n

1. INTRODUCTION TO
CLINICAL
CHEMISTRY
CHAPTER ONE

2. LEARNING OBJECTIVES
Upon completion of this chapter the student will be able to:
1. Define clinical chemistry
2. Explain the significance of clinical chemistry
3. Describe units of measurements
4. List common apparatus and equipments used in the clinical
chemistry laboratory

3. Outline of introduction to clinical chemistry
lecture
 Define clinical chemistry and other important terminologies
 Significance of clinical chemistry
 Measurement units
 Apparatuses and equipments used in the clinical chemistry
laboratories

4. DEFINITIONS
Clinical Chemistry
 is defined as an area in laboratory sciences that deals
with chemical analysis of body fluids such as blood,
urine, spinal fluid as well as feces, tissue, calculi and
other materials.
• Links the knowledge of general, organic, Inorganic
analytical & biochemistry with an understanding of
human physiology.

5. DEFINITIONS… CONT’D
 Analyte
 Analysis
 Types of Analysis
 Qualitative
 Semi-quantitative
 Quantitative

6. DEFINITIONS… CONT’D
 Analyzer
 Reagent
 Quality assurance

7. The main purpose of Clinical Chemistry tests
 To assess the physiological function of our body systems or
organs
 To diagnose and monitor diseases
 To follow up response to treatment

8. Significance of Clinical Chemistry
Provides biochemical testing of patient sample
 Glucose
 Protein
 Bilirubin
 Creatinine
 Lipids
 Enzymes
 Electrolytes
 And Other biochemicals

9. Clinical Chemistry Equipment and
Supplies

10. Analysis in Clinical Chemistry

11. Clinical Chemistry Instrument
Spectrophotometer

12. Typical Clinical Chemistry Analyzer

13. Making Reagents for Clinical Chemistry

14. Clinical Chemistry Reagents

15. Reagents for Testing in Clinical
Chemistry

16. SPECIMENS TESTED IN
CLINICAL CHEMISTRY

17. Clinical Chemistry Testing

18. Quality Control in Clinical Chemistry

19. Recording Clinical Chemistry
Results

20. Clinical Chemistry Analysis

21. UNIT OF MEASURMENT
Basic units: Describes one fundamental physical quantity

23. Derived unit I: is related mathematically
to the basic or supplemental units

24. Summary
 Definition of clinical chemistry and other terms used in
clinical chemistry
 Significance of clinical chemistry for laboratory diagnostics
 Different apparatuses and equipments used in the clinical
chemistry laboratories

25. CHAPTER TWO
Solutions

26. OBJECTIVE
At the end of the chapter, the student will be able to
• Define solution, types of solutions and expressing
concentration of solutions
• Discuss purpose and how to prepare dilutions
• Explain standard, primary, and secondary solutions
• Discuss concept, preparation, and calculations of pH, Strong
and weak acid solutions, and Strong and weak basic solutes.
• Discuss concept, and calculations of dissociation constants.

27. OUTLINE OF SOLUTION LECTURE
• Definition of solution
• Types of solutions
• Expressing concentration of solutions
• Chemical units: molarities, normality
• Dilution: Definition and preparation
• Standard, primary, and secondary solutions
• Concepts and characteristics of acids and bases
• pKa, dissociation constant of strong acids & bases, dissociation constant of
weak acids and bases
• pH- concept and calculations
• Henderson-Hasselbach equation
• Buffer solutions: Characteristics and classifications.

28. SOLUTION
2.1 Definition
• A solution is a homogeneous mixture of one or
more substances dispersed molecularly in a
sufficient quantity of dissolving medium
• A solution is made up of solute and solvent
• Solute + solvent = Solution

29. 2.2 TYPES OF SOLUTION
•Solid solution:
•Liquid solution:
•Gaseous solution:

30. Solid solution:
a mixture of two or more solids & they are dispersed or mixed randomly
throughout one another. E.g. alloys.
Liquid solution: in liquid solution the solvent is always liquid where as
the solute can be liquid, solid, or gas.
Gaseous solution: this is a homogenous mixture of two or more gases. Eg air
Note: If any type of solute (gas, liquid, or solid ) is dissolved in water, the
solution is said to be aqueous.
Many measurements in clinical chemistry laboratory concerns the
determination of dissolved solutes (analyte) in a solvent.

31. 2.3. EXPRESSING CONCENTRATION OF
SOLUTIONS
Strength of solutions can be categorized based on the relative amount of the
solute to the solvent.
Expression of the strength of solutions is broadly divided into two:
2.3.1 Relative expression
a. Dilute solution:
Solution which contains small amount of solute in large amount of solvent
b. Concentrated solution
Solution which contains large amount of solute in a small amount of
solvent.

