Base (chemistry)



Acids and bases:
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In 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 ammonia.

Bases can be thought of as the chemical opposite of salts (or their solutions).

Definitions

Main article: acid-base reaction theories

A strong base is a base which 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.

Properties

Some general properties of bases include:

  • Bitter taste (opposed to sour taste of aldehydes and ketones)
  • Slimy or soapy feel on fingers, due to saponification of the lipids in human skin
  • Concentrated or strong bases are 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 phenolphtalein red

Bases and pH

The hydroxide ions (OH−), according to the following equation:

2H2O(l) → H3O+(aq) + OH-(aq)

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 activity is equivalent to the concentration, which is not realistic at concentrations over 0.1 mol dm-3.

The basicity constant or Kb is a measure of basicity. pKb is the negative log of Kb 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.

Common Bases

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.

Alkalinity of non-hydroxides

Bases are generally compounds that can neutralize an amount of acids. Both ammonia are bases, although neither of these substances contains OH− groups. Both compounds accept H+ when dissolved in water:

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

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

CO32- + H+ → HCO3-
NH3 + H+ → NH4+

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)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:[citation needed]

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 superbases 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.

CH3CH2O- + H2O --> CH3CH2OH + OH-


Bases as heterogeneous catalysts

Basic substances can be used as Michael reaction, and many other reactions.

See also

References

  1. ^ The Origin of the Term Base William B. Jensen Journal of Chemical Education • 1130 Vol. 83 No. 8 August 2006
 
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