BASES

A base in chemistry is a substance that can accept hydrogen ions 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. The oldest Arrhenius theory defines bases as hydroxide anions, 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.

ACIDS

An acid (from the Latin acidus/acēre meaning sour) is a substance which reacts with a base. Commonly, acids can be identified as tasting sour, reacting with metals such as calcium, and bases like sodium carbonate. Aqueous acids have pHs of less than 7, and turn blue litmus paper red. Chemicals or substances having the property of an acid are said to be acidic.

Common examples of acids include acetic acid (in vinegar), sulfuric acid (used in car batteries), and tartaric acid (used in baking). As these three examples show, acids can be solutions, liquids, or solids. Gases such as hydrogen chloride can be acids as well.

There are three common definitions for acids: the Arrhenius definition, the Brønsted-Lowry definition, and the Lewis definition. The Arrhenius definition states that acids are substances which increase the concentration of hydronium ions (H₃O⁺) in solution. The Brønsted-Lowry definition is an expansion: an acid is a substance which can act as a proton donor. Most acids encountered in everyday life are aqueous solutions, or can be dissolved in water, and these two definitions are most relevant. The reason why pHs of acids are less than 7 is that the concentration of hydronium ions is greater than 10⁻⁷ moles per liter. Since pH is defined as the negative logarithm of the concentration of hydronium ions, acids thus have pHs of less than 7. By the Brønsted-Lowry definition, any compound which can easily be deprotonated can be considered an acid. Examples include alcohols and amines which contain O-H or N-H fragments.

In chemistry, the Lewis definition of acidity is frequently encountered. Lewis acids are electron-pair acceptors. Examples of Lewis acids include all metal cations, and electron-deficient molecules such as boron trifluoride and aluminum trichloride. Hydronium ions are acids according to all three definitions. Interestingly, alcohols and amines mentioned above as examples of Brønsted-Lowry can function as Lewis bases at the same time.

Properties

It is easy to describe acids in terms of what they do. Among their most familiar properties, they taste sour and corrode metals. Vinegar, for example, gets its tart taste from acetic acid.

Vinegar is a fairly dilute form of acetic acid. At full strength, this acid blisters the skin. Sulfuric acid is so corrosive that even a drop can cause deeply serious burns. In some very weak acids such as carbonic acid, the properties of sourness and corrosion are hardly noticeable to the 'nonchemist'. Some weak acids are even used as medicines, including acetylsalicylic acid, better known as aspirin.

Bases are best known for bitter taste and the soapy feel they impart to water. In fact, many bases such as sodium hydroxide are essential ingredients in a number of familiar household and industrial soaps. The base ammonia is a household and industrial cleaner in its own right. Strong bases, like strong acids, can be quite corrosive and burn the skin.

In general, acids are compounds which, when dissolved in water

• form solutions that conduct electricity
• react with metals like zinc and magnesium to produce salts and a gas which is usually hydrogen
• turns blue litmus paper to red
• react with bases to form salt and water.

On the other hand, bases are compounds which, when dissolved in water

• yield solutions that also conduct electric current
• turns red litmus paper to blue
• react with acids to form salt and water.

ACIDS, BASES, and SALTS: An Introduction

The substances known as acids, bases, and salts are among the most important and best studied in Chemistry. They are widespread in nature, in industry, and in the home.

Hydrochloric acid, present in small quantities in our stomachs, helps digest our food. Ascorbic acid, or vitamin C, is among our most important nutrients. Carbonic acid is familiar to us in the form of soda water. And the much stronger sulfuric acid is used to manufacture fertilizers, paints, synthetic fibers, and plastics.

Among the more familiar bases are lye, or caustic soda, and ammonia water. By far, the best-known salt is sodium chloride, which is commonly known as table salt. It is a basic ingredient of seawater as well as of our own blood.