Pharmaceutical analytical chemistry I
Lecture 5
Dr. Reem Youssif Shahin
Aqueous solutions
An aqueous solution is a solution in which the solvent is water
Seawater Rain Cola
Juice Vinegar
Aqueous Solutions and Solubility
• How do solids such as salt and sugar dissolve in water?
Water dissolve salts by separation of cations and anions and formation of new interactions
between water and ions: when sodium chloride is put into water, it separates into Na+ cations and
Cl - anions .
As a result of the polar nature of the water molecule, the oxygen atom in water is electron-
rich, giving it a partial negative charge (Ƃ- ), while The hydrogen atoms, are electron-poor,
giving them a partial positive charge (Ƃ+),thus a dipole is formed.
As a result, the positively charged sodium ions are strongly attracted to the oxygen side
of the water molecule and the negatively charged chloride ions are attracted to the
hydrogen side of the water molecule (ion dipole interaction).
Hydration shell (hydration sphere) is formed surrounding solute molecules with water molecules.
the attraction between the separated ions and the water molecules overcomes the attraction of
sodium and chloride ions to each other, and the sodium chloride finally dissolves in the water
In summary, water can dissolve
ionic substances (universal
solvent) by two mechanism ion-
dipole interaction and hydrogen
bond formation.
Sodium Chloride Dissolving in Water :The
attraction between water molecules and the ions of sodium
chloride causes NaCl to dissolve in the water.
• In contrast to sodium chloride, sugar is a molecular compound. Most
molecular compounds dissolve in water as intact molecules.
• Sugar dissolves because the attraction between sugar molecules
and water molecules overcomes the attraction of sugar molecules to
each other
A Sugar Solution
Electrolyte and Nonelectrolyte Solutions
• Electrolyte: substance able to conduct electricity as it ionizes greatly in water.
Two types of electrolytes are well known strong and weak electrolytes.
Item Strong electrolyte Weak electrolyte
Ionization Higher Lower
Electric conduction Higher Lower
Acids HCl, H2SO4, HNO3, HClO4 HCN, H3BO3, CH3COOH,
H3PO4, H2CO3
Bases NaOH, KOH Al(OH)3
• Non-electrolyte: Substance not ionized in water and no ability to conduct
electricity e.g. sugars and glycerol.
Note that: All salts are strong electrolyte such as NaCl, KCl, CH3COONa. 6
comparison between electrolytes and non
electrolytes.
Electrical conductivity of aqueous solutions: The circuit will be completed and will allow current to flow only when there are
charge carriers (ions) in the solution.
(a) A hydrochloric acid solution, which is a strong electrolyte, contains ions that readily conduct the current and give a brightly
lit bulb. (b) An acetic acid solution, which is a weak electrolyte, contains only a few ions and does not conduct as much current
as a strong electrolyte. The bulb is only weakly lit. (c) A sucrose solution, which is a nonelectrolyte, contains no ions and does
not conduct a current. The bulb remains unlit.
` 2-Weak electrolytes: when dissolved
in water , it is weakly ionized or
ionized to a much less extent, so it is a
weak conductor for electricity
1-Strong electrolytes: when .Example: acetic acid: CH3COOH
dissolved in water , it is
completely ionized as NaCl
HCl is assumed to be 100
percent dissociated into ions in
solution 3-Non electrolytes: dissolve
as molecules in water, so do
so it is a good conductor for not produce ions, and do not
electricity conduct electricity. Example
A solution’s sugar
ability to conduct
electricity
depends on
number of ions it
contains
Acids and bases
Definitions of Acids and Bases:
I- The Arrhenius Definition:
Acid: A substance that produces H+ ions in
aqueous solution.
Base: A substance that produces OH- ions in
aqueous solution.
According to the Arrhenius definition, HCl is
an acid because it produces H+ ions in
solution :
HCl(aq)→ H+ (aq) + Cl-(aq)
According to the Arrhenius definition, NaOH
is a base because it produces OH- ions
in solution:
NaOH(aq)→ Na+ (aq) + OH-(aq)
ü Limitations of Arrhenius theory:
1. Acids and bases e.g. HCl, H2SO4, NaOH according to these
definitions, include only electrically neutral substances, and do
not include ions such as NH4+ (ammonium ion), HSO4- (acid
sulphate), and HCO3- (bicarbonate).