32. c. Saturated solution
solution in which a given volume of solvent has dissolved all the
solutes at a given temperature & pressure.
If the temperature or pressure of the solution is altered, the solution
is no longer saturated.
Therefore, a saturated solution may contain excess solute
d. Super saturated solution
is a solution which holds more solute than it can hold normally at a
given temperature & pressure.

33. To prepare supersaturated solution
First, prepare a saturated solution. Heat the solution; and add more
solute, and dissolve completely to create unsaturated solution; Slowly
cool the solution to keep the dissolved solute in solution; While
cooling, excess solutes may crystallize out of the solution;
To check whether the solution is saturated, unsaturated or
supersaturated, add a small crystal of the solute.
If the crystal dissolves, the solution is unsaturated or saturated or
if the crystal is observed, the solution is supersaturated.

34. 2.3.2 QUANTITATIVE EXPRESSION OF
THE CONCENTRATION
The concentration of the solutes in a solution can be expressed quantitatively in
physical units or chemical units.
A. Physical units
– Parts per hundred (Percentage)
%W/W,
%W/V
% V/V
– Parts per unit
– Parts per million

35. Parts per hundred (Percentage)
%W/W This is the number of parts of solute by weight per 100 parts of solution by
weight.
Eg. 37 % w/w HCl means 100g of HCl solution contains 37g of HCl & the rest
63g is the solvent (water)
% W/V This is the number of parts of solute by weight per 100 parts of solution by
volume.
Eg. 98% w/v glucose solution means 98g of glucose was dissolved in 100 ml of
solution.
%V/V This is the number of parts of solute by volume per 100 parts of solution.
Eg. 70% v/v ethanol means 70ml of ethanol was mixed with 30ml of water.

36. b. Parts per unit
It is to express the number of parts by volume or by weight of solutes per given volume
or weight of the solution.
For example; prepare 1000 mL of a mixture of 1 part of acetic acid with 3 part of
ethanol.
Formula: V= C/A+B
V = 1000mL
1 part of acetic acid + 3 part of ethanol
= 1000mL
4 = 250 mL ---- one part of solute
c. Parts per million
Is the number of parts of solute by weight or volume per 1 million parts of solution by
weight or volume.

37. B. Chemical units
• Mole is the amount of a given substance in grams.
• terms like normality, molarity, or molality are used.
• If the quantity of the solute is very small milli-mole,
micromole or nanomoles may be used.
 1 mole = 1,000 milimoles
 1mole = 1,000,000 micromoles
 1mole = 1,000,000,000 nanomoles

38. A. MOLAR SOLUTION (MOLARITY, M)
• number of moles of solute per liter of solution or the gram
molecular weight of a compound per liter of solution.
E.g.. Prepare an aqueous glucose solution with two moles of
glucose per liter
• Solution: 1 mole= its molecular weight in grams
 mole of glucose = 180g
 2 moles = 2X180g = 360g
• 360g glucose is dissolved in 1 liter of water to yield 2 moles/l
of glucose.

39. Formula:
Molarity = Actual weight in grams
Molecular Weight x volume of solution in liter.
E.g.. What is the molarity of NaOH if 20g NaOH is dissolved in
200ml of solution?
• Solution: actual mass = 20g, molecular weight = 40
Volume = 0.2 liters
Molarity = 20g = 2.5 M NaOH
40g X0.2 liter

40. a. Normal solution (Normality, N)
• is the number of grams equivalents of solute by weights per liter of
solution (not solvent).
• An equivalent weight is equal to the molecular weight of a
substance divided by its valance.
• The valance is the number of unit that can combine with or replace
one mole of hydrogen ion
• Normality (N) = actual weight in grams
Equivalent weight x Vol. in liter

41. Eg. What is the normality of KOH if 5.6g KOH is
dissolved in 1 liter of solution?
• Solution: actual mass = 5.6g
Volume of solution= 1 liter
KOH equivalent weight = (56/1) = 56
• Therefore 5.6g = 0.1 N KOH
56 X 1 liter
• Exercise. How do you make 0.4 N solution of KOH.

42. 2.4.MAKING DILUTION
2.4.1.Introduction:
• Dilution is defined as a process by which the
concentration of a given solution is decreased by the
addition of solvents to make a weaker solution
• It represents the ratio of concentrated material to the total
volume of a solution
• Total volume consists of the volume or weight of the
concentrated plus the volume of the diluents

43. DILUTIONS
• This ratio of concentrated or stock solution to the total volume
equals the dilution factor
Dilution factor (df) = volume of stock
Total volume of solution
• The df is inversely related to the concentration thus, the
dilution factor increases as the concentration decreases.

