Base (chemistry)

For the term in genetics, see base (genetics)

Acids and Bases
Acid dissociation constant Acid-base extraction Acid–base reaction Acid–base titration Dissociation constant Acidity function Buffer solutions pH Proton affinity Self-ionization of water
Acid types
Brønsted · Lewis · Mineral Organic · Strong Superacids · Weak
Base types
Brønsted · Lewis · Organic Strong · Superbases Non-nucleophilic · Weak
v · d · e

A base in chemistry is a substance that can accept hydrogen ions (protons) or more generally, donate electron pairs. A soluble base is referred to as an alkali if it contains and releases hydroxide ions (OH⁻) quantitatively. The Brønsted-Lowry theory defines bases as proton (hydrogen ion) acceptors, while the more general Lewis theory defines bases as electron pair donors, allowing other Lewis acids than protons to be included.^[1] The oldest Arrhenius theory defines bases as hydroxide anions,^[2] which is strictly applicable only to alkali. In water, by altering the autoionization equilibrium, bases give solutions with a hydrogen ion activity lower than that of pure water, i.e. a pH higher than 7.0 at standard conditions. Examples of common bases are sodium hydroxide and ammonia. Metal oxides, hydroxides and especially alkoxides are basic, and counteranions of weak acids are weak bases.

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 and acids are typically found in aqueous solution forms. Aqueous solutions of bases react with aqueous solutions of acids to produce water and salts in aqueous solutions in which the salts separate into their component ions. If the aqueous solution is a saturated solution with respect to a given salt solute any additional such salt present in the solution will result in formation of a precipitate of the salt.

Definitions

Main article: acid-base reaction theories

A strong base is a base which hydrolyzes completely, raising the pH of the solution toward 14. Concentrated bases, like concentrated acids, attack living tissue and cause serious burns. The reaction of bases upon contact with skin is different from that of acids. So while either may be quite destructive, strong acids are called corrosive, and strong bases are referred to as 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 at standard conditions. An alkali is a special example of a base, where in an aqueous environment, hydroxide ions 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. Rouelle considered that such a substance serves as a base for the salt, giving the salt a "concrete or solid form”.^[3]

Properties

Some general properties of bases include

• Slimy or soapy feel on fingers, due to saponification of the lipids in human skin.
• Concentrated or strong bases are caustic 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, phenolphthalein pink, keep bromothymol blue in its natural colour of blue, and turns methyl orange yellow.
• The pH level of a basic solution is higher than 7.
• Bases are bitter in taste.^[4]

Bases and pH

The pH of an aqueous sample (water) is a measure of its acidity. In pure water, about one in ten million molecules dissociate into hydronium ions and hydroxide ions according to the following equation:

2H₂O(l) ⇌ H₃O⁺(aq) + OH⁻(aq)

The concentration, measured in molarity (M or moles per litre), of the ions is indicated as [H₃O⁺] and [OH⁻]; their product is the dissociation constant which has the value of 10⁻¹⁴ M². The pH is defined as −log [H₃O⁺]; thus, pure water has a pH of 7. (These numbers are correct at 23 °C and are slightly different at other temperatures.)

A base accepts protons from hydronium ions, or donates hydroxide ions to the solution. Both actions will lower the concentration of hydronium ions, and thus raise the pH. By contrast, an acid donates protons to water or accepts OH⁻, thus increasing the concentration of hydronium and lowering the pH.

For example, if 0.1 mol (4 g) of sodium hydroxide (NaOH) are dissolved in water to make 1 litre of solution, the concentration of hydroxide ions becomes [OH⁻] = 0.1 mol/L. As the ionic product remains a constant value, [H⁺] = 1×10⁻¹⁴/[OH⁻] =  1×10⁻¹³ mol/L, and pH = −log 10⁻¹³ = 13. The base dissociation constant, K_b, is a measure of basicity. It is related to the acid dissociation constant, K_a, by the simple relationship pK_a + pK_b = 14, where pK_b and pK_a are the negative logarithms of K_b and K_a, respectively.

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

Neutralization of acids

When dissolved in water, the strong base sodium hydroxide ionizes 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 neutralize exactly, leaving only NaCl, effectively table salt, in solution.

Weak bases, such as baking 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⁻ 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:

CO₃^2- + H⁺ → HCO₃^-

NH₃ + H⁺ → NH₄⁺

Carbon can act as a base as well as nitrogen and oxygen. This occurs typically in compounds such as butyl lithium, alkoxides, and metal amides such as sodium amide. Bases of carbon, nitrogen and oxygen without resonance stabilization are usually very strong, or superbases, 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. 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. Here is a list of several strong bases:

• Potassium hydroxide (KOH)
• Barium hydroxide (Ba(OH)₂)
• Caesium hydroxide (CsOH)
• Sodium hydroxide (NaOH)
• Strontium hydroxide (Sr(OH)₂)
• Calcium hydroxide (Ca(OH)₂)
• Magnesium hydroxide (Mg(OH)₂)
• Lithium hydroxide (LiOH)
• Rubidium hydroxide (RbOH)

The cations of these strong bases appear in the first and second groups of the periodic table (alkali and earth alkali metals).

Acids with a pK_a of more than about 13 are considered very weak, and their conjugate bases are strong bases.

Superbases

Main article: Superbase

Group 1 salts of carbanions, amides, and hydrides tend to be even stronger bases due to the extreme weakness of their conjugate acids, which are stable hydrocarbons, amines, and dihydrogen. 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 they 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⁻

Here are some superbases:

• Butyl lithium (n-BuLi)
• Lithium diisopropylamide (LDA) (C₆H₁₄LiN)
• Lithium diethylamide (LDEA)
• Sodium amide (NaNH₂)
• Sodium hydride (NaH)
• Lithium bis(trimethylsilyl)amide (((CH₃)₃Si)₂NLi)

Bases as catalysts

Basic substances can be used as insoluble heterogeneous catalysts for chemical reactions. Some examples are metal oxides such as magnesium oxide, calcium oxide, and barium oxide as well as potassium fluoride on alumina and some zeolites. Many 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-Verley reduction, the Michael reaction, and many other reactions.

See also

• Acids
• Acid-base reactions
• Base-richness (used in ecology, referring to environments)
• Conjugate base
• Titration

References

1. ^ Chemistry, 9th Edition. Kenneth W. Whitten, Larry Peck, Raymond E. Davis, Lisa Lockwood, George G. Stanley. (2009) ISBN 0495391638. Page 363
2. ^ Chemistry. Page 349
3. ^ The Origin of the Term Base William B. Jensen Journal of Chemical Education • 1130 Vol. 83 No. 8 August 2006
4. ^ http://www.merriam-webster.com/dictionary/base