PHARMACEUTICAL ANALYSIS A
Lecture 1
Analyte is the compound to be analyzed in the chemical analysis. Analyte is present
in the sample.
Precipitation: neerslag
Classic wet chemistry: gravimetry, titrimetry
Gravimetry is a classic, wet-chemical quantitative method used to determine the
concentration of the analyte in a sample based on its mass. In this method, a known
amount of a reagent is added to the sample to form a precipitate/insoluble solid,
which is then rinsed, dried, and weighed. The original concentration of the analyte
can then be calculated.
Example: Ag+ + Cl- ⇄ AgCl(s) → used for Cl- analyte analysis
Titrimetry is another classic, wet-chemical quantitative method
used to determine the unknown amount of analyte A. The amount
of analyte in the sample is determined by measuring the
consumption of a reagent T that reacts with A.
A + T → AT
A = analyte (in sample S)
T = reagent that reacts with A (the reagent is contained in a
solution called titrant)
Analyte A is reacted/converted by addition of small, precisely
measured volumes of titrant containing a known concentration of reactant T.
Equivalence point: the amount of reagent T (and thus titrant) added is exactly the
same as the amount required to convert all of the analyte A to product AT. The
equivalence point is an important concept in titration. The point at which the
equivalence point is reached can be determined using an indicator, which is a
substance that changes color in response to changes in pH, or by using a pH meter
to measure the pH directly.
When a reaction has reached equilibrium, the forward and reverse rate of the
reactions are equal. In other words, the concentrations of A, T and AT remains the
same.
Choice of titration reaction is crucial:
- Reaction must go to completion – All titrant (T) added must react with available
analyte (A)
- Reaction must be fast
1
,The equilibrium constant, K:
ABCD: concentration of
chemical compounds
(reactants and products).
abcd: indicate how many molecules of a given compound take part in the reaction.
They are also called stoichiometry coefficients.
Stoichiometry: ratio of substances participating in a chemical reaction.
Ag(I)2 → Ag2+ + 2 I-
In this case is the stoichiometry Ag : I = 1 : 2
if a reacting compound is a gas, concentration will be given as PA instead of [A] in
units of pressure.
The value of K indicates the relative amounts of reactants and products present at
equilibrium and depends on the temperature. A large value of K (>1) indicates that
the products are favored at equilibrium, while a small value of K (<1) indicates that
the reactants are favored. If K is equal to 1, the reaction is at equilibrium with roughly
equal amounts of reactants and products.
Equilibrium constant of multiple reactions:
1st law of equilibrium: Law Of Mass Action
2nd law of equilibrium: The Equilibrium Law or Le
Châtelier’s Principle
Law of mass action: at a given temperature, a
chemical system reaches a state in which a
particular ratio of reactant and product
concentration has a constant value.
The equilibrium law or le Châtelier’s principle: when a system at equilibrium is
subjected to a change that disturbs it, it will re-adjust itself to partly counteract the
effect of the applied change so it can proceed back to equilibrium.
K always remains constant at a constant temperature!
Reaction proceeds to the right, when:
- Product is removed
- More reactant is added
Reaction proceeds to the left, when:
- Reactant is removed
- More product is added
Until equilibrium is reached again
2
,Don’t need to add H2O in the formula of K.
An example of the equilibrium law is the
addition of dichromate. So there is an
increased value in the numerator, so there
must also be an increase in the denominator
to re-establish the constant K value. So the
reaction system re-adjusts to the LEFT. The increase in product concentration is now
counteracted.
The reaction quotient (Q) is a measure of the relative concentrations of the products
and reactants in a chemical reaction at a given point in time. So you take the formula
of K and then fill in the concentration values. By comparing the value of Q to the
equilibrium constant (K), one can determine whether a reaction is at equilibrium or if it
will proceed in the forward or reverse direction to reach equilibrium.
- Q = K – the reaction is at equilibrium
- Q > K – the reaction will proceed to the left until equilibrium is re-established
- Q < K – the reaction will proceed the right until equilibrium is re-established
If we take the last example, we saw
that the concentration of dichromate
is increased: Le Châtelier - [Cr2O72-]
is increased from 0.10 M to 0.20 M.
If we calculate the reaction quotient Q, it can be determined that Q > K after the
increase. It can now be concluded that the reaction will proceed to the left. Product
concentrations will go down (numerator decreases) and reagent concentrations will
go up (denominator increases) to re-establish K.