44. REASONS FOR DILUTION
• To prepare a weaker solution from stronger one
• If the specimen at hand is less than a procedure calls
for
• If the concentration of substances (analyte) is too
high to be accurately measured.
• To remove undesirable substances from the specimen

45. 2.4.2.SIMPLE/SINGLE DILUTION
• Is defined as one part of the original solution to the total part
of final solution, which include both solute & solvent.
• In making a simple dilution, the laboratory technician must
decide on the total volume desired & amount of stock
solution to use

46. A. USING PROPORTION
• It is used when reagents are prepared by adding a specific
amount of one solution to a specific amount of another
solution.
V =
Where: C – total volume of final reagent
A – total parts of solution A
B – total parts of solution B
V – volume of each part

47. PROPORTIONS
• Example 1: a buffer made by adding two parts of ‘solution A’ to five
parts of solution B would be required to make 70 mL of the buffer.
• Formula: V = 70mL required
2 parts of A + 5 part of B
= 70 mL
7 part
= 10 volume of one part
A = 2 x 10 = 20 mL
B = 5 x 10 = 50 mL
C
A + B

48. EXERSISE
a 100mg/mL N2 standard is diluted 1:10. then the concentration
of the resulting solution is
100mg/mL x 1/10 = 10 mg/mL

49. B. USING C1V1 = C2V2
• This formula is useful only if the units for concentration
& volume are the same & if three of the four variables are
known.
• Example what volume is needed to make 500ml of 0.1M
solution of tris-buffer from a solution of 2 M tris-buffer?

50. 2.4.3. SERIAL DILUTION
• Large dilutions may be difficult to make because of
the amount of diluent that needs to be added.
• For example, a 1/1000 dilution may be difficult to
create accurately even with 0.1mL of serum & 99.9
mL of diluent.
• A series of dilutions, also called serial dilutions, may
be a better way to make the dilutions.
• Multiple progressive dilution process .

51. SERIAL DILUTION,
CONTINUED
• It is required for certain quantitative tests
• Serial dilution is extremely useful when the volume of the
concentrate &/or diluents is in short supply

52. 1 mL 1 mL 1 mL 1 mL -
discard
Sample: serum tube 1 tube 2 tube 3
Dilution: 1:2 1:4 1:8
Concentration: 1 0.5 0.25
1ml 1ml 1ml
1ml

53. DILUTION
The dilution fold of a system can be determined by the formula:
1 = volume transferred
dilution fold total volume
Volume transferred = is equal to the constant volume
transferred to each successive tubes in the serial dilution system.
Total volume = is equal to the volume being transferred plus the
volume of diluents already in the tube.

54. EXAMPLE
What is the dilution fold of the following serial dilution system
consisting of five tubes? The following amount of diluents have
been added to the tubes; 0.5 mL to tube 1 & 0.5 mL to tube 2 to
5. Next, 0.5 mL of patient serum is added to tube 1 and 0.5 mL
is serialy transferred through tube 5. finally, 0.5 mL is discarded
from tube 5.
• 1/Y = 0.5/1.0
• Y x 0.5 = 1
• Y = 1/0.5 = 2

55. DILUTION EXAMPLE 2
It is often desirable to determine the dilution of a given tube (Y) in a serial
dilution system. This dilution can be calculated by
Solution of tube 1 = dilution of Y x [ 1/dilution fold]y-1
• What is the dilution of tube 3 in the preceding example?
Y= ½ x (½) (Y-1)
= ½ x (½)2
= ½ x ½ x ½
= 1/8 ;
The dilution of serum in the tube 3 is 1/8.

56. 2.5.STANDARD SOLUTION
• Is a solution whose concentration is exactly known.
• There are two types of standard solutions:
– primary standard
–secondary standard solution.

