INTRODUCTION TO ANALYTICAL CHEMISTRY
Analytical Chemistry - science of inventing and applying the concepts, principles, and strategies for measuring the
characteristics of chemical systems
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measurement science
area of chemistry responsible for characterizing the composition of matter, both qualitatively and quantitatively.
science of measurements
Role of Analytical Chemistry
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medicinal chemistry
clinical chemistry
toxicology
forensic chemistry
materials science
geochemistry and environment
Analytical Approach
1. Identify and define the problem
a. Problem’s context
b. Type of information needed
2. Design the experimental plan/procedure
a. Establish design criteria
b. Identify potential interferents
c. Establish validation criteria
d. Select analytical method
e. Establish sampling strategy
3. Conduct the experiments to produce data relevant to the problem
a. Calibrate instruments and equipment
b. Standardize reagents
c. Gather data
4. Analyze the experimental data
a. Reduce and transform data
b. Complete statistical analysis
c. Verify results
d. Interpret results
5. Propose a solution to the problem
a. Is the answer sufficient?
b. Does answer suggest a new problem
1. Identify and define the problem
 The second aspect of analytical chemistry practice hinges on the importance of chemical measurements to
address technical problems related to manufacturing and regulation
 best accomplished through discussions with others involved in answering the problem
2. Design the experimental plan/procedure.
 Selection of the method of analysis
 Sampling
 Sample preparation
3. Perform analysis
 The samples will be collected, treated and subjected to analysis according to the plan developed in step 2.
 The first step in performing the measurements is to verify that the instrument is working properly
4. Data Analysis and Interpretation
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Once the data has been acquired, it must be converted into a format that leads to meaningful
interpretation
Note: Analytical chemists often use a variant of the scientific method, called the analytical approach to problem
solving.
Common Analytical Problems
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1. Quantitative Analyses - most common analytical problem
amount of each substance in a sample.
how much is present
 elemental analyses of a newly synthesized compound
 concentration of glucose in the blood
 amount of contaminant such as pesticide
2. Qualitative Analyses - identity of the elements and compounds in a sample.
what is present in the sample
 Screening an athlete’s urine for possible performance – enhancement drugs
 Determining the spatial distribution of Chromium, Lead and some other toxic substances on the surface or
airborne particulates
3. Characterization Analyses
 Methods for characterizing physical and chemical properties
 Determination of chemical structure
 Determination of equilibrium constant of particle size and of surface structure
4. Fundamental Analysis
 It tells us how does this method work and How can it be improved?
APPLICATIONS OF CHEMICAL ANALYSIS
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Assaying
Quality Control
Important factors which must be taken into account when selecting an appropriate method of analysis:
a. nature of the information which is sought
b. size of sample available and the proportion of the constituent to be determined
c. purpose for which the analytical data are required
TYPES OF ANALYSIS
1. Proximate analysis - in which the amount of each element in a sample is determined with no concern as to
the actual compounds present;
2. Partial analysis - deals with the determination of selected constituents in the sample;
3. Trace constituent analysis - specialized instance of partial analysis in which we are concerned with the
determination of specified components present in very minute quantity;
4. Complete analysis - the proportion of each component of the sample is determined
On the basis of sample size, analytical methods are often classified as:
1.
2.
3.
4.
5.
Macro – 0.1 g or more
Meso or semimicro – 10-2 – 10-1 g
Micro – 10-2 – 10-3 g
Submicro – 10-3 – 10-4 g
Ultramicro – below 10-4 g
COMMON TECHNIQUES
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Gravimetric analysis - substance being determined is converted into an insoluble precipitate which is
collected and weighed, or in the special case of electrogravimetry electrolysis is carried out and the
material deposited on one of the electrodes is weighed.
