Monday, February 15, 2010

ACID BASE Reactions


Acid base reactions
Strength of an acid
Concepts of pH & pOH
Henderson Hasselbalch Equation


The word "acid" comes from the Latin acidus meaning "sour,"
Chemical compound when combined with water gives a pH less than 7

Different definitions are given by

Solvent-system definition:
Acetic acid
In vinegar
Sulphuric acids
In car batteries

Lavoisier's definition

The first scientific definition was proposed by the French chemist Antoine Lavoisier. Since Lavoisier's knowledge of strong acids was mainly restricted to oxoacids, which tend to contain central atoms in high oxidation states surrounded by oxygen, such as HNO3 and H2SO4, and since he was not aware of the true composition of the hydrohalic acids, HCl, HBr, and HI, He defined acids in terms of their containing oxygen which in fact he named from Greek words meaning "acid-former". When the elements chlorine, bromine, and iodine were identified and the absence of oxygen in the hydrohalic acids was established by Sir Humphry Davy in 1810, this definition had to be rejected.

Acid - Arrhenius:

Swedish chemist Svante Arrhenius
An acid is a substance that increases the concentration of hydronium ion (H3O+) when dissolved in water, while Bases are substances that increase the concentration of hydroxide ions (OH-).
This definition limits acids and bases to substances that can dissolve in water.
Around 1800, many French chemists, including Antoine Lavoisier, incorrectly believed that all acids contained oxygen. English chemists, including Sir Humphry Davy at the same time believed all acids contained hydrogen. Arrhenius used this belief to develop this definition of acid.

Acid - Brønsted-Lowry

An  acid is a proton (hydrogen nucleus) donor and
A  base is a proton acceptor.
The acid is said to be dissociated after the proton is donated.
An acid and the corresponding base are referred to as conjugate acid-base pairs.
Brønsted and Lowry independently formulated this definition, which includes water-insoluble substances not in the Arrhenius definition.
The protonic (Brønsted-Lowry)
The Brønsted-Lowry definition, formulated independently by its two proponents
Johannes Nicolaus Brønsted , Martin Lowry in 1923 revolves around an acid's ability to donate protons (H+) to another compound, called a base, in a chemical reaction.
A base is a proton acceptor.
In Brønsted-Lowry acid-base reactions, there is a "competition" between two bases for a proton, so that if X and Y are two species, the equilibrium
HX + Y- ↔ HY + X- occurs.
Both HX and HY are Brønsted-Lowry acids;
both X- and Y- are Brønsted-Lowry bases.
If the reaction runs mostly to the left, then HY is the stronger acid and X- the stronger base;
if the reaction runs mostly to the right, then HX is the stronger acid and Y- the stronger base.

Acid-solvent-system definition

According to this definition, an acid is a substance that, when dissolved in an autodissociating solvent, increases the concentration of the solvonium cations, such as
H3O+ in water,
NH4+ in liquid ammonia,
NO+ in liquid N2O4,
SbCl2+ in SbCl3, etc.
Base is defined as the substance that increases the concentration of the solvate anions, respectively
OH-, NH2-, NO3-, or SbCl4-.
This definition extends acid-base reactions to nonaqueous systems and even some aprotic systems, where no hydrogen nuclei are involved in the reactions.
This definition is not absolute, a compound acting as acid in one solvent may act as a base in another.

Acid -Lewis

According to this definition developed by Gilbert N. Lewis,
an acid is an electron-pair acceptor and
a base is an electron-pair donor.
(These are frequently referred to as "Lewis acids" and "Lewis bases," and are electrophiles and nucleophiles, respectively, in organic chemistry; Lewis bases are also ligands in coordination chemistry.)
Lewis acids include substances with no transferable protons (ie H+ hydrogen ions), such as iron(III) chloride, and hence the Lewis definition of an acid has wider application than the Brønsted-Lowry definition.
The Lewis definition can also be explained with molecular orbital theory. In general, an acid can receive an electron pair in its lowest unoccupied orbital (LUMO) from the highest occupied orbital (HOMO) of a base.  That is, the HOMO from the base and the LUMO from the acid combine to a bonding molecular orbital. Although not the most general theory, the Brønsted-Lowry definition is the most widely used.

The Usanovich definition

The most general definition is that of the Russian chemist Mikhail Usanovich, and can basically be summarized as defining an acid as anything that accepts negative species or donates positive ones, and
a base as the reverse. This tends to overlap the concept of redox (oxidation-reduction), and so is not highly favored by chemists

The Lux-Flood definition

This definition, proposed by German chemist Hermann in 1939, further improved by Håkon Flood circa 1947 commonly used in modern geochemistry and electrochemistry of molten salts, describes an acid as an oxide ion acceptor and a base as an oxide ion donor. For example:
MgO (base) + CO2 (acid) → MgCO3
CaO (base) + SiO2 (acid) → CaSiO3
NO3- (base) + S2O72- (acid) → NO2+ + 2SO42-[16]

The Pearson definition

In 1963 Ralph Pearson proposed an advanced qualitative concept known as Hard Soft Acid Base principle, later made quantitative with help of Robert Parr in 1984. 'Hard' applies to species which are small, have high charge states, and are weakly polarizable. 'Soft' applies to species which are large, have low charge states and are strongly polarizable. Acids and bases interact and the most stable interactions are hard-hard and soft-soft. This theory has found use in both organic and inogranic chemistry.
Base - Chemical compound that absorbs H3O+ ion in aqueous solution or a proton accepter

Strength of an acid

The strength of an acid is its ability to donate H+ ions
It is best described by acid dissociation constant
Acid dissociation constant, denoted by Ka, is an equilibrium constant for the dissociation of a weak acid.
 According to the Brønsted-Lowry theory of acids and bases an acid is only recognized by its reaction with a base.
In aqueous solution, the base is water itself.
HA +H2O   A- + H3O+
Acid dissociation constants are also known as
the acidity constant or
the acid-ionization constant.
The term is also used for pKa,
which is equal to negative  logarithm of Ka
Acid dissociation constant (Ka)
When an acid, HA, dissolves in water,
some molecules of the acid 'dissociate' to form
hydronium ions (H3O+ as H+) and
the conjugate base, (A-), of the acid.
HA ----> H+ + A-
The dissociation constant Ka can be written as
Ka = [H+][A-] / [HA]
where the square brackets are usually taken to signify concentration.
H2O is omitted from these expressions because in dilute solution the concentration of water may be assumed to be constant.
pKa = -log Ka
Weak and strong acids
A weak acid may be defined as an acid with pKa greater than about -2.
An acid with pKa = -2 would be 99% dissociated at pH 0, that is, in a 1M solution.
Any acid with a pKa less than about -2 is said to be a strong acid.
Strong acids are said to be fully dissociated.
On the pKa scale of acid strength, a large values indicates a very weak acid, and a small value indicates a not so weak one.

Ionization of water

1 L of water at 25 degrees will have 10-7  moles of H+ and 10-7 moles of OH-
H2O ------> H+ + OH-
Equilibrium constant
Keq = [H+] [OH-] / [H2O]
Kw, the ion product of water
 =Kw = 55.5 Keq
= 10-14 = [H+][OH-]

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