Base (chemistry)

Encyclopedia : B : BA : BAS : Base (chemistry)

Acids and bases:
Acid-base reaction theories
pH
Self-ionization of water
Buffer solutions
Systematic naming
Electrochemistry
Acids:
Strong acids
Weak acids
Bases:
Strong bases
Weak bases
The common definition of a base is a chemical compound that absorbs hydronium ions when dissolved in water (a proton acceptor). An alkali is a special example of a base, where in an aqueous environment, hydroxide ions are donated. 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. Arrhenius bases are water-soluble and these solutions always have a pH greater than 7.

There are other more generalized and advanced definitions of acids and bases.

Contents

1 Bases and pH
2 Neutralization of acids
3 Alkalinity of non-hydroxides
4 Bases as heterogeneous catalysts
5 See also
6 External links

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 the equivalent moles per liter, of the ions is indicated as [H₃O⁺] and [OH⁻] their product is the dissociation constant of water with and has the value 10⁻¹⁴ 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.

The pH of a solution can be calculated. 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.

Bases are slightly less viscous than pure water, have a bitter taste and are soapy to the touch. They react with acids to form salts.

Neutralization of acids

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

[\mbox\to \mbox^++\mbox^-]

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

[\mbox + \mbox_2\mbox\to \mbox_3\mbox^++\mbox^-.]

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

[\mbox_3\mbox^++\mbox^-\to\mbox_2\mbox]

If equal quantities of NaOH and HCl are dissolved, the base and the acid exactly neutralize, leaving only NaCl, effectively table salt, in solution. Acids have pH levels below 7

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:

[\mbox_2\mbox_3+\mbox_2\mbox\to2\mbox^++\mbox_3^-+\mbox^-]

[\mbox_3+\mbox_2\mbox\to\mbox_4^++\mbox^-.]

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, Meerwein-Ponndorf-Verlay reduction, the Michael reaction, and many other reactions.

External links

[Acid-Base equilibrium diagrams, pH calculation and titration curves simulation and analysis - freeware]