Electrical methods of analysis - measurement of current, voltage or resistance in relation to the
concentration of a certain species in solution
o Techniques which can be included under this general heading are:
i. Voltammetry - measurement of current at a micro-electrode at a specified voltage
ii. Coulometry - measurement of current and time needed to complete an electrochemical
reaction or to generate sufficient material to react completely with a specified reagent
iii. Potentiometry - measurement of the potential of an electrode in equilibrium with an ion to be
determined
iv. Conductimetry - measurement of the electrical conductivity of a solution
Titrimetric analysis or volumetric analysis - substance to be determined is allowed to react with an
appropriate reagent added as a standard solution, and the volume of solution needed for complete
reaction is determined.
Common types of reaction which are used in titrimetry are:
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a. neutralization (acid-base) reactions
b. complex-forming reactions
c. precipitation reactions
d. oxidation-reduction reactions
Volumetry - concerned with measuring the volume of gas evolved or absorbed in a chemical reaction.
Absorption methods are usually classified according to the wavelength involved as:
a.
b.
c.
d.
visible spectrophotometry (colorimetry)
ultraviolet spectrophotometry
infrared spectrophotometry
Atomic absorption spectroscopy - involves atomizing the specimen, often by spraying a solution of the
sample into a flame, and then studying the absorption of radiation from an electric lamp producing the
spectrum of the element to be determined.
e. Turbidimetric and Nephelometric methods - involve measuring the amount of light stopped or scattered
by a suspension
f. Emission methods - subjecting the sample to heat or electrical treatment so that atoms are raised to
excited states causing them to emit energy: it is the intensity of this emitted energy which is measured.
The common excitation techniques are:
a. emission spectroscopy - sample is subjected to an electric arc or spark plasma and the light emitted
(which may extend into the ultraviolet region) is examined
b. Flame photometry - a solution of the sample is injected into a flame
c. Fluorimetry - a suitable substance in solution (commonly a metal-fluorescent reagent complex) is excited
by irradiation with visible or ultraviolet radiation.
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 Chromatography - separation process employed for the separation of mixtures of substances.
widely used for the identification of the components of mixtures.
often possible to use the procedure to make quantitative determinations, particularly when using: Gas
Chromatography (GC) and High-Performance Liquid Chromatography (HPLC).
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INSTRUMENTAL METHODS - much faster than purely chemical procedures, they are normally applicable
at concentrations far too small to be amenable to determination by classical methods, and they find
wide application in industry.
Despite the advantages possessed by instrumental methods in many directions, their widespread adoption has not
rendered the purely chemical or 'classical' methods obsolete; the situation is influenced by three main factors:
1. The apparatus required for classical procedures is cheap and readily available in al1 laboratories, but
many instruments are expensive and their use will only be justified if numerous samples have to be
analyzed, or when dealing with the determination of substances present in minute quantities (trace,
subtrace or ultratrace analysis).
2. With instrumental methods it is necessary to carry out a calibration operation using a sample of
material of known composition as reference substance.
3. Whilst an instrumental method is ideally suited to the performance of a large number of routine
determinations, for an occasional, non-routine, analysis it is often simpler to use a classical method
than to go to the trouble of preparing requisite standards and carrying out the calibration of an
instrument.
BASIC TOOLS OF ANALYTICAL CHEMISTRY
Uncertainty in Measurements
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 International Vocabulary of Basic and General Terms in Metrology (VIM) defines uncertainty as:
parameter associated with the result of a measurement, that characterizes the dispersion of the values that
could reasonably be attributed to the measured value.
1. Calculate the average value of all the measurements
2. Calculate the deviation of each measurement, which is the absolute value of the difference between
each measurement and the average value:
3. Add all the deviations and divide by the number of measurements to obtain the average deviation:
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Significant Figures - numbers that describe the value without exaggerating the degree to which it is known
to be accurate
Rules
1.
2.
3.
4.
Any nonzero digit is significant. E.g. 123456789
Any zeros between nonzero digits are significant. E.g. 205, 2002
Any zeros used as a placeholder preceding the first nonzero digit are not significant. E.g. 0.03
When a number does not contain a decimal point, zeros added after a nonzero number may or may not be
significant. E.g. 100, which may be interpreted as having one, two, or three significant figures.
