PH and Pk
• Acidic and basic are two extremes that describe chemicals,
just like hot and cold are two extremes that describe
temperature.
• pH is the negative log of the activity of the hydrogen ion in an
aqueous solution.
• Solutions with a pH less than 7 are said to be acidic
• and solutions with a pH greater than 7 are basic or alkaline.
• Pure water has a pH of 7.
• It has also been suggested that the "p" stands for
the German Potenz (meaning "power").
• Others refer to French puissance (also meaning "power",
based on the fact that the Carlsberg Laboratory was Frenchspeaking).
• Another suggestion is that the "p" stands for the Latin
terms pondus hydrogenii, potentia hydrogenii, or potential
hydrogen.
• A pH level of is 7.0 at 25°C is defined as 'neutral' because
the concentration of H3O+equals the concentration of OH− in
pure water.
• pH is a measure of hydrogen ion concentration;
• pH is calculated by the formula
• pH = - log [H+]
History
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The concept of p[H] was first introduced by Sørensen 1909.
It was revised to the modern pH in 1924.
The exact meaning of the "p" in "pH" is disputed,
But according to the Carlsberg Foundation pH stands for
"power of hydrogen".
• Knowledge of the dissociation of weak acids and bases is
important as it provide base to understand the influence of
intracellular pH on structure and biologic activity.
• Many biochemicals possess functional groups that are weak
acids or bases.
• Carboxyl groups, amino groups, and phosphate esters, whose
second dissociation falls within the physiologic range, are
present in proteins
• and nucleic acids, most coenzymes, and most intermediary
metabolites.
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PH = 1/[H+]
PH = log 1 – log H+
As log 1 =0 so
PH= – log H+
Example 1: What is the pH of a solution whose hydrogen ion
concentration is 3.2x10– 4 mol/L?
PKa
• We term the species which donate H+ ( R—NH3+) the acid and
those species which accept H+ (R—NH2) its conjugate base.
• Similarly, we may refer to a base ( R—NH2) and its conjugate
acid ( R—NH 3+).
• We express the relative strengths of weak acids and bases in
terms of their dissociation constants.
• Shown below are the expressions for the dissociation constant
(K a) for two representative weak acids,
• R—COOH and R—NH3+.
Definition of Pka
• pKa is used to express the relative strengths of both acids and
bases.
• For any weak acid, its conjugate is a strong base.
• Similarly, the conjugate of a strong base is a weak acid. The
relative strengths of bases are expressed in terms of the pKa
of their conjugate acids.
• pK a is related to Ka
• as pH is to [H+].
• The stronger the acid, the lower is its pKa value.
• The strength of an acid can be determined by considering the
extent to which it dissociates in a given solvent.
• The equilibrium constant in water is given by:
• Ka is known as the acidity constant of the acid.
• Generally, the numerical value of Ka will be small, and for this
reason, Ka is usually converted into pKa by the following
equation:
• Thus, the smaller the value of pKa, the stronger the acid.
• Similarly, the strength of a base is determined by using
pKb. For the example of BOH, a base, in water, Kb can be
determined by the following formula:
Buffer Solutions
• A buffer solution is one in which the pH of the solution is
"resistant" to small additions of either a strong acid or strong
base.
•
Types of Buffer Solution
• Acidic buffer solutions:
• An acidic buffer solution is simply one which has a pH less
than 7.
• Acidic buffer solutions are commonly made from a weak acid
and one of its salts - often a sodium salt.
• A common example would be a mixture of Acetic acid and
sodium acetate in solution.
• Alkaline buffer solutions:
• An alkaline buffer solution has a pH greater than
7. Alkaline buffer solutions are commonly made
from a weak base and one of its salts.
• A frequently used example is a mixture of
ammonia solution and ammonium chloride
solution.
How do buffer solutions work?
• A buffer solution has to contain things which will remove any
hydrogen ions or hydroxide ions that you might add to it
• Acidic buffer solutions
• We'll take a mixture of Acetic acid and sodium acetate as
typical.
• Acetic acid is a weak acid, and the position of this equilibrium
will be well to the left:
CH3COOH
CH3COO- + H+
Adding an acid to this buffer
solution
• The buffer solution must remove most of the new hydrogen
ions otherwise the pH would drop markedly.
• Hydrogen ions combine with the Acetate ions to make Acetic
acid.
• Although the reaction is reversible, since the Acetic acid is a
weak acid, most of the new hydrogen ions are removed in this
way.
• CH3COO- + H+
CH3COOH
Adding an alkali to this buffer
solution
• Alkaline solutions contain hydroxide ions and the buffer
solution removes most of these H+.
• The most likely acidic substance which a hydroxide ion is going
to collide with is an acetic acid molecule.
• They will react to form acetate ions and water.
• CH3COOH + OH -
CH3COO- + H2O
Adding Acid to alkaline buffer
• The most likely basic substance which a hydrogen ion is going
to collide with is an ammonia molecule.
• They will react to form ammonium ions.
Adding an alkali to this
buffer solution
• The hydroxide ions from the alkali are removed by a simple
reaction with ammonium ions.
References;
• Harper's Illustrated Biochemistry,
• Color atlas of biochemistry by J. Koolman and K.H. Roehm
• Wikipedia.org
Thank You