57. 2.5.1.PRIMARY STANDARD SOLUTION
• a highly purified chemical that can be measured directly
to produce a substance of known concentration in a given
solution.
• The solution must satisfy the following criteria:
• Must be essentially free of impurities
• Should be stable in both solid & solution form
• The concentration should be accurate.
• Should not be hygroscopic or vaporize at 20oc

58. EXAMPLE
• Sodium carbonate(NaCO3)
• Sodium oxalate
• Sodium chloride,
• Potassium dichromate

59. 2.5.2 SECONDARY STANDARD SOLUTION
• is defined as a substance of lower purity & whose concentration is
determined by comparison to a primary standard.
• To prepare secondary standard
• Weigh some amount of the substance using analytical balance.
• Dissolve in a given volume of solvent.
• Using primary standard solution, determine the exact
concentration of the substance.
• If necessary dilute to some volume to get the concentration
required

60. EXAMPLE SECONDARY STANDARD
• Oxalic acid
• Nitric acid
• Hydrochloric acid
• Sulfuric acid.

61. 2.6 CONCEPT OF ACIDS AND BASES
A. Acids & Bases
• According to Bronsted Lowry theory and Arrhenus
theory respectively
Acids – are proton donors
- dissociate to furnish hydrogen ions
Bases – are proton acceptors
-dissociate to furnish hydroxyl ions in aqueous
solution

62. Characteristics of acids
-sharp taste (don’t try this) – refers to sourness of a solution
-change the color of blue litmus paper in to red
- ability to neutralize bases
-ability to trigger the evolution of carbon dioxide when added to
carbonate.
Characteristics of bases (alkaline solution)
soapy feeling – refers to bitterness of a solution
ability to change the color of red litmus paper to blue.
Ability to neutralize acids

63. DISSOCIATION CONSTANT
 HA + H2O H3O+ + A-
K = [H3O+ ] [A-]
[HA ] [H2O]
Ka = K [H2O] = [H3O+ ] [A-]
[HA ]
• For weak acids Ka < 1
• For strong acids Ka > 1

64. pKa of acids or bases
• Strength of an acid or base is determined by its pKa
which is defined as the negative logarithm of the
ionization constant.
 HA H+ + A-
 K = [H+] [A-]
[HA]
 pKa = -log Ka
• Strong acids has a low pka value where as a weak
acids has a high pKa value.

65. DISSOCIATION OF STRONG ACIDS OR
BASES
• Strong acids are those which dissociate completely into
their respective cations (H+) &anions in aqueous (A-)
solution.
 HA H+ + A-
 HNO3 H+ + NO3-
 H2SO4 2H+ + SO4-

66. • Strong bases are those which dissociate completely
into their respective hydroxyl ions (OH-) & respective
cations (B-) in aqueous solution.
• BOH B+ + OH-
• NaOH Na+ + OH-

67. DISSOCIATION OF WEAK ACIDS OR
BASES
• Weak acids (HA) are those which partially dissociate
when they are in aqueous solution
• HA H+ + A-
• Examples: Acetic acid, Carbonic acid, phosphoric
acids, lactic acids, benzoic acids etc.

68. • Weak bases are those which dissociate partially when
dissolved in aqueous solution.
 BOH B+ + OH-
 Eg NH3, Ca (OH)2, Al2(OH)3

69. HYDROGEN ION CONCENTRATION
(PH)
To express hydrogen ion concentration, the symbol pH was
introduced by Sorensen in 1909.
• pH is defined as the negative logarithm of the Hydrogen
ion concentration.
pH = - log [H+] in mol/L
pH of pure water = 7
pH of 1 M solution of strong acid = 0
pH of 1 M solution of strong base = 14

70. EXERSISE
1. What is the pH of a solution whose hydrogen ion
concentration is 3.2 x 10-4 mol/L
2. What is the pH of a solution whose hydroxide ion
concentration is 4 x 10-4
3. Calculate the hydrogen ion concentration of a
solution of HNO3 with pH 3.4.

71. HANDERSON-HASSELBALCH
EQUATION
• HA H+ + A-
• K = [H+] [A-]
[HA]
• [H+] = Ka x [HA]/ [A-]
• Log [H+] = log Ka + log [HA]/ [A-]
• - Log [H+] = -log Ka - log [HA]/ [A-]
• PH = PKa + log [A-]/ HA]

72. • The pH of the system is determined by the pKa of the
acid & the ratio of [A-] to [HA].
 pH = pKa + log101
 pH = pKa + 0
 pH = pKa
• The buffer has greatest buffer capacity at its pKa, i.e., that
at which the [A-] = [HA].

73. • The capacity of the buffer decreases as the ratio
deviates from 1.
• In general, a buffer should not be used at a PH > 1
from its PKa.
• What is the pH of a solution whose hydrogen ion concentration is 3.2 ×
10-4 mol/l?