2. Limited to aqueous media.
• II- The Brønsted–Lowry Definition:
Acid D H+ + Base
ü An acid is a substance, either ion or molecule, that can
donate a proton (H+) to another substance. Substances
which are acids according to this concept are: HCl, HNO3,
H2SO4, H3O+, NH4+, HSO4-, HCO3- and H2O.
ü A base is a substance, also either ion or molecule, that is
capable of accepting a proton. The following are typical
bases: NH3, OH-, CH3COO-, CN-, HCO3- and H2O.
HCl is an acid because, in solution, it donates a
proton to water:
HCl(aq) + H2O(l )→H3O+ (aq) + Cl-(aq)
According NH3 is a base because it accepts a proton
from water:
to this
NH3(aq) + H2O↔NH+4(aq) + OH-(aq)
definition
some substances—such as water in the previous
two equations—can act as acids or bases
(amphoteric).
Conjugate Acid–Base Pairs
The substance that was the base (NH3) has become the acid (NH4 +) and vice versa. NH4 + and NH3
are called conjugate acid–base pair, two substances related to each other by the transfer of a
proton
• A conjugate acid is : any base to which a proton has been added.
• A conjugate base is: any acid from which a proton has been removed.
Summarizing the Brønsted–Lowry Definition of Acids and Bases:
■ A base accepts a proton and becomes a conjugate acid.
■ An acid donates a proton and becomes a conjugate base.
Auto ionization of water and pH
• Water is amphoteric; it can act as an acid or a base. Even when pure,
water acts as an acid and a base with itself, a process called auto
ionization:
• H2O + H2O ↔ H3O+ + OH-
Acid Base
(proton (proton
donor) acceptor)
So water ionization equation is : 2H2O ↔ H3O+ + OH-
For simplicity, it can be written as: H2O ↔ H+ + OH-
and the ionization constant can be expressed as:
• In dilute solutions concentration of water is a constant, and we may combine
[H2O] with constant Kc,
Thus Kc [H2O] = [H+] [OH-]
Kc [H2O] is called ion product of water or water dissociation constant, and is
given the symbol Kw.
Kw = [H+] [OH-] = 1.0 x 10-14
• In pure water, [H+] = [OH-] = x (molar conc.)
• Kw = [H+] [OH] = X2 = 1.0 x 10-14
x = 1.0 x 10-7 M
In neutral solution and in pure water [H+] = [OH-] = 1.0 x 10-7 M at 25oC
● In an acid medium the [H+] is > 1.0 x 10-7 > [OH-]
● In an alkaline medium the [OH-] is > 1.0 x 10-7 > [H+]
The pH Concept and Scale for Acidity and Basicity. Acidity of a solution is expressed in terms
of its hydrogen ion concentration, and not in terms of concentration of the dissolved acid.
The pH Concept:●Concentration of H+ in a solution may be expressed in terms of the
pH scale.
In general p Anything = - log Anything
●pH of a solution is defined as
pH = log1/ [H+]= -log [H+]
for example In a solution of [H+] = 10-3 M
pH = - log [H+] = - log [10-3] = - (-3) = 3
●For a neutral solution [H+] = 1 x 10-7
[OH-] = 1 x 10-7
\pH = - log [H+]= -log 10-7 = 7
pOH = -log [OH-]= - log 10-7 = 7
16 therefore, pH + pOH = 14
The pH Concept:
Solution
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pH meters are used to measure acidity
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Problem on pH
●Example 1: Calculate the pH of HNO3 whose hydrogen ion concentration is
0.76 M
Solution: pH = -log [ 0.76 ] = - ( – 0.119) = 0.119 ~0.12M
● Example 2: In a NaOH solution [OH-] is 2.93 x 10-4 M. Calculate the pH of
the solution.
