Analytical Errors and Validation of Analytical procedures.pdf

1.3. Analytical Errors and Validation of
Analytical procedures
1
By: zenebe k

1. Errors in pharmaceutical analysis
• The terminology ‘error’ invariably refers to the difference in the
numerical values between a measured value and the true value.
 Simply accepting the analytical result could lead to rejection or
acceptance of a product on the bases of faulty analysis.
 For this reason it is usual to make several repeated measurements
of the same sample in order to determine the degree of agreement
between them.
• Consequently, the differences thus obtained between the standard
values and those by the new analytical methods are then treated as
‘errors’ in the latest procedure
2

Errors in pharmaceutical…
• Analytical errors may be broadly categorized into two heads, namely :
(i) Determinate (systematic) Errors, and
(ii) Indeterminate (random) Errors, (iii) growth error(sometimes)
1) Determinate (systematic) errors
• These are errors that possess a definite value together with a reasonable
assignable cause;
• however, in principle these avoidable errors may be measured and
accounted for conveniently. 3

 The most important errors belonging to systemic error are:
(a) Personal Errors : They are exclusively caused due to ‘personal equation’ of an analyst and
have no bearing whatsoever either on the prescribed procedure or methodology involved.
(b) Instrumental Errors : These are invariably(never changed) caused due to faulty and
uncalibrated instruments(measuring device), such as : pH meters, single pan electric balances,
UV spectrophotometers,
(c) Reagent Errors : The errors that are solely introduced by virtue of the individual
reagents(substance used to produce a chemical rxn), for instance : impurities inherently
present in reagents ; unwanted introduction of ‘foreign substances’ caused by the action of
reagents on either porcelain or glass apparatus.
4

(d) Constant Errors : They are observed to be rather independent of the
magnitude of the measured amount ; and turn out to be relatively less
significant as the magnitude enhances.
• Example : Assuming a constant equivalence point error of 0.10 ml is
introduced in a series of titrations, hence for a specific titration
needing only 10.0 ml of titrant shall represent a relative error of 1%
and only 0.2% for a corresponding 50 ml of titrant consumed

(f) Errors due to Methodology : Both improper (incorrect) sampling and
incompleteness of a reaction often lead to serious errors.
• A few typical examples invariably encountered in titrimetric and gravimetric
analysis are cited below :
6

2. Indeterminate (random) errors
• As the name suggests, indeterminate errors cannot be pin-pointed to any specific well-
defined reasons.
• They are usually manifested due to the minute variations which take place inadvertently
in several successive measurements performed by the same analyst, using utmost care,
under almost identical experimental parameters.
• These errors are mostly random in nature and ultimately give rise to high as well as low
results with equal probability.
• They can neither be corrected nor eliminated, and therefore, form the ‘ultimate
limitation’ on the specific measurements.
7

• Example : Figure below, represents the absolute errors in nitrogen
analysis by means of micro Kjeldahl’s Method, Here, each vertical line
labelled ( x1 – xt) designates the absolute deviation of the mean of the
set from the true value.
 In A represents ( x1 – xt) the absolute error obtained by ‘analyst-1’ for
the assay of benzyl-iso-thioureahydrochloride, whereas B represents (
x2 – xt) the absolute error obtained by ‘analyst-2’ for the assay of the
same compound.

9
 Larger indeterminate errors seem to be linked with the performance of
‘analyst 2’ than with that of ‘analyst-1’

Minimizing systematic errors
• Systematic errors may be reduced substantially and significantly by adopting
one of the following procedures rigidly, such as :
(i) Calibration of Instruments, Apparatus and Applying Necessary
Corrections
• Most of the instruments, commonly used in an analytical laboratory, such as :
UV-Spectrophotometer, IR-Spectrophotometer, single—pan electric balance,
pH-meter and the like must be calibrated duly, before use so as to eliminate any
possible errors. 10

• In the same manner all apparatus(equipments for particular activity), namely :
pipettes, burettes, volumetric flasks, thermometers, weights etc., must be
calibrated duly(as required or expected), and the necessary corrections
incorporated to the original measurements.
• In some specific instances where an error just cannot be avoided it may be
convenient to enforce an appropriate correction for the effect that it ultimately
causes ;
– for instance : the inherent impurity present in a weighed precipitate can be
estimated first and then deducted duly from its weight.