74. BUFFER SOLUTION & ITS ACTION
• Buffers are chemical system that prevents change in the
concentration of another chemical substance.
• For example, a proton donor & acceptor system of solutions
are used as buffers to prevent change in hydrogen ion
concentration.
• It consists of a mixture of weak acid (or base) & its conjugate
base (or acid).
• E.g. CH3COOH(ethanoic acid) & CH3COONa (its salt)

75. • Buffers resists changes in pH when small quantities
of an acid or a base is added to it.
• A buffer act as like an base if acid is added & like an
acid if base is added.
• The useful buffer range is at PKa + 1

76. • A good buffer is prepared by mixing equal amount of the acid
(base) & its salt.
• The action of the buffer & their role in the maintenance of solutions
PH is explained with the aid of Henderson-Hasselbatch equation
HA H+ + A-
BOH B+ + OH-
HA - is a weak acid, indicate that H is combined with an other
chemical element or compound A.
A – is a conjugate base since it can accept H+ & act as a base

77. • Maximum buffering capacity is achieved when
[A-] = HA, such that the ratio of [A-]/ HA] is 1 & pH = pKa.
• A good buffer maintains its pH change within the limit of
+ 1 PH unit.
• Similarly a buffer solution is effective if the pH of
prepared buffer is adjusted between + 1 pKa or pKb.
• E.g. the pKa acetic acid at 25 0C is 4.76. to prepare
effective acetate bufer, the pH should be between 3.76 &
5.76.

78. TYPES OF BUFFERS
1.Neutral buffer – phosphate buffer
2. Acidic buffer – citrate buffer
3. Alkaline buffer – borate buffer

79. NEUTRAL BUFFERS
• It is a buffer solution that has pH value of around 7.
• Used for maintaining reaction around neutral pH

80. ACIDIC BUFFER SOLUTIONS
• It is a buffer solution which has a pH less than 7.
• They are commonly made from a weak acid & one of its
salts - often a sodium salt.
• For example, a mixture of ethanoic acid & sodium ethanoate
in solution.
• In this case, if the solution contained equal molar
concentrations of both the acid & the salt, it would have a
pH of 4.76 (its PKa).
• You can change the pH of the buffer solution by changing
the ratio of acid to salt, or by choosing a different acid & one
of its salts.

81. ALKALINE BUFFER SOLUTIONS
• It is a buffer solution which has a pH greater than 7.
• they are commonly made from a weak base & one of its
salts.
• For example, a mixture of ammonia & ammonium
chloride solutions.
• If they are mixed in equal molar proportions, the solution
would have a pH of 9.25.

82. HOW DO BUFFER SOLUTIONS WORK?
• A buffer solution has to contain things which will remove any H+ or OH-
ions that you might add to it ,otherwise the pH will change.
• Acidic & alkaline buffer solutions achieve this in different ways.

83. ACIDIC BUFFER SOLUTIONS
PREPARATION
o We'll take a mixture of ethanoic acid & sodium ethanoate.
o Ethanoic acid is a weak acid, & the position of this
equilibrium will be well to the left:
o Adding sodium ethanoate to this adds lots of extra
ethanoate ions.
o According to Le Chatelier's Principle, that will tip the
position of the equilibrium even further to the left.

84. • The solution will therefore contain these important
things:
lots of un-ionised ethanoic acid;
lots of ethanoate ions from the sodium ethanoate;
enough hydrogen ions to make the solution acidic.

86. ALKALINE BUFFER SOLUTIONS
PREPARATION
• We'll take a mixture of ammonia and ammonium
chloride solutions as typical.
• Ammonia is a weak base, and the position of this
equilibrium will be well to the left:
• Adding ammonium chloride to this adds lots of extra
ammonium ions. According to Le Chatelier's
Principle, that will tip the position of the equilibrium
even further to the left.

87. • The solution will therefore contain these important
things:
lots of unreacted ammonia;
lots of ammonium ions from the ammonium chloride;
enough hydroxide ions to make the solution alkaline.

89. EXERSISE
1. Calculate the pH of a solution containing 0.08 M
acetic acid & 0.02 M sodium acetate. [pKa = 4.7]
2. Calculate the pH of 500 ml of 0.1 M weak acid after
addition of 100 ml of 0.1 M KOH. [pKa = 5]

90. SUMMARY
 Solution are homogeneous mixture of one or more substances( solutes)
dispersed molecularly in a sufficient quantity of dissolving medium
(solvents)
 Solutions amount may expressed as dilutions, concentrations, saturated,
and super saturated solution.
 Molarity of solution (Molarity, M) number of moles of solute per
liter of solution or the gram molecular weight of a compound
per liter of solution, where as a Normality of a solution
(Normality, N) is the number of grams equivalents of solute
by weights per liter of solution (not solvent

