Properties and Reactions of Bases
- Concentrated or strong bases are caustic on organic matter and react violently with acidic substances.
- Aqueous solutions or molten bases dissociate into ions and conduct electricity.
- Bases turn red litmus paper blue, phenolphthalein pink, keep bromothymol blue in its natural blue color, and turn methyl orange-yellow.
- The pH of a basic solution at standard conditions is greater than seven.
- Bases are bitter.
Neutralization and Alkalinity of Bases
- Bases react with water to produce a conjugate acid and a conjugate base.
- The equilibrium constant (Kb) for this reaction can be found using the equation: Kb = [BH+][OH-]/[B].
- Bases that react with water have relatively small equilibrium constant values.
- Bases with lower equilibrium constant values are weaker.
- Some bases, like ammonia, react with water to increase the concentration of hydroxide ions.
- Bases react with acids to neutralize each other at a fast rate.
- When dissolved in water, strong bases like sodium hydroxide ionize into hydroxide and sodium ions.
- When solutions of a base and an acid are mixed, the H+ and OH- ions combine to form water molecules.
- Equal quantities of a base and an acid neutralize, leaving only a salt in solution.
- Weak bases, like baking soda, should be used to neutralize acid spills.
- Bases can neutralize acids even if they don't contain OH groups, like sodium carbonate and ammonia.
- Both sodium carbonate and ammonia accept H+ when dissolved in water.
- pH can be calculated for aqueous solutions of bases.
- Carbon, nitrogen, and oxygen can act as bases by accepting an electron pair bond.
- Compounds with resonance stabilization, like sodium acetate, are weaker bases.
Strong Bases
- Strong bases can remove a proton from even a very weak acid in an acid-base reaction.
- Examples of strong bases include hydroxides of alkali metals and alkaline earth metals.
- Some strong bases, like alkaline earth hydroxides, have low solubility.
- Low solubility allows for the use of metal hydroxides in antacids to prevent harm to tissues.
- Strong bases can stop an increase in the concentration of the hydroxide ion.
- Superbases are stronger bases than the hydroxide ion.
- Examples of superbases include ethoxide ion, meta-diethynylbenzene dianion, para-diethynylbenzene dianion, and lithium monoxide anion.
Weak Bases and Lewis Bases
- A weak base does not fully ionize in an aqueous solution.
- Protonation of a weak base is incomplete.
- Ammonia is an example of a weak base.
- The equilibrium constant for the reaction between ammonia and water is small.
- Weak bases have a limited extent of reaction or degree of ionization.
- A Lewis base is a molecule with high-energy lone pairs of electrons.
- Lewis bases can form adducts with low-energy vacant orbitals in acceptor molecules.
- Electron-pair acceptors include neutral molecules and high oxidation state metal ions.
- Adducts involving metal ions are described as coordination complexes.
- Lewis acid-base reactions involve the sharing of an electron pair and the creation of a high dipole moment.
Uses and Applications of Bases
- Basic substances can be used as insoluble heterogeneous catalysts.
- Examples include metal oxides like magnesium oxide, calcium oxide, and barium oxide.
- Potassium fluoride on alumina and some zeolites also serve as basic catalysts.
- Transition metals make good catalysts, many of which are basic substances.
- Basic catalysts are used in various chemical reactions such as hydrogenation and the Michael reaction.
- Bases are used in the production of various chemicals and materials.
- They are utilized in the manufacturing of detergents and cleaning agents.
- Bases play a role in wastewater treatment to neutralize acidity.
- Many biological processes, such as enzyme activity, are pH-dependent.
- Bases are used in agriculture to adjust soil pH for optimal plant growth.
- Sodium hydroxide is used in soap, paper, and rayon manufacturing.
- Calcium hydroxide is used in the manufacture of bleaching powder and for cleaning sulfur dioxide in power plants.
- Magnesium hydroxide is used as an antacid for stomach acidity.
- Sodium carbonate is used as washing soda and for water softening.
- Sodium bicarbonate is used in cooking, baking powders, and as an antacid.
In chemistry, there are three definitions in common use of the word "base": Arrhenius bases, Brønsted bases, and Lewis bases. All definitions agree that bases are substances that react with acids, as originally proposed by G.-F. Rouelle in the mid-18th century.
In 1884, Svante Arrhenius proposed that a base is a substance which dissociates in aqueous solution to form hydroxide ions OH−. These ions can react with hydrogen ions (H+ according to Arrhenius) from the dissociation of acids to form water in an acid–base reaction. A base was therefore a metal hydroxide such as NaOH or Ca(OH)2. Such aqueous hydroxide solutions were also described by certain characteristic properties. They are slippery to the touch, can taste bitter and change the color of pH indicators (e.g., turn red litmus paper blue).
In water, by altering the autoionization equilibrium, bases yield solutions in which the hydrogen ion activity is lower than it is in pure water, i.e., the water has a pH higher than 7.0 at standard conditions. A soluble base is called an alkali if it contains and releases OH− ions quantitatively. Metal oxides, hydroxides, and especially alkoxides are basic, and conjugate bases of weak acids are weak bases.
Bases and acids are seen as chemical opposites because the effect of an acid is to increase the hydronium (H3O+) concentration in water, whereas bases reduce this concentration. A reaction between aqueous solutions of an acid and a base is called neutralization, producing a solution of water and a salt in which the salt separates into its component ions. If the aqueous solution is saturated with a given salt solute, any additional such salt precipitates out of the solution.
In the more general Brønsted–Lowry acid–base theory (1923), a base is a substance that can accept hydrogen cations (H+)—otherwise known as protons. This does include aqueous hydroxides since OH− does react with H+ to form water, so that Arrhenius bases are a subset of Brønsted bases. However, there are also other Brønsted bases which accept protons, such as aqueous solutions of ammonia (NH3) or its organic derivatives (amines). These bases do not contain a hydroxide ion but nevertheless react with water, resulting in an increase in the concentration of hydroxide ion. Also, some non-aqueous solvents contain Brønsted bases which react with solvated protons. For example in liquid ammonia, NH2− is the basic ion species which accepts protons from NH4+, the acidic species in this solvent.
G. N. Lewis realized that water, ammonia, and other bases can form a bond with a proton due to the unshared pair of electrons that the bases possess. In the Lewis theory, a base is an electron pair donor which can share a pair of electrons with an electron acceptor which is described as a Lewis acid. The Lewis theory is more general than the Brønsted model because the Lewis acid is not necessarily a proton, but can be another molecule (or ion) with a vacant low-lying orbital which can accept a pair of electrons. One notable example is boron trifluoride (BF3).
Some other definitions of both bases and acids have been proposed in the past, but are not commonly used today.
