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Analytical Chemistry Lecture Notes
Medical Technology (BSMT1)
Emilio Aguinaldo College
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Introduction to Analytical Chemistry
Introduction
Everything is made of chemicals. Analytical chemistry determine what and how much. In other words analytical chemistry is concerned with the separation, identification, and determination of the relative amounts of the components making up a sample. Analytical chemistry is concerned with the chemical characterization of matter and the answer to two important questions what is it (qualitative) and how much is it (quantitative). Analytical chemistry answering for basic questions about a material sample: · What? · Where? · How much? · What arrangement, structure or form?
Applications of Analytical Chemistry
Analytical chemistry used in many fields: · In medicine, analytical chemistry is the basis for clinical laboratory tests which help physicians diagnosis disease and chart progress in recovery. · In industry, analytical chemistry provides the means of testing raw materials and for assuring the quality of finished products whose chemical composition is critical. Many household products, fuels, paints, pharmaceuticals, etc. are analysed by the procedures developed by analytical chemists before being sold to the consumer. · Enviermental quality is often evaluated by testing for suspected contaminants using the techniques of analytical chemistry. · The nutritional value of food is determined by chemical analysis for major components such as protein and carbohydrates and trace components such as vitamins and minerals. Indeed, even the calories in a food are often calculated from the chemical analysis. · Forensic analysis - analysis related to criminology; DNA finger printing, finger print detection; blood analysis. · Bioanalytical chemistry and analysis - detection and/or analysis of biological components (i., proteins, DNA, RNA, carbohydrates, metabolites, etc.).
Applications of analytical chemistry in pharmacy sciences.
· Pharmaceutical chemistry. · Pharmaceutical industry (quality control). · Analytical toxicology is concerned with the detection, identification and measurement of drugs and other foreign compounds (and their metabolites in biological and related specimens. · Natural products detection, isolation, and structural determination.
Steps in a Chemical Analysis
§ Define the problem. § Select a method. § Sampling (obtain sample). § Sample preparation (prepare sample for analysis). § Perform any necessary chemical separations § Analysis (perform the measurement). § Calculate the results and report.
The Language of Analytical Chemistry
Qualitative analysis: An analysis in which we determine the identity of the constituent
species in a sample.
Quantitative analysis: An analysis in which we determine how much of a constituent
species is present in a sample.
Analytes: The constituents of interest in a sample.
Matrix: All other constituents in a sample except for the analytes.
A selective reaction or test is one that can occur with other substances but exhibits a
degree of preference for the substance of interest.
A specific reaction or test is one that occurs only with the substance of interest.
Note: few reactions are specific but many exhibit selectivity.
Detection limit: A statistical statement about the smallest amount of analyte that can be
determined with confidence.
Precision and Accuracy Precision describes the reproducibility of a result. If you measure a quantity several times and the values agree closely with one another, your measurement is precise. If the values vary widely, your measurement is not very precise. Accuracy describes how close a measured value is to the “true” value. If a known standard is available, accuracy is how close your value is to the known value.
(neither precise nor accurate) (accurate but not precise) (accurate and precise) (precise but not accurate)
Classifying Analytical Techniques Classical techniques Mass, volume, and charge are the most common signals for classical techniques, and the corresponding techniques are: 1- Gravimetric techniques. 2- Volumetric techniques. 3- Coulometeric techniques.
Instrumental techniques 1- Spectroscopic methods - measuring the interaction between the analyte and electromagnetic radiation (or the production of radiation by an analyte).
2- Electroanalytic methods - measure an electrical property (i., potential, current, resistance, amperes, etc.) chemically related to the amount of analyte.
Pipette is used to deliver a specified volume of solution. Several different styles of pipets are available.
Burette is volumetric glassware used to deliver variable, but known volumes of solution. A burette is a long, narrow tube with graduated markings, and a stopcock for dispensing the solution.
Equipment for Drying Reagents, precipitates, and glassware are conveniently dried in an oven at 110°C. Many materials need to be dried prior to their analysis to remove residual moisture. Depending on the material, heating to a temperature of 110–140 °C is usually sufficient. Other materials need to be heated to much higher temperatures to initiate thermal decomposition. Both processes can be accomplished using a laboratory oven capable of providing the required temperature. Commercial laboratory ovens are used when the maximum desired temperature is 160–325 °C (depending on the model). Higher temperatures, up to 1700° C, can be achieved using a muffle furnace.