5. Integers obtained either by counting objects or from definitions are exact numbers, which are considered to
have infinitely many significant figures. E.g. 100g atoms
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 SOLUTION - mixture of two or more substances that is identical throughout (homogeneous)
can be physically separated
composed of solutes and solvents
 Solutes - substance being dissolved; smallest part of the solution
 Solvents - substance that dissolves the solute; largest part of the solution
Types of Solutions:
IDENTIFYING COMPONENTS OF SOLUTIONS
1. Gaseous solutions – air = Oxygen + Nitrogen
2. Liquid solutions – drinks = mix + water
3. Solid solutions – alloys = steel, brass
Solution
Solute
Air in balloon
O
2
How does solution form:
Ammonia water
Solvent
N
2
NH
HO
3
1. Solvent molecules are attracted to surface ions.
2. Each ion is surrounded by solvent molecules.
3. Enthalpy ( H) changes with each interaction broken
or formed.
 The ions are solvated (surrounded by solvent).
 If the solvent is water, the ions are hydrated.
 The intermolecular forces.
 Solvation - interaction of a solute with the
solvent, which leads to stabilization of the solute
in the solution.
Rubbing alcohol
(70%)
Rubbing Alcohol (40%)
Tincture of
Iodine
HO
2
Isopropyl alcohol
2
HO
Ethyl alcohol
Iodine
2
Alcohol
Sea water
Salt
HO
2
Types of solution based on solute concentration:
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Hypotonic - lower solute concentration to the solution
Hypertonic - higher solute concentration to the solution
Isotonic - solutions are equal in their solute concentrations
 Dissolution - process by which a solid, liquid or gas forms a solution in a solvent.
Dissolution is a physical change—you can get back the original solute by evaporating the solvent. If not, it
already reacted.
Dissolution process in solid:
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In solids this can be explained as the breakdown of the crystal lattice into individual ions, atoms or
molecules and their transport into the solvent.
Dissolution process in liquid and gas:
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For liquids and gases, the molecules must be compatible with the solvent for a solution to form.
Solubility - maximum amount of solute, expressed in grams, that can be dissolved in 100 g of water at a specific
temperature & pressure.
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Soluble - substance that dissolves in a solvent
Insoluble - substance that does not dissolve in a solvent
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Immiscible - 2 liquids that are insoluble
Miscible - 2 liquids that are soluble
Types of Saturation:
Kinds of Saturation
Definition
Saturated Solution
A solution with solute that dissolves until it is unable to
dissolve anymore, leaving the undissolved substances at
the bottom.
Unsaturated Solution
A solution (with less solute than the saturated solution)
that completely dissolves, leaving no remaining
substances.
Supersaturated Solution
A solution (with more solute than the saturated solution)
that contains more undissolved solute than the saturated
solution because of its tendency to crystallize and
precipitate.
Examples of Saturated Solutions:
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Carbonated water is saturated with carbon; hence it gives
off carbon through bubbles.
Adding sugar to water until it no longer dissolves creates a
saturated solution.
Continuing to dissolve salt in water until it will no longer
dissolve creates a saturated solution.
The Earth's soil is saturated with nitrogen.
Mixing powdered soap into water until it will not dissolve
creates a saturated solution.
In beer or sparkling juices there is a saturation of carbon
dioxide that is let off as a gas.
Coffee powder added to water can create a saturated
solution.
Salt added to vinegar can create a saturated solution when
the salt no longer dissolves.
Chocolate powder added to milk can create saturation at
the point that no more powder can be added.
Sugar dissolved into vinegar until it will no longer do so
creates a saturate solution.
Water can be saturated with juice powder to create a
beverage.
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Milk can be saturated with flour at which
point no more flour can be added to the
milk.
Melted butter can be saturated with salt
when the salt will no longer dissolve.
Bathing salts can saturate water when there
is no more ability to dissolve them.
Sugar can be added to milk to the point of
saturation.
Processed tea powders can be added to
water to saturate the water.
Protein powder could be used to create a
saturated solution with milk, tea, or water.