Heat can also be regarded as a reagent or product. The temperature, and thus the
heat, can change the value of K.
Endothermal reaction: heat + reagent ⇋ product
- Increase in temperature, will lead the reaction to the right. K gets bigger!
Exothermal reaction: reagent 1 + reagent 2 ⇋ product + heat
- Increase in temperature, will lead the reaction to the left. K gets smaller!
What kinds of equilibria are good for titrimetric analysis?
- Solubility reactions
- Complex formation
- Acid-base chemistry
3
, 1. Solubility reactions
A sparingly soluble salt has a bad solubility.
- SP: solubility product (KSP)
AB2 (S) ⇋ A2+ (aq) + 2 B- (aq)
In solubility reactions, the product will be equal to 1!
KSP is used when an excess of a sparingly soluble salt is immersed in water. Both
reactants and products are then available. The salt dissolves until the concentrations
of cations and anions have reached their maximum values → the solution is
saturated. Under these conditions, the cation and anion concentrations multiplied
together equal the KSP.
Molar solubility: mol salt dissolved per L of aqueous solution
EXAMPLE: calculate the molar
solubility of barium iodate, Ba(IO3)2.
This dissolution reaction has a stoichiometry of
- 1 molecule salt: 1 metal cation: 2 iodate anions
So: molar solubility = mol salt
dissolved/L = [metal cation]
COMMON ION EFFECT:
HgCl2 (s) ⇋ Hg2+ (aq) + 2 Cl- (aq)
KSP = [Hg2+] [Cl-]2 = 1.2 x 10-18
→ [Hg2+] = 6.7 x 10-7 M & [Cl-] = 1.3 x 10-6 M
When KCl is added, the [Cl-] increases and [Hg2+] thus must decrease to satisfy the
equilibrium conditions (KSP). The amount of dissolved HgCl2 decreases, so the
solubility of HgCl2 also decreases and HgCl2 then precipitates.
In chemistry, precipitation occurs when two aqueous solutions containing ions are
mixed, and one or more insoluble ionic compounds form and precipitate out of
solution. This process is known as precipitation reaction.
4
Lecture 1
Analyte is the compound to be analyzed in the chemical analysis. Analyte is present
in the sample.
Precipitation: neerslag
Classic wet chemistry: gravimetry, titrimetry
Gravimetry is a classic, wet-chemical quantitative method used to determine the
concentration of the analyte in a sample based on its mass. In this method, a known
amount of a reagent is added to the sample to form a precipitate/insoluble solid,
which is then rinsed, dried, and weighed. The original concentration of the analyte
can then be calculated.
Example: Ag+ + Cl- ⇄ AgCl(s) → used for Cl- analyte analysis
Titrimetry is another classic, wet-chemical quantitative method
used to determine the unknown amount of analyte A. The amount
of analyte in the sample is determined by measuring the
consumption of a reagent T that reacts with A.
A + T → AT
A = analyte (in sample S)
T = reagent that reacts with A (the reagent is contained in a
solution called titrant)
Analyte A is reacted/converted by addition of small, precisely
measured volumes of titrant containing a known concentration of reactant T.
Equivalence point: the amount of reagent T (and thus titrant) added is exactly the
same as the amount required to convert all of the analyte A to product AT. The
equivalence point is an important concept in titration. The point at which the
equivalence point is reached can be determined using an indicator, which is a
substance that changes color in response to changes in pH, or by using a pH meter
to measure the pH directly.
When a reaction has reached equilibrium, the forward and reverse rate of the
reactions are equal. In other words, the concentrations of A, T and AT remains the
same.
Choice of titration reaction is crucial:
- Reaction must go to completion – All titrant (T) added must react with available
analyte (A)
- Reaction must be fast
1
,The equilibrium constant, K:
ABCD: concentration of
chemical compounds
(reactants and products).
abcd: indicate how many molecules of a given compound take part in the reaction.
They are also called stoichiometry coefficients.
Stoichiometry: ratio of substances participating in a chemical reaction.
Ag(I)2 → Ag2+ + 2 I-
In this case is the stoichiometry Ag : I = 1 : 2
if a reacting compound is a gas, concentration will be given as PA instead of [A] in
units of pressure.