Conventional laboratory oven used for drying materials. Example of a muffle furnace used for heating samples to maximum temperatures of 1100–1700 °C.
After drying or decomposing a sample, it should be cooled to room temperature in a desiccator to avoid the readsorption of moisture. A desiccator is a closed container that isolates the sample from the atmosphere. A drying agent, called a desiccant, is placed in the bottom of the container. Typical desiccants include calcium chloride and silica gel.
(a) Ordinary desiccator. (b) Vacuum desiccator
Filtration In gravimetric analysis, the mass of product from a reaction is measured to determine how much unknown was present. Precipitates from gravimetric analyses are collected by filtration. Liquid from which a substance precipitates or crystallizes is called the mother liquor. Liquid that passes through the filter is called filtrate.
Volumetric Methods of Analysis
Titrimetric Analysis
Introduction The term titrimetric analysis refers to quantitative chemical analysis carried out by determining the volume of a solution of accurately known concentration which is required to react quantitatively with a measured volume of a solution of a substance to be determined. The solution of accurately known concentration is called standard solution.
The term volumetric analysis was used for this form of quantitative determination but it has now been replaced by titrimetric analysis. In titrimetric analysis the reagent of known concentration is called titrant and the substance being titrated is termed the titrand.
The standard solution is usually add from a long graduated tube called burette. The process of adding the standard solution until the reaction is just complete is termed titration. The point at which this occurs is called equivalence point or the theoretical (or stoichiometric) end point. The completion of the titration is detected by some physical change, produced by the standard solution itself or, more usually, by the addition of an auxiliary reagent, known as an indicator ; alternatively some other physical measurement may be used. After the reaction between the substance and the standard solution is practically complete, the indicator should give a clear visual change (either a color change or the formation of turbidity) in the liquid being titrated. The point at which this occurs is called the end point of the titration. In the ideal titration the visible end point will coincide with the stoichiometric or theoretical end point. In practice, however, a very small difference usually occurs this represents the titration error. The indicator and experimental conditions should be so selected that the difference between the visible end point and equivalence point is as small as possible. For use in titrimetric analysis a reaction must have the following conditions: 1- There must be a simple reaction which can be expressed by a chemical equation; the substance to be determined should react completely with the reagent in stoichiometric or equivalent propties. 2- The reaction should be relatively fast. (Most ionic reaction satisfy this condition.) In some cases the addition of a catalyst may be necessary to increase the speed of a reaction. 3- There must be an alteration in some physical or chemical property of the solution at the equivalence point. 4- An indicator should be available which, by a change in physical properties (color or formation of a precipitate), should sharply define the end point of the reaction.
Definition of some terms Titration Titration is the process in which the standard reagent is added to a solution of an analyte until the reaction between the analyte and reagent is complete.
Equivalence point and End point The equivalence point of a titration is a theoretical point that can not be determined experimentally. Instead, we can only estimate its position by observing some physical change associated with the condition of equivalence. This change is called the end point for titration.
Titration error The difference between the observed end point and the true equivalence point in a titration.
Indicators Indicators are often added to analyte solution in order to give an observable physical change (end point) at or near the equivalence point. In other wards indicator is a compound having a physical property (usually color) that changes abruptly near the equivalence point of a chemical reaction.
End Points in Volumetric Analysis Detection of an end point involves the observation of some property of the solution that change in a characteristic way at or near the equivalent point. The properties that have been used for this purpose are numerous and varied; they include: 1. Color due to the reagent, the substance being determined, or an indicator substance. 2. Turbidity changes resulting from the formation or disappearance of solid phase. 3. Electric conductivity of the solution. 4. Electric potential between a pair of electrodes immersed in the solution. 5. Refractive index of the solution. 6. Temperature of the solution. 7. Electric current passing through the solution.
Primary standard A primary standard is a highly purified compound that serve as a reference material in all volumetric method. The accuracy of method is critically dependent on the properties of this compound. Important requirements for primary standard are: 1- High purity. 2- Stability toward air. 3- Absence of hydrated water. 4- Ready availability at modest cost. 5- Reasonable solubility in titration medium. 6- Reasonably large molar mass so that the relative error associated with weighing the standard is minimized. Compound that meet or even approach these criteria are very few , and only a limited number of primary standard substances are available to the chemist.