Solution: pOH = -log [ 2.93 x 10-4 ] = 3.54
pH + pOH = 14
pH = 14- 3.54 = 10.46
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Problem on pH
●If pH = 1.46 and you are asked to calculate [H+].
pH = -log [H+]
1.46 = -log [H+]
We must take the antilog
So [H+ ] = 10 -1.46 =0.035
●Example 3: pH of rainwater collected in a certain region was 4.82. Calculate the
H+ ion concentration of the rainwater.
Solution:
4.82 = -log [H+]
We must take the antilog
So [H+] = 10 – 4.82 = 0.000015 = 1.5 x 10-5
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pH scale
[H+] 10-1 10-2 10-3 10-4 10-5 10-6 10-7 10-8 10-9 10-10 10-11 10-12 10-13 10-14
pH 1 2 3 4 5 6 7 8 9 10 11 12 13 14
Increasing acidity Increasing basicity
Neutral
v Many common substances are either acidic or basic, and their degree of acidity or basicity is
conveniently expressed in terms of pH .
v Amino acids are organic acids that contain NH2 group and acid COOH group.
v Neutral amino acids (monoamino and monocarboxylic; e.g glycine, and alanine), Acidic amino
acids (monoamino and dicarboxylic; e.g glutamic acid and aspartic acid), and basic amino acids
(diamino and monocarboxylic; e.g. ornithine and lysine).
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Strength of Acids and Bases
• Strong acids : are strong electrolytes that ionize completely in
water.
• Strong acids are inorganic acids:
hydrochloric acid (HCl), nitric acid (HNO3), perchloric acid
(HClO4), sulfuric acid (H2SO4), hydrobromic acid (HBr) and
hydroiodic acid (HI).
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• Weak acids: ionize only to a limited extent in water. Examples
of weak acids are hydrofluoric acid (HF), acetic acid
(CH3COOH), and the ammonium ion (NH4+).
HF(aq) + H2O(l) ↔ H3O+(aq) + F-(aq)
Strong acid as HCl is completely or 100 % ionized
Weak acid as Hf is weakly ionized
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• Strong bases: are strong electrolytes that ionize completely in
water: Hydroxides of alkali metals and certain alkaline earth
metals are strong bases.
• Weak bases, like weak acids, are weak electrolytes that ionize partially
in water. Ammonia is a weak base. It ionizes to a very limited extent in
water:
NH3(aq) + H2O(l) ↔ NH4 +(aq) + OH-(aq)
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Dissociation of weak acids
• Consider a weak monoprotic acid, HA.
• Its ionization in water is represented by:
HA(aq) + H2O(l) ↔ H3O+(aq) + A-(aq)
or simply HA(aq) + H2O(l) ↔ H+(aq) + A-(aq) e.g. acetic acid (CH3COOH):
CH3COOH ↔ H+ + CH3COO-
The equilibrium expression for this ionization is:
Ka=[H+][CH3COO-]/[CH3COOH]
where Ka, the acid ionization constant (the equilibrium constant for the ionization of an
acid).
• The strength of the acid HA is measured quantitatively by the magnitude of Ka:
- The larger Ka, the stronger the acid—that is, the greater the concentration of H+ ions
at equilibrium due to its ionization.
- Note that: only weak acids have Ka values associated with them.
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Dissociation of weak bases
• The ionization of weak bases is treated in the same way as the
ionization of weak acids:
• When ammonia dissolves in water:
NH3 + H2O ↔ NH+4 + OH-
So K = [NH+4 ][OH-]/[NH3][H2O]
Compared with the total concentration of water, very few water
molecules are consumed by this reaction, so we can treat [H2O] as a
constant.
Thus Kb = K[H2O] =
The
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smaller the value of Kb, the smaller the extent of ionization and the weaker the
base.
Dissociation of weak acids and bases:
Relative strengths of acids and bases can also be indicated by their pKa's.
●The smaller the value of pKa or pKb the stronger is the acid or base.
● For example, the pKa's for:
HC2H3O2 (acetic acid) pKa = 4.74
HC2H2ClO2 (chloroacetic acid) pKa = 2.85
HC2HCl2O2 (dichloroacetic acid) pKa = 1.30
● Order of increasing acidity is therefore,
dichloroacetic acid ˃ chloroacetic ˃ acetic
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