Chemistry

Introduction to Acid and Base Concepts



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Acids and bases have various definitions accordingto their structure and function. There are three basic definitions of acids and bases. These are the Bronsted definition of acids and bases and the definition by lewis and the third definition is by Lux and Flood. Bronsted acid is defined as a compound which can donate a proton ( a positively charged hydrogen atom).

The strength of the acid is defined by its PH or the negative logarith of the hydrogen ion concentration. The PH scale can range from 1 to 14, where a PH=7 means neutral solution. This means it is neather acidic nor basic. Very strong acids are called superacids and can have even a negative PH. PH lower than 7 is considered acidic and PH greater than 7 is considered basic.

Bronsted definition of bases is that a molecule which can donate a hydroxyl group or -OH. For Example H2O or water can donate H+ or -OH so it is considered as an acid or a base in the same time. Our body has many acids. However we do not die because of the concentrated acid in the body because there are buffers in the body which control the acidity of the blood. A buffer is a solution of an acid with its conjugate base which keeps the PH within a very limited range.

Examples of Bronsted acids are HCl and H2SO4. Examples of Bronsted base are KOH or NaOH or water. Bronsted Acids can be neutralized using bases in which case the products are a salt and water. Lewis acids are defined as compounds which have empty or vacant P or d orbitals which are available for bonding. An example of a lewis acid is BH3 or AlCl3. In the first example boron has a low lying P orbital which can accept a lone pair.

In the second example Al has a vacant and low lying P orbital which can accept a donor of electrons such as water. The definition of lewis bases is that any compound which has lone pair of electrons which are free to participate in bonding. An example is water which its central atom the oxygen has two lone pairs of electrons which are available for bonding. Therefor water is considered as a lewis base.

The third definition of acids and bases is a different one than the definition by Bronsted and lewis. It describes acids and bases using oxides. Oxygen donors in oxides such as CaO is considered as a base, while oxygen acceptor such as SiO2 is considered as an acid. The reaction between the base CaO and the acid SiO2 gives the salt calcium silicate or CaSiO3. Bronsted acids and bases have found use in organic chemistry as catalysts such as the acidic or basic hydrolysis of esters. Acid halides are also reactive towards -OH and are hydrolyzed to the corresponding acid.

Tertiary alcohols react with bronsted acids to form carbocations and water. carbocations are lewis acids which have a low lying vacant or empt P orbital. Example of the reaction of alcohol with a bronsted acid is the reaction between ethylalcohol and HCl. Lewis acids are also used as caalysts in organi chemistry. They function by binding to a lone pair on a heteroatom such as oxygen thus satbilizing the transition state for the reaction.

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