Secondary standard A secondary standard is a compound whose purity has been established by chemical analysis and serves as the reference material for titrmetric method of analysis. Compound such as sodium hydroxide or hydrochloric acid cannot be considered as primary standard since their purity is quite variable. So for instance sodium hydroxide solution must be standardized against potassium hydrogen phethalate (primary standard), which is available in high purity. The standardized sodium hydroxide solution (secondary standard) may be used to standardize solutions.
Standard solution Standard solution is the reagent of exactly known concentration that is used in titrimetric analysis. Standard solutions play a central role in all titrimetric method of analysis. Therefore we need to consider the desirable properties for such solutions, how they are prepared and how their concentration are expressed.
Desirable properties of standard solutions The ideal standard solution for titrmetric method will: 1- be sufficiently stable so that it is only necessary to determine the concentration once,
2- Precipitation reaction. These depend upon the combination of ions to form a simple precipitate as in the titration of silver ion with solution of chloride. No change in oxidation state occurs. 3- Complex formation reaction. These depend upon the combination of ions, other than hydrogen or hydroxide ion, to form a soluble slightly dissociated ion or compound, as in the titration of a solution af a cyanide with silver nitrate. Ethylendiaminetera-acetic acid, largely as the disodium salt of EDTA, is a very important reagent for complex formation titration and has become on of the most important reagents used in titrimetric analysis. 4- Oxidation-reduction reaction. Under this heading are included all reactions involving change in oxidation number or transfer of electrons among the reactive substance. The standard solutions are either oxidizing or reducing agents.
Titration Curves
To find the end point we monitor some property of the titration reaction that has a well- defined value at the equivalence point. For example, the equivalence point for a titration of HCl with NaOH occurs at a pH of 7. We can find the end point, therefore, by monitoring the pH with a pH electrode or by adding an indicator that changes color at a pH of 7.
Acid-base titration curve for 25 mL of 0 M HCI with 0 M NaOH.
Suppose that the only available indicator changes color at a pH of 6. Is this end point close enough to the equivalence point that the titration error may be safely ignored? To answer this question we need to know how the pH changes during the titration.
A titration curve provides us with a visual picture of how a property, such as pH, changes as we add titrant. We can measure this titration curve experimentally by suspending a pH electrode in the solution containing the analyte, monitoring the pH as titrant is added. We can also calculate the expected titration curve by considering the reactions responsible for the change in pH. However we arrive at the titration curve, we may use it to evaluate an indicator's likely titration error. For example, the titration curve in the above figure shows us that an end point pH of 6 produces a small titration error. Stopping the titration at an end point pH of 11, on the other hand, gives an unacceptably large titration error.
A titration curve is a plot of reagent volume added versus some function of the analyte concentration. Volume of added reagent is generally plotted on the x axis. The measured parameter that is a function of analyte concentration is plotted on the y axis.
Two general titration curve types are seen:
- Sigmoidal curve - a "z" or "s"-shaped curve where the y axis is a p-function of the analyte (or the reagent reacted with the analyte during titration) or the potential of an ion-specific electrode.
The equivalent point is observed in the middle of the "middle" segment of the "z" or "s." Examples of Sigmoidal titration curves
Complexation titration Redox titration Precipitation titration.
- Linear-segment curve - a curve generally consisting of two line segments that intersect at an angle.
Titrations Based on Acid-Base Reactions
The earliest acid-base titrations involved the determination of the acidity or alkalinity of solutions, and the purity of carbonates and alkaline earth oxides. Various acid-base titration reactions, including a number of scenarios of base in the burette, acid in the reaction flask, and vice versa, as well as various monoprotic and polyprotic acids titrated with strong bases and various weak monobasic and polybasic bases titrated with strong acids. A monoprotic acid is an acid that has only one hydrogen ion (or proton) to donate per fomula. Examples are hydrochloric acid, HCl, a strong acid, and acetic acid, HC 2 H 302 , a weak acid. A polyprotic acid is an acid that has two or more hydrogen ions to donate per formula. Examples include sulfuric acid, H 2 S0 4 , a diprotic acid, and phosphoric acid, H 3 P0 4 , a triprotic acid. A monobasic base is one that will accept just one hydrogen ion per formula. Examples include sodium hydroxide, NaOH, a strong base; ammonium hydroxide, NH 4 OH, a weak base; and sodium bicarbonate, NaHC0 3 , a weak base. A polybasic base is one that will accept two or more hydrogen ions per formula. Examples include sodium carbonate, Na 2 CO 3 , a dibasic base, and sodium phosphate, Na 3 P0 4 , a tribasic base.