(ii) Performing a parallel control determination
• It essentially comprises of performing an altogether separate estimation under
almost identical experimental parameters with a quantity of a standard
substance that consists of exactly the same weight of the component as is
present in the unknown sample.
• Thus, the weight of the component present in the unknown sample may be
calculated with the help of the following expression :
12
where, X = Weight of the component present in the Unknown Sample.

(iii) Blank Determination :
• In order to ascertain the effect of the impurities present in the
reagents employed and reaction vessels used ; a blank
determination is an absolute necessity.
• It may be accomplished by performing a separate parallel
estimation, without using the sample at all, and under identical
experimental parmeters as employed in the actual analysis of the
given sample. 13

(iv) Method of Standard Addition
• Here, a small known quantity of the component under estimation is added to
the sample, which is subsequently subjected to analysis for the total amount of
component present.
• The actual difference in the quantity of components present in samples with or
without the added component ultimately gives the recovery of the quantum
added component.
• Note : The method of ‘standard addition’ is particularly useful to
physicochemical techniques of analysis, for instance : spectrophotometry,
turbidimetry.

2. Validation of Analytical procedures

2. Validation of Analytical procedures
• Validation is the conformation of a newly developed analytical methods that
is going to be used in QC to produce analytical results as designed and
expected.
• The object of validation of an analytical procedure is to demonstrate that it is
suitable for its intended purpose” determined by means of well-documented
experimental studies.
• Accuracy and reliability of the analytical results is crucial for ensuring
quality, safety and efficacy of pharmaceuticals. For this reason, regulatory
requirements have been published for many years. 16

Validation of Analytical……
General recommendation in method validation
• The principle of the test procedure should be described briefly
• Procedure should be sufficiently detailed to make repletion by
experts or other analyst and the following description should
be provided
Parameters evaluated or tested
Reagents preparation techniques
Calculation technique of assessed parameters
Equipment and parameters
Reference standard used
Precaution to be taken 17

• The common parameters(a limit defining the scope of a process or activity) that
should be verified in method validation are:
a) Linearity
• The ICH defines the linearity of an analytical procedure as the ability (within a given
range) to obtain test results of variable data which are directly proportional to the
concentration (amount of analyte) in the sample.
• The equation of a straight line takes the form: y = ax + b, wher b intercept of y-axis, a
slope of the line.
• At least five concentration levels should be used. Under normal circumstances,
linearity is achieved when the coefficient of determination (r2) is ≥0.997
18

b) Range
• The range of a method is related to its sensitivity, although there are methods
such as immunoassay w/c are capable of measuring very small amount of
materials, but are not very sensitive in that they measure over a restrict range
of low conc.
• Thus, some typed of detection have very wide dynamic range and other may
only function over a restriction range before linearity is lost. E.g. A UV
detector has a dynamic range of about 1x103 and for a particular cpd it might
measure conc b/n 0.1-100μg/ml.
19

c) Accuracy
• The ICH defines the accuracy of an analytical procedure as the closeness of
agreement between the values that are accepted either as conventional true
values or an accepted reference value and the value found.
• Accuracy is usually reported as percent recovery by assay, using the proposed
analytical procedure, of known amount of analyte added to the sample
• Typical accuracy of the recovery of the drug substance in the mixture is
expected to be about 98 to 102%.
• Values of accuracy of the recovery data beyond this range need to be
investigated.
20

c) Precision
• The precision of an analytical procedure expresses the closeness of agreement
(degree of scatter) between a series of measurements obtained from multiple
samples of the same homogeneous sample under prescribed conditions.
• it does not imply anything with respect to their relation to the ‘true value’
• Precision is usually investigated at three levels: repeatability, intermediate
precision, and reproducibility
• Example : A sample of pure Peppermint Oil is known to contain 31.10 ± 0.03 per cent of
Menthone. The results obtained by two Analysts-1 and 2 are as stated below :
21