91. SUMMARY….
• Standard Solution is a solution whose concentration
is exactly known. Primary standard solution highly
purified chemical that can be measured directly to
produce a substance of known concentration in a
given solution. Secondary standard solution is
defined as a substance of lower purity & whose
concentration is determined by comparison to a
primary standard

92. • Acids – are proton donors, or dissociate to furnish
hydrogen ions; where as Bases – are proton acceptors or
dissociate to furnish hydroxyl ions in aqueous solution
• pH is defined as the negative logarithm of the Hydrogen
ion concentration. pH = - log [H+] in mol/L
• Buffers are chemical system that prevents change in the
concentration of another chemical substance. There are
three types of buffers: Neutral buffer eg. phosphate
buffer; Acidic buffer eg. citrate buffer; and Alkaline
buffer eg.– borate buffer

93. End of
the session

94. CHAPTER THEREE
Introduction to radiant energy

95. OBJECTIVES
At the end of this chapter the student will be
able to:
• Definition of terms
• Discuss radiant energy
• Describe properties of EMR
• Explain about interaction of EMR with matter
• Discuss basic law of absorption: Beer-Lambert’s law

96. OUTLINE OF RADIANT ENERGY
LECTURE
• Introduction to radiant energy
• Properties of EMR
• Interaction of EMR with matter
• Electromagnetic spectrum
• Absorption measurements: Beer-Lambert’s law, stray light

97. INTRODUCTION TO RADIANT ENERGY
Electromagnetic radiation
• Radiation showing electric & magnetic characteristics in
the form of waves or photons is termed as
electromagnetic radiation
• It travels at approx. 3 x 105 km/s in the vacuum of space.
• In materials which are transparent to electromagnetic
radiation, the velocity is slightly less than the velocity in a
vacuum.

99. DUAL NATURE OF EMR ENERGY
• Energy transfers in the physical world either by
waves or particles
• In general, electromagnetic radiation behaves:
as a wave when moving through space,
as a particle when it interacts with matter

100. WAVE PROPERTIES
• Wave is the way of transferring an energy from one
place to another
• Consists of discrete packets of energy or quanta
called photons
• It Can be described by:
1. Velocity (c )
2. Amplitude
3. Wave length (λ)
4. Frequency (ν)

102. 1. Amplitude
 The height the wave crest or troughs from the baseline
 Governs brightness of light
2. Wavelength
 distance between two wave crests or troughs
 One full wave cycle (crest top to next crest top).
3. Frequency
 how fast it oscillates (goes up & down) measured in cycles (remember
crest to crest) per second.
 The number of wave crests per second, that is the number of wave crests
that passes by a given point in one second.
1cycle/second = 1 Hz (hertz) = 1s-1.

103. 4. Speed of Light (velocity)
 Speed of the wave
 For example;
• water - few meters per second
• Sound wave – 340 m/sec
• Light - 3 x 108 meters/second

104. RELATIONSHIP BETWEEN C, V & λ .
 The longer the wavelength the lower the frequency, or
the shorter the wavelength the higher the frequency.
 This relationship is expressed in the formula
ν = c/λ where: - ν - frequency of light in cycle/sec.
- c - speed of EM wave in vacuum
- λ - wave length in cm

105.  E = hv
 Where h = planck’s constant (6.62x10-27 erg.sec)
ν = frequency
E = energy
 E = h c /λ

106. INTERACTION OF EMR WITH MATTER
• In order to use photometric instruments correctly &
to be able to develop & modify spectroscopic
techniques it is necessary to understand the
principle of interaction of radiation with matter.
• The only way to observe electromagnetic radiation
is by its interaction with matter.

107. • It involves:
 Diffraction
 Reflection
 Refraction
 Dispersion
 Absorption & transmission

108. 1. Diffraction
Is the change of direction of the EM beam when it
strikes the edge of an opaque body or it passes
through a small hole

109. 2. Refraction
• bending of light as it passes through materials of
different optical density

110. 3. Reflection
 When radiation falls on silver coated glasses, the
beam of the radiation returns towards the source of
radiation.

111. 4. Dispersion
• is the change in refractive index with a change in
wavelength
• The velocity of light in a materials & its refractive
index depends on the wavelength of the light.
• This causes the light to be refracted by different
amounts according to the wavelength (or color).
• This gives rise to the colors seen through a prism.