Acid-Base Titration Curves In the overview to the titration we noted that the experimentally determined end point should coincide with the titration’s equivalence point. For an acid-base titration, the equivalence point is characterized by a pH level that is a function of the acid-base strengths and concentrations of the analyte and titrant. The pH at the end point, however, may or may not correspond to the pH at the equivalence point. To understand the relationship between end points and equivalence points we must know how the pH changes during a titration. In this section we will learn how to construct titration curves for several important types of acid-base titrations.
Titrating Strong Acids and Strong Bases For our first titration curve let’s consider the titration of 50 mL of 0 M HCl with 0 M NaOH. For the reaction of a strong base with a strong acid the only equilibrium reaction of importance is
H 3 O+(aq) + OH-(aq) = 2H 2 O(l) The first task in constructing the titration curve is to calculate the volume of NaOH needed to reach the equivalence point. At the equivalence point we know from reaction above that
Moles HCl = moles NaOH
or MaVa = MbVb
where the subscript ‘a’ indicates the acid, HCl, and the subscript ‘b’ indicates the base, NaOH. The volume of NaOH needed to reach the equivalence point, therefore, is
MaVa (0 M)(50 mL) Veq = Vb = --------------- = ------------------------------ = 25 mL Mb (0 M)
Before the equivalence point, HCl is present in excess and the pH is determined by the concentration of excess HCl. Initially the solution is 0 M in HCl, which, since HCl is a strong acid, means that the pH is
pH = -log[H 3 O+] = -log[HCl] = -log(0) = 1. The equilibrium constant for reaction is (Kw)–1, or 1 × 10 14. Since this is such a large value we can treat reaction as though it goes to completion. After adding 10 mL of NaOH, therefore, the concentration of excess HCl is
moles excess HCl MaVa - MbVb [HCl] = ---------------------------- = ------------------------ total volume Va + Vb
(0 M)(50 mL) - (0 M)(10 mL) = ------------------------------------------------------------ = 0 M 50 mL + 10 mL
giving a pH of 1.
At the equivalence point the moles of HCl and the moles of NaOH are equal. Since neither the acid nor the base is in excess, the pH is determined by the dissociation of water.
Kw = 1 × 10-14 = [H 3 O+][OH–] = [H 3 O+] 2
[H 3 O+] = 1 × 10–7 M
Thus, the pH at the equivalence point is 7.
Finally, for volumes of NaOH greater than the equivalence point volume, the pH is determined by the concentration of excess OH–. For example, after adding 30 mL of titrant the concentration of OH– is
moles excess NaOH MbVb - MaVa [OH] = --------------------------------- = ----------------------- total volume Va + Vb
(0 M)(30 mL) - (0 M)(50 mL) = ---------------------------------------------------------- = 0 M 50 mL + 30 mL
To find the concentration of H 3 O+, we use the Kw expression
Kw 1 × 10- [H 3 O+] = ---------- = ------------------ = 8 × 10- [OH-] 0.
giving a pH of 12.
x = [H 3 O+] = 1 × 10 –
We can use the following equation
[ H 3 O+]= Kac(HA )
At the beginning of the titration the pH is 2. Adding NaOH converts a portion of the acetic acid to its conjugate base. CH 3 COOH(aq) + OH-(aq) = H 2 O(l) + CH 3 COO–(aq)
Any solution containing comparable amounts of a weak acid, HA, and its conjugate weak base, A–, is a buffer. As we learned before, we can calculate the pH of a buffer using the Henderson-Hasselbalch equation.
