronium ions, and thus raise pH. By contrast, an acid donates H3O+ ions to the solution or accepts OH−, thus lowering pH.

For example, if 1 mole of sodium hydroxide (40 g) is dissolved in 1 litre of water, the concentration of hydroxide ions becomes [OH−] = 1 mol/L. Therefore [H+] = 10−14 mol/L, and pH = −log 10−14 = 14.

The basicity constant or pKb is a measure of basicity and related to the pKa by the simple relationship pKa + pKb = 14.

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

[edit] Neutralization of acids

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

NaOH → Na+ + OH-

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

HCl + H2O → H3O+ + Cl-

When the two solutions are mixed, the H3O+ and OH− ions combine to form water molecules:

H3O+ + OH- → 2 H2O

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.

[edit] Alkalinity of non-hydroxides

Both sodium carbonate and ammonia are bases, although neither of these substances contains OH− groups. That is because both compounds accept H+ when dissolved in water:

Na2CO3 + H2O → 2 Na+ + HCO3- + OH-
NH3 + H2O → NH4+ + OH-

Carbon can act as a base as well as nitrogen and oxygen. This occurs typically in compounds such as butyl lithium

[edit] 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)2. 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:

The cations of these strong bases appear in the 1st and 2nd groups of the Periodic Table (Alkali and Alkali-Earth 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.

[edit] Bases as heterogeneous catalysts

Basic substances can be used as insoluble heterogeneous catalysts for chemical reactions. Examples are metal oxides such as magnesium oxide, calcium oxide, and barium oxide as well as potassium fluoride on alumina and some zeolites. A great deal of transition metals make good catalysts, many of which form basic substances. Basic catalysts have been used for hydrogenations, the migration of double bonds, in the Meerwein-Ponndorf-Verlay reduction, the Michael reaction, and many other reactions.