Analytical Chemistry
1st Class 1st Semester 2012/2013
Reference text:1- Fundamentals of Analytical Chemistry by Stook and West
--------------------------------------------------------------------------------------Lecture No. 1
Date : 1/12 /2012
Dr. Mohammed Hamed
Historic acid–base theories
Lavoisier's oxygen theory of acids
The first scientific concept of acids and bases was provided by Lavoisier circa
1776. Since Lavoisier's knowledge of strong acids was mainly restricted to
oxoacids, such as HNO3 (nitric acid) and H2SO4 (sulphuric acid), which tend to
contain central atoms in high oxidation states surrounded by oxygen.
The Lavoisier definition was held as absolute truth for over 30 years, until the 1810
article and subsequent lectures by Sir Humphry Davy in which he proved the lack
of oxygen in H2S, H2Te, and the hydrohalic acids. However, Davy failed to develop
a new theory, concluding that "acidity does not depend upon any particular
elementary substance, but upon peculiar arrangement of various substances". One
notable modification of oxygen theory was provided by Berzelius, who stated that
acids are oxides of nonmetals while bases are oxides of metals.
Liebig's hydrogen theory of acids
This definition was proposed by Justus von Liebig circa 1838, based on his
extensive works on the chemical composition of organic acids. This finished the
doctrinal shift from oxygen-based acids to hydrogen-based acids, started by Davy.
According to Liebig, an acid is a hydrogen-containing substance in which the
hydrogen could be replaced by a metal.
In 1884 Svante Arrhenius suggested that salts such as NaCl dissociate when they
dissolve in water to give particles he called ions.
H2O
NaCl(s)
Na+(aq) + Cl-(aq)
Three years later Arrhenius extended this theory by suggesting that acids are
neutral compounds that ionize when they dissolve in water to give H+ ions and a
corresponding negative ion. According to his theory, hydrogen chloride is an acid
because it ionizes when it dissolves in water to give hydrogen (H+) and chloride
(Cl-) ions .
Arrhenius argued that bases are neutral compounds that either dissociate or ionize
in water to give OH- ions and a positive ion. NaOH is an Arrhenius base because it
dissociates in water to give the hydroxide (OH-) and sodium (Na+) ions.
Arrhenius acid - Any substance that ionizes when it dissolves in water to give the
H+ ion.
Arrhenius base - Any substance that ionizes when it dissolves in water to give the
OH- ion.
Arrhenius acids include compounds such as HCl, HCN, and H 2SO4 that ionize in
water to give the H+ ion. Arrhenius bases include ionic compounds that contain the
OH- ion, such as NaOH, KOH, and Ca(OH)2.
Arrhenius's theory helped to explain why acids have similar characteristics, since
they all give H+ ions when they dissolve in water. It also explained why acids are
neutralized by bases and why bases are neutralized by acids; the H+ ions from
acids combine with the OH- ions from bases to form water:
Though the Arrhenius theory helped to explain more about acids and bases, there
were still several drawbacks to this theory.
 The theory can only classify substances when they are dissolved in water since
the definitions are based upon the dissociation of compounds in water.
 It does not explain why some compounds containing hydrogen such as HCl
dissolve in water to give acidic solutions and why others such as CH4 do not.
 The theory can only classify substances as bases if they contain the OH- ion and
cannot explain why some compounds that don't contain the OH- such as Na2CO3
have base-like characteristics.
To extend the Arrhenius theory a little further, consider the formation of water from
the combination of an H+ ion and an OH- ion:
This reaction is actually reversible, represented by the forward/backward arrow in
the following reaction:
Based on the fact that the above reaction is reversible, we can conclude the
following operational definitions for acids and bases.
 Acid - Any substance that increases the concentration of the H+ ion when it
dissolves in water.
Base - Any substance that increase the concentration of the OH- ion when it
dissolves in water.
NOTE: A common shorthand notation for the concentration of a substance is
placing it in brackets:

[H+] = concentration of the H+ ion
Now, substances not containing H+ or OH- ions can be classified as acids or bases if
they alter the [H+] or [OH-] when they dissolve in water.
e.g. CO2 cannot dissociate to give H+ but it does increase [H+] when it is dissolves
in water.
e.g. CaO cannot dissociate to give OH- but it does increase [OH-] when it
dissolves in water.
In 1923, Johannes Brønsted and Thomas Lowry separately proposed a new set of
defintions for acids and bases which are known as either Brønsted acids and bases
or Brønsted-Lowry acids and bases.
Brønsted Acid - Any substance that can donate a proton, H+ ion to a base.
aka: hydrogen-ion donors or proton donors
Brønsted Base - Any substance that can accept a proton, H+ ion from an acid.
aka: hydrogen-ion acceptor or proton acceptor
In the above reaction, the H from HCl is donated to H 2O which accepts the H to
form H3O+, leaving a Cl- ion.
The dissociation of water can be represented as follows :
The Brønsted-Lowry model of acids and bases brings rise to the concept of
conjugate acid-base pairs. The part of the acid remaining when an acid donates a
H+ ion is called the conjugate base. The acid formed when a base accepts a H+ ion
is called the conjugate acid. For the generic acid HA:
For the generic base A-:
More examples of conjugate acid-base pairs:
In the following reactions, it is shown how H2PO4- and H2O can act as both acids
and bases. Such compounds are said to be amphoteric.
Amphoteric.substance: A molecule or ion that can behave as either a proton
donor or a proton accepter.
- Strong acids have weak conjugate bases.
- Strong bases have weak conjugate acids.
Water has the tendency to equalize the strengths of all strong acids and strong
bases, regardless of the strength of the acid itself. This is known as the leveling
effect. Acids are limited to the strength of the H3O+ ions that they form when they
lose H+ ions when they dissolve in water. Likewise, bases are limited to the
strength of the OH- ions that they form when they gain H+ ions when they dissolve
in water.
The relative strength of an acid is described as an acid-dissociation equilibrium
constant.
Acid-dissociation equilibrium constant (Ka) - For the generic acid reaction with
water:
The acid-dissociation equilibrium constant is the mathematical product of the
equilibrium concentrations of the products of this reaction divided by the
equilibrium concentration of the original acid:
Strong acids dissociate almost completely in water and therefore have relatively
large Ka values. Weak acids, however, dissociate only slightly in water and
therefore have relatively small Ka values. To distinguish between strong and weak
acids, the following guidelines are used:
Advantages to the Brønsted-Lowry model of acids and bases
- Acids and bases can now be ions or neutral molecules.
- Acids and bases can now be any molecule with at least one pair of nonbonding
electrons.
- It explains the role of water in acid-base reactions; H2O accepts H+ ions from
acids to form H3O+ ions.
- It can be applied to solutions with solvents other than water and even in reactions
that occur in the gas or solid phases.
- It relates acids and bases to each other with conjugate acid-base pairs and can
explain their relative strengths.
- It can explain the relative strengths of pairs of acids or pairs of bases.
- It can explain the leveling effect of water.
Proposed another method of defining acids and bases.
Lewis acid - Any substance that can accept a pair of nonbonding electrons.
aka: electron-pair acceptor
- Lewis base - Any substance that can donate a pair of nonbonding electrons.
aka: electron-pair donor
In the following example, the Al3+ ion acts as an acid, accepting electron pairs
from water which acts as a base, an electron-pair donor. The two combine to form
Al(H2O)63+,
an
acid-base
complex
or
a
complex
ion.