[ ]
[ ]
log
HA
A
pH pKa
= +
The equilibrium constant for the above reaction is large (K = Ka/Kw = 1 x 109 ), so we can treat the reaction as one that goes to completion. Before the equivalence point, the concentration of unreacted acetic acid is
moles unreacted CH 3 COOH MaVa - MbVb [CH 3 COOH] = ---------------------------------------------- = ------------------------ total volume Va + Vb
and the concentration of acetate is
moles NaOH added MbVb [CH 3 COO– ] = ---------------------------------- = ---------------- total volume Va + Vb
For example, after adding 10 mL of NaOH the concentrations of CH 3 COOH and CH 3 COO– are
(0 M)(50 mL) - (0 M)(10 mL) [CH 3 COOH] = -------------------------------------------------------- = 0 M 50 mL + 10 mL
(0 M)(10 mL) [CH 3 COO– ] = ---------------------------- = 0 M 50 mL + 10 mL
giving a pH of
4. 16
[ 0. 0667 ]
[ 0. 0167 ]
pH= 4. 76 +log =
A similar calculation shows that the pH after adding 20 mL of NaOH is 4.
At the equivalence point, the moles of acetic acid initially present and the moles of NaOH added are identical. Since their reaction effectively proceeds to completion, the predominate ion in solution is CH 3 COO–, which is a weak base. To calculate the pH we first determine the concentration of CH 3 COO–.
moles CH 3 COOH (0 M)(10 mL) [CH 3 COO– ] = ----------------------------- = ---------------------------- = 0 M total volume 50 mL + 50 mL
The pH is then calculated for a weak base.
CH 3 COO-(aq) + H 2 O(l) = OH–(aq) + CH 3 COOH(aq)
[ OH - ]= Kbc(B )
[OH-] = 5 X I0-6 M The concentration of H 3 O+, therefore, is 1 × 10–9, or a pH of 8.
After the equivalence point NaOH is present in excess, and the pH is determined in the same manner as in the titration of a strong acid with a strong base. For example, after adding 60 mL of NaOH, the concentration of OH– is
moles excess NaOH MbVb - MaVa [OH] = --------------------------------- = ----------------------- = 0 total volume Va + Vb
(0 M)(60 mL) - (0 M)(50 mL) = ---------------------------------------------------------- = 0 M 50 mL + 60 mL
giving a pH of 11. The table and figure below show additional results for this titration.
The calculations for the titration of a weak base with a strong acid are handled in a similar manner except that the initial pH is determined by the weak base, the pH at the equivalence point by its conjugate weak acid, and the pH after the equivalence point by the concentration of excess strong acid.
Analytical Chemistry Lectures Notes ............................................................. 19
Precipitation Titrations
Thus far we have examined titrimetric methods based on acid-base reactions. A reaction in which the analyte and titrant form an insoluble precipitate also can form the basis for a titration. We call this type of titration a precipitation titration. One of the earliest precipitation titrations, developed at the end of the eighteenth century, was for the analysis of K 2 CO 3 and K 2 SO 4 in potash. Calcium nitrate, Ca(N0 3 ) 2 , was used as a titrant, forming a precipitate of CaCO 3 and CaSO 4 The end point was signaled by noting when the addition of titrant ceased to generate additional precipitate. The importance of precipitation titrimetry as an analytical method reached its zenith in the nineteenth century when several methods were developed for determining Ag+ and halide ions.
Precipitation Reactions
A precipitation reaction occurs when two or more soluble species combine to form an insoluble product that we call a precipitate. The most common precipitation reaction is a metathesis reaction, in which two soluble ionic compounds exchange parts. When a solution of lead nitrate is added to a solution of potassium chloride, for example, a precipitate of lead chloride forms. We usually write the balanced reaction as a net ionic equation, in which only the precipitate and those ions involved in the reaction are included. Thus, the precipitation of PbCl 2 is written as
Pb2+(aq) + 2Cl-(aq) = PbCl 2 (s) In the equilibrium treatment of precipitation, however, the reverse reaction describing the dissolution of the precipitate is more frequently encountered.
PbCl 2 (s) = Pb2+(aq) + 2Cl-(aq)
The equilibrium constant for this reaction is called the solubility product, Ksp, and is given as
Ksp = [Pb2+] [Cl-] 2 = 1 X I0- Note that the precipitate, which is a solid, does not appear in the Ksp expression. It is important to remember, however, that equation is valid only if PbCl 2 (s) is present and in equilibrium with the dissolved Pb2+ and Cl.
Titration Curves
The titration curve for a precipitation titration follows the change in either the analyte's or titrant's concentration as a function of the volume of titrant. For example, in an analysis for I using Ag