112. • Rainbows are caused by dispersion of light inside the
raindrop & total internal reflection of light from the
back of raindrops.
• The following is a chart giving the index of refraction
for various wavelengths of light in glass

114. The electromagnetic Spectrum
• Spectrum is an ordered arrangement of radiant energy
according to the wavelength.
A. Continuous spectrum
• A spectrum which is composed of visible lights of all
wavelengths are called Continuous spectrum
• It is a continuous spectrum because one color fades into
another.
• E.g. sun light or light from ordinary incandescent bulb

115. COMPONENT ENERGIES OF THE
ELECTROMAGNETIC SPECTRUM
a. Radio waves
 The longest- from a few meters to longer than the
size of the earth.
 They can travel long distance in the atmosphere
b. Microwaves
 The wavelength is from about 1 millimeter to 1
meter.
 Used in communication, radar & cooking

116.  Wavelength from 10-3 to 10-6 (micron)
 Ranges from approx 12,500 – 50 cm-1
Used in toxicology and molecular structure determination
• 4000 to 1000 cm -1 – used for the analysis of organic
compounds
• 1000 to 400 cm -1 - is used for the analysis of inorganic
compounds
• 12,500 to 4000 cm -1 is not helpful for such analysis
c. Infra red

117. d. Visible
• It is a very small portion of the total EM spectrum visible
to human eye
• It ranges from 700nm (at red light) to 400nm (violet)
• the visible color of a solution corresponds to the
wavelength of the light that are transmitted, not absorbed,
by the solution.

118. • The different colors have different wavelengths &
frequencies.
• The rest of the EM spectrum is not visible to the human eye
• Source of visible light:
tungsten lamp.

120. Note:
• A substance that absorbs green light at 500 nm reflects or
transmits all other lights or wavelengths & appears as
purple.
• To measure the concentration of a blue solution, light at
about 590 nm is passed through the solution
• The amount of yellow light absorbed varies directly as the
concentration of the absorbing substance in solution

121.  The absorbed color is the complementary of the
transmitted color.
 Thus to make absorption measurement, one must use the
wavelength at which a colored solution absorbs light.
 For example, a red solution absorbs green light &
transmitted red light. Therefore, a red solution should be
measured at 490 to 550nm

122. • In photometer using filter used as a monochromator,
the filter chosen is usually complementary to the
color of the solution to be measured.
 blue solution – yellow filter
 Yellow solution –blue filter
 Red solution – blue green filter
 Blue green solution – red filter

123. e. UV Light
• It ranges from 400 to 100 nm
• It is dangerous to tissue & cells (common sun burns)
• It is obtained by energy transition in the valence electrons of
the molecules.
• Widely used in the quantitative & qualitative determination of
clinical chemistry tests.

124. f. X- ray
• It ranges from 100 to 0.1 nm
• High frequency, high energy waves that can penetrate
several centimeters into most solid matter.
• Used in radiological diagnosis

125. g. Gamma rays
 it ranges from 0.1 to less than 10-16
 It is the highest energy ray in the EM spectrum
 It is generally produced in nuclear reactions & not as common in nature,

127. B. LINE/ATOMIC EMISSION SPECTRUM
• a spectrum with only certain colors. (NOT continuous
like sunlight)
• Samples of elements emit light when they are vaporized
(heated) or electricity passes through them.
• Every element has a unique line

128. • The wavelength of the line are characteristics of a
particular element
• It can be used for qualitative identification & quantitative
determination of elements in an unknown mixture
• For example, flame photometer, atomic absorption
spectrophotometer.

129. 5. ABSORPTION & TRANSMISSION
• when some radiant energy passing through a solution,
transparent glass, or semitransparent substances
– Some amount of light is transmitted &
– Some is absorbed or trapped by the medium.

130. ABSORPTION MEASUREMENT
• Many determinations in clinical chemistry are based
on the measurement of the radiant energy
 Emitted (e.g Fluorometer)
 Transmitted
 Absorbed (absorption spectrophotometer)
 Reflected (reflectance photometer)
 Refracted (refractometer)

131. Light transmittance
• Is defined as the proportion of the incident light that
is transmitted.

132. • Transmittance (T) = I/Io
where I = transmitted light
Io = original incident
• Usually this ratio is described as a percentage:
• %T = I/ Io x 100%
• As the concentration of a compound in solution increases,
more light is absorbed by the solution & less light is
transmitted.

133. LIGHT ABSORBANCE
• The relationship between %T & concentration is not
linear but varies inversely & logarithmically.
• As a result it is more convenient to use the concept of
absorbance to avoid the use of logarithmic units.
• However the concept of transmittance is important
because only transmitted light can be measured.