The value of K indicates the relative amounts of reactants and products present at
equilibrium and depends on the temperature. A large value of K (>1) indicates that
the products are favored at equilibrium, while a small value of K (<1) indicates that
the reactants are favored. If K is equal to 1, the reaction is at equilibrium with roughly
equal amounts of reactants and products.
Equilibrium constant of multiple reactions:
1st law of equilibrium: Law Of Mass Action
2nd law of equilibrium: The Equilibrium Law or Le
Châtelier’s Principle
Law of mass action: at a given temperature, a
chemical system reaches a state in which a
particular ratio of reactant and product
concentration has a constant value.
The equilibrium law or le Châtelier’s principle: when a system at equilibrium is
subjected to a change that disturbs it, it will re-adjust itself to partly counteract the
effect of the applied change so it can proceed back to equilibrium.
K always remains constant at a constant temperature!
Reaction proceeds to the right, when:
- Product is removed
- More reactant is added
Reaction proceeds to the left, when:
- Reactant is removed
- More product is added
Until equilibrium is reached again
2
,Don’t need to add H2O in the formula of K.
An example of the equilibrium law is the
addition of dichromate. So there is an
increased value in the numerator, so there
must also be an increase in the denominator
to re-establish the constant K value. So the
reaction system re-adjusts to the LEFT. The increase in product concentration is now
counteracted.
The reaction quotient (Q) is a measure of the relative concentrations of the products
and reactants in a chemical reaction at a given point in time. So you take the formula
of K and then fill in the concentration values. By comparing the value of Q to the
equilibrium constant (K), one can determine whether a reaction is at equilibrium or if it
will proceed in the forward or reverse direction to reach equilibrium.
- Q = K – the reaction is at equilibrium
- Q > K – the reaction will proceed to the left until equilibrium is re-established
- Q < K – the reaction will proceed the right until equilibrium is re-established
If we take the last example, we saw
that the concentration of dichromate
is increased: Le Châtelier - [Cr2O72-]
is increased from 0.10 M to 0.20 M.
If we calculate the reaction quotient Q, it can be determined that Q > K after the
increase. It can now be concluded that the reaction will proceed to the left. Product
concentrations will go down (numerator decreases) and reagent concentrations will
go up (denominator increases) to re-establish K.
Heat can also be regarded as a reagent or product. The temperature, and thus the
heat, can change the value of K.
Endothermal reaction: heat + reagent ⇋ product
- Increase in temperature, will lead the reaction to the right. K gets bigger!
Exothermal reaction: reagent 1 + reagent 2 ⇋ product + heat
- Increase in temperature, will lead the reaction to the left. K gets smaller!
What kinds of equilibria are good for titrimetric analysis?
- Solubility reactions
- Complex formation
- Acid-base chemistry
3
, 1. Solubility reactions
A sparingly soluble salt has a bad solubility.
- SP: solubility product (KSP)
AB2 (S) ⇋ A2+ (aq) + 2 B- (aq)
In solubility reactions, the product will be equal to 1!
KSP is used when an excess of a sparingly soluble salt is immersed in water. Both
reactants and products are then available. The salt dissolves until the concentrations
of cations and anions have reached their maximum values → the solution is
saturated. Under these conditions, the cation and anion concentrations multiplied
together equal the KSP.
Molar solubility: mol salt dissolved per L of aqueous solution
EXAMPLE: calculate the molar
solubility of barium iodate, Ba(IO3)2.
This dissolution reaction has a stoichiometry of
- 1 molecule salt: 1 metal cation: 2 iodate anions
So: molar solubility = mol salt
dissolved/L = [metal cation]
COMMON ION EFFECT:
HgCl2 (s) ⇋ Hg2+ (aq) + 2 Cl- (aq)
KSP = [Hg2+] [Cl-]2 = 1.2 x 10-18
→ [Hg2+] = 6.7 x 10-7 M & [Cl-] = 1.3 x 10-6 M
When KCl is added, the [Cl-] increases and [Hg2+] thus must decrease to satisfy the
equilibrium conditions (KSP). The amount of dissolved HgCl2 decreases, so the
solubility of HgCl2 also decreases and HgCl2 then precipitates.
In chemistry, precipitation occurs when two aqueous solutions containing ions are
mixed, and one or more insoluble ionic compounds form and precipitate out of
solution. This process is known as precipitation reaction.
4