Re
d) Repeatability
• is a measure of the precision under the same operating conditions
over a short interval of time.
• It is sometimes referred to as intraassay precision
The arithmetic mean stands at 31.12%
The arithmetic mean is 31.44%

e) Intermediate Precision.
• Intermediate precision is defined as the variation within the same laboratory.
• The extent to which intermediate precision needs to be established depends on the circumstances
under which the procedure is intended to be used.
• Typical parameters that are investigated include day-to-day variation, analyst variation, and
equipment variation
f) Reproducibility.
• Reproducibility measures the precision between laboratories
• This parameter should be considered in the standardization of an analytical procedure (e.g.,
inclusion of procedures in pharmacopoeias and method transfer between different laboratories).
23

 To validate this characteristic, similar studies need to be performed at
other laboratories using the same homogeneous sample lot and the same
experimental design.
 the most common approach is the direct method transfer from the
originating laboratory to the receiving laboratory.

g) Selectivity
• The selectivity of a method is a measure of how capable it is of measuring the analyte alone in the
presence of other copds contained in the sample.
• The most selected analytical methods involves a chromatography separation.
• Detection methods can be ranked according to their selectivity
• A simple comparison is b/n fluorescence and UV spectrophotometery; there are many more cpds
w/c exhibit UV absorption than fluorescence, thus fluorescence spectrophotometry is moer
selective methods.
• B/c selective methods are based on more complex principles than non selective methods they may
be less robust, e.g. fluorescence spectrophotometry is more affected by changes in the analytical
method than UV.
25

h) Robustness
• Refers to how resistance the precision and accuracy of an assay is to small variation in
the method. E.g. changes of instrumentation, slight variation in extraction procedure etc.
• Robust assays may not be capable of the highest precision or specificity but they are
regarded as fit for the purpose for w/c they are designed.
i) Sensitivity
• It indicates how responsive the method is to a small change in the conc of the analyte.
• It can be viewed as the slope on a response curve and may be a function of the method it
self or of the way in w/c the instrument in has been calibrated
26

1.4. Basic calculations in
pharmaceutical analysis
27

• Molality(M):
• It is the number of moles of solute in 1000g
of the solvent.
Molality (m) =
Number of moles of solute
Mass of solvent(kg)
MOLES=
𝒈𝒊𝒗𝒆𝒏 𝒎𝒂𝒔𝒔
𝒎𝒐𝒍𝒆𝒄𝒖𝒍𝒂𝒓 𝒘𝒆𝒊𝒈𝒉𝒕

1.4. Basic calculations……
Molarity
 The molar concentration (Cx) of the solution of the chemical species X is the number of
moles of that species that is contained in one liter of the solution (not one liter of the
solvent).
 The unit of molar concentration is molarity, M, which has the dimensions of mol L-1.
Cx= no mole solute = no m mole solute
no L solution no ml solution
 One liter of one molar solution will consist of one mole of solute plus enough solvent to
make a final volume of one liter.
Example:
 Calculate the molar concentration of ethanol in an aqueous solution that contains 2.30g of
C2H5OH (46.07 g/mol) in 3.50L of solution.
 Describe the preparation of 2.00L of 0.108M Bacl2 from BaCl2.2 H2O (FW= 244.3g /mol )
28

Normality
• Normality (N) is defined as the number of equivalents weight of solute
dissolved in one liter of solution.
N = no of equivalent weight no of equivalent weight = weight of solute
Volume of so/n in Liter equivalent weight
• An equivalent weight is defined as the ratio of a chemical species’ formula
weight (FW) to the number of its equivalents
• The number of equivalents, n, is based on a reaction unit, which is that
part of a chemical species involved in a reaction
• Normality makes use of the chemical equivalent, which is the amount of
one chemical species reacting stiochiometrically with another chemical
species. 29

• Note that this definition makes an equivalent, and thus normality, a
function of the chemical reaction in which the species participates.
• Although a solution of H2SO4 has a fixed molarity, its normality
depends on how it reacts.
• In an acid–base reaction, the reaction unit is the number of H+ ions
donated by an acid or accepted by a base. For the reaction between
sulfuric acid and ammonia
n = 2 for H2SO4 and n = 1 for NH3