134. • % T can be related light of absorbance of a solution
by:
• Absorbance, A = log10 I0 / I
A = log10 1 / T
A = log10 100 / %T
A = 2 - log10 %T

136. RELATIONSHIP BETWEEN ABSORBANCE &
TRANSMITTANCE

137. ABSORPTION SPECTROPHOTOMETRY
Fundamental law of absorption
a. Introduction
• When a radiant energy, Io, passes through a solution
to be analyzed,
– some of the radiant energy will be absorbed
– some of it will be transmitted

138.  The transmitted light, I, is affected by factor such as:
 Incident light
 Optical path length
 Concentration of solution

141. 2. BEERS- LAMBERT'S LAW
It is commonly referred to as Beer’s Law
1. Beer’s law
 It states that concentration of a substance is:
 directly proportional to the amount of light absorbed by the solution &
 inversely proportional to the logarithm of transmittance
 A = C
 A = a. c

143. 2. LAMBERT'S LAW
• It states that the amount of radiant energy absorbed is
directly proportional to the thickness of the
medium through which the light pass.
A = b
A = a.b

145. 3. BEERS-LAMBERT’S LAW
 A = a.b.c
 It is the combination of the two laws.

146. • Beer-Lambert's law indicates a direct
proportionality between A and c only if:
incident radiation is monochromatic
each molecule in solution acts as an independent absorbing species in
solution
Absorption takes place in a solution of uniform cross-section (a well
mixed solution)

147. • Limitations of Beer’s law-cause for deviation from the law
 non-monochromatic light
 Elevated concentration
 Solvent absorption
 Transmitted light by other mechanisms
 Non-parallel sides of cuvets

148. DEVIATIONS FROM BEER’S LAW
a. Spectral interference
• The beers-lambert’s law express the linear relationship
b/n the concentration of the sample & the absorbance
value recorded.
• However, the relationship is only an experimental one
& not a fundamental law of nature.
• As a result, the linearity is only true under certain
limiting conditions

149. • Some amount of radiation will be
– reflected from the surface of the sample holder,
– absorbed by the material of which the cell is composed or
– The solvent may also absorb or reflect radiation.
• Io = absorbed + transmitted + others

151. • To focus attention on the compound of interest,
elimination of these factors is necessary.
• This is done through the use of blank or reference
solution
• This blank should be identical to the test sample in all
aspects except the presence of the test substance
• Blank reading = Io - other loses

152. • Hence:
• Absorbed = blank – transmitted.
• Types of blank solution
1. reagent blank
2. Sample blank
3. Water saline or air blank

153. REAGENT BLANK
• reagent + solvent or
• A solution of reagents with out sample
• Used to correct high absorbance of the reagent

154. SAMPLE BLANK
• Sample + diluent
• A solution of sample & reagents missing a key
reagents that initiate the rxn or cause formation of
final rxn product.

156.  T = Is/ IR
 A = log 1/T
 A = log 1/ Is/ IR
 A = log IR / Is
 A = - log Is / IR

157. b. Stray light
• Quantitative radiation rely on radiation that reach the
detector passing through the sample
• But it is impractical, because it is difficult to design
instrument which are capable of effectively eliminating
all extraneous radiation
• Much of these unwanted radiation arises from the
scattering of the incident radiation by irregularities in
surfaces (by faults in manufacturers) or scratches

158. • Such light, stray light, results in a deviation from Beer’s law
&
• The effect is that absorbance measurements are lower than
they should be.
• It is possible to asses the proportion of stray light by
measuring the amount of radiation transmitted by samples
which are optically opaque at the wavelength to be assessed
but which transmit radiation of other wavelengths

159. • The instrument is set to zero %T with blocking light
path & 100% transmittance with a reagent blank in
the normal way & opaque substance introduced into
the sample compartment.
• The amount of light transmitted by the sample,
measured in percentage transmittance is quoted as the
stray light at a specified wavelength.

160. SUMMARY
• Radiant energy: radio waves (longest) to gamma rays (shorts)
• Properties of Electromagnetic radiation as wave and particle.
Visible light is 350-700 nm.
• Interaction of EMR with matter is one of six types:
diffraction, refraction, reflection, dispersion, absorption or
transmission.
• Basic law of absorption: Beer-Lambert’s law which is A = a.b.c

161. QUIZ
Q.1 Define the following terms
1. Analysis
2. Quality assurance
3. Reagents
4. Solution
5. Dilution
Q.2. Define qualitative, semi-quantitative and quantitative
analysis
Q.3. List the basic components of spectrophotometer
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