Laxative powders could saturate juice or
water with which they are mixed.
Cocoa powder could be mixed into water to
the point of saturation.
Sugar could be mixed into tea to the point
that the tea is saturated.
Coffee can be saturated with sugar when no
more will mix in to the coffee.
Factors affecting solubility:
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Nature of Solute and Solvent
o like dissolves like
o Polar substances dissolve in polar solvents.
o Nonpolar substances dissolve in nonpolar solvents.
Effect of Temperature
o solubility of solid solutes in liquid solvents increases with increasing temperature.
o The opposite is true of gases.
o Higher temperature drives gases out of solution.
o Carbonated soft drinks are more “bubbly” if stored in the refrigerator.
o Warm lakes have less O2 dissolved in them than cool lakes.
 Effect of Pressure
o Small changes in pressure have little effect on the solubility of solids in liquids or liquids in liquids
but have a marked effect on the solubility of gases in liquids.
o Henry’s Law – amount of gas dissolved in a solution is directly proportional to the pressure of the
gas over the solution. (William Henry - English chemist and physicist)
Sg = kPg
Where:
• Sg is the solubility of the gas;
• k is the Henry’s law constant for that gas in that solvent;
• Pg is the partial pressure of the gas above the liquid.
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Molecular Size
o The larger the molecules of the solute are, the larger is their molecular weight and size.
o It is more difficult for solvent molecules to surround bigger molecules.
o If all of above-mentioned factors excluded, general rule - larger particles are generally less soluble.
o If the pressure, and temperature are the same than out of two solutes of the same polarity, the
one with smaller particles is usually more soluble.
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Effect of stirring
o Stirring only increases the speed of the process - it increases movement of solvent that exposes
solute, thus enabling solubility.
o molecules in liquid substances are in constant movement - process takes place, but more time.
Intermolecular forces = H-bonds; dipole-dipole; dispersion
o Ions in water also have ion-dipole forces.
o The stronger the intermolecular attractions between solute and solvent, the more likely the
solute will dissolve.
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Concentration - amount of solute dissolved in a solvent at a given temperature
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ratio of the amount of solute to the amount of the solvent
 dilute - low concentration of solute dissolved
 concentrated - has a high concentration of solute dissolved
 Molarity (M) - concentration of a solution expressed in moles of solute per Liter of solution.
 Formality (F) - Substance total concentration without regard to its specific chemical form
calculated based on the formula weight of a substance per liter of solution.
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calculated based on the formula weight of a substance per liter of solution.
calculated based on the formula weight of a substance per liter of solution.
 Normality (N) - equivalent concentration of a solution
mainly used as a measure of reactive species in a solution and during titration reactions or particularly in
situations involving acid-base chemistry
number of gram or mole equivalents of solute present in one liter of a solution
used mostly in three common situations:
o In determining the concentrations in acid-base chemistry. For instance, normality is used to indicate
hydronium ions (H3O+) or hydroxide ions (OH–) concentrations in a solution
o It is used in precipitation reactions to measure the number of ions which are likely to precipitate in a
specific reaction.
o It is used in redox reactions to determine the number of electrons that a reducing or an oxidizing agent
can donate or accept.
 Molality (m) - amount of a substance dissolved in a certain mass of solvent.
moles of a solute per kilograms of a solvent.
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Base - Z= # of moles of H+ that would react with 1 mole of a base
NaOH + H+ → Na+ + H2O
Z=1
40/1 = 40
Ca (OH)2 + 2H+ → Ca2+ + 2H2O
Z= 2
74/2 = 37
Ionic Reaction (Precipitation Reaction) - Value of Z is based on the value of ionic charge
Ca2+ + CO32- (CaCO3) = MM = 100/2 = 50 equivalent wt.
Al2(SO4) 3 = 2Al3+ + 3SO4-2
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Equiv. 6
 Redox - Oxidation number
Eq. wt. is the number of electrons taking part in the half reaction
O2 + 4H+ + 4e -→ 2H20 Eq. 4