• In a precipitation reaction, for example, the reaction unit is the charge of
the cation or anion involved in the reaction; thus for the reaction
n = 2 for Pb2+ and n = 1 for I–.
• For a complexation reaction, the reaction unit is the number of electron
pairs that can be accepted by the metal or donated by the ligand. In the
reaction between Ag+ and NH3
n for Ag+ is 2 and that for NH3 is 1
31

• Finally, in an oxidation–reduction reaction the reaction unit is the number
of electrons released by the reducing agent or accepted by the oxidizing
agent; thus, for the reaction
n = 1 for Fe3+ and n = 2 for Sn2+
Example:
 The equivalent weight of H2 SO4 is (FW=98 gm/mol)
 If there is a one liter solution that contains 78.32 grams H2 SO4 , the
number of equivalents is=---------------? answer=n=0.8

Weight, Volume, and Weight-to-Volume Ratios
• Weight percent (% w/w), volume percent (% v/v) and weight-to-volume percent
(% w/v) express concentration as units of solute per 100 units of sample
• Percent composition of a solution can be expressed as:
• Weight percent (w/w) = mass of solute X 100%
mass of soln
• Volume percent (v/v) = volume of solute X 100%
volume of solution
• Weight /volume percent (w/v) = mass of solute g X 100%
• volume soln ml
33
2.Percentage (% concentration Expression)

• Weight percent is frequently employed to express the concentration of commercial
aqueous reagents. E.g. 37% hydrochloric solution – this means the reagent contains
37g of HCl per 100g of solution.
• Volume percent is commonly used to specify the concentration of a solution prepared
by diluting a pure liquid with another liquid.
E.g. 5% aqueous solution of methanol – usually means a solution prepared by diluting
5.0ml of pure methanol to give 100ml of solution with enough water.
• Weight /volume percent is after employed to indicate the composition of dilute
aqueous solutions of solid reagents.
E.g., 5% aqueous silver nitrate often refers to a solution prepared by dissolving 5g of
silver nitrate in sufficient water to give 100ml of solution
34

Parts per million & parts per billion
 For very dilute solutions, parts per million (PPM) is convenient way to express
concentration:
Cppm = Mass of solute X 106 ppm
Mass of solution
 The units of mass in the numerator & denominator must agree.
 For even more dilute solutions we use parts per billion.
Cppb = Mass of solute X109 ppb
Mass of solution
 If we approximate the density of an aqueous solution as 1.00 g/mL, then solution
concentrations can be expressed in parts per million or parts per billion using the following
relationships.
 Example
 What is the morality of K+ in aqueous solution that contains 63.3 ppm of K3 Fe
(CN)6 (329.2 g/mol)? 35

1.5. Physical and chemical
properties of drug molecules
36

1.5. Physical and chemical……
• The physical properties of a drug molecules along with simple chemical derivatization and
degradation reaction play an important part in the development of analytical methods.
• Drug molecules can be complex, containing multiple functional group that in combination
produce the overall properties of the drug
Calculation of pH value of aqueous solution of strong and weak acids and
bases
• The pH of a solution is defined as – log [H+], where [H+] is the conc of hydrogen ion in
so/n
• In pure water the conc. of hydrogen ion governed by the equilibrium:
• Ka is the dissociation constant for the equilibrium, is known as Kw in the case of the
dissociation of water
•
37

• Since the conc. of water does not change appreciably as a result of ionization its conc. Can be
regarded as not having an effect on the equilibrium and it can be omitted from the equation and
this mean that in pure water:
• If an acid is introduced into an aqueous so/n the [H+] increase
• Strong acid is completely ionized in water and [H+] is equal to its Molarity e.g. 0.1M HCl contains
0.1M H+ and has a pH of log [0.1]=1
• For a so/n of a strong base such as 0.1NaOH, [OH-]=1M and
[H+] [OH-]=1x 10-14, therefore [H+]=1x10-13 and pH=13
• Weak acids are not completely ionized in aqueous so/n and are in equilibrium with the
undissociated acid, as is the case for water, w/c is a very weak acid.
38

• The dissociation constant Ka is given by the expression below:
• For instance in a 0.1M so/n of acetic acid (Ka=1.75 x 10-5) the equilibrium can be written
as follows:
• The pH can be calculated as follows:
• Since the dissociation of the acetic acid does not greatly change the conc. of
the unionized acid the above expression can be approximated to:
• In comparison the pH of 0.1M HCl is 1
39

40

Acidic and basic strength and pKa
• The pKa value of a cpd is defined as :pKa= - log Ka
• If pKa is used as a measure of acidic or basic strength, for an acid the smaller the pKa
vallue the strongest the acid. For a base the largest the pKa value the stronger the base
• For an acid the forward reaction used
• In the case of a base it is the protonated form of the base that act as a proton donor
Buffer solution
• A solution containing a weak acid/ base and its conjugate base/acid that is resistant to a
change in pH when a strong acid or strong base is added.
• Adding as little as 0.1 mL of concentrated HCl to a liter of H2O shifts the pH from 7.0 to
3.0. The same addition of HCl to a liter solution that is 0.1 M in both a weak acid and its
conjugate weak base, however, results in only a negligible change in pH
41

 A mixture of acetic acid and sodium acetate is one example of an acid/base buffer
 The equilibrium position of the buffer is governed by the reaction
 The relationship between the pH of an acid–base buffer and the relative amounts of
CH3COOH and CH3COO– is derived by taking the negative log of both sides of the
above equation and solving for the pH
 Buffering occurs because of the logarithmic relationship between pH and the ratio of
the weak base and weak acid concentrations.
 For example, if the equilibrium concentrations of CH3COOH and CH3COO– are equal,
the pH of the buffer is 4.76.
 If sufficient strong acid is added such that 10% of the acetate ion is converted to
acetic acid, the concentration ratio [CH3COO–]/[CH3COOH] changes to 0.818, and the
pH decreases to 4.67.
42

• A more useful relationship relates the buffer’s pH to the initial concentrations of weak
acid and weak base.
• A general buffer equation can be derived by considering the following reactions for a
weak acid, HA, and the salt of its conjugate weak base, NaA.
• After several rearrangement, it provide us a general formula known as henderson-
Hasselbalch equation
• Where CNaA conc of salt, CHA conc of the acid
• Hasselbalch equation provides a simple way to calculate the pH of a buffer and to
determine the change in pH upon adding a strong acid or strong base.
• For a base, henderson-Hasselbalch equation is written as
43

• Using hasselbach equation it is possible to determine degree of ionization of a
drug at a given pH. E.g. degree of ionization of acetic acid at pH of 4.76
acetic acid ionized 50% at pH of 4.76
EXAMPLE :
• Calculate the pH of a buffer that is 0.020 M in NH3 and 0.030M in NH4Cl. What is
the pH after adding 1.00 mL of 0.10 M NaOH to 0.10 L of this buffer? (Ka=5.7x10-10)
• Calculate the percentage of ionization of diphenhydramine at pH of 7
44

Analytical Errors and Validation of Analytical procedures.pdf

Stability of drugs
 Many drugs are quite stable but functional groups such as esters and lactam rings
w/c occur in some drugs are susceptible to hydrolysis and functional groups such as
catechols and phenols are quite readily oxidized.
 The most common type of degradation w/c occur and formulated drugs obey zero or
first order kinetics
Zero order degradation
 In zero order kinetics the rate of degradation is independent of the conc of the
reactant.
 Thus, if the rate constant for the zero order degradation of a subs is 0.01mole/h
then after 10hr 0.1 mole of the subs will have degraded
 This type of degradation is typically of hydrolysis of drugs in suspension or tablets
where the drugs is initially in the solid state and gradually dissolve at the same rate
as the drug in so/n id degraded
45

First order degradation
• This type of degradation would be typical of hydrolysis of a drug in so/n
• In first order kinetics the rate constant k has units h-1 or s-1 and the rate of the
reaction for a drug is governed by the expression
• From this expression by integration and rearrangement the following expression arises:
•
The half life of the drug (the time taken for 50% of a sample drug to degrade, i.e. where
x is a/2) is thus given by the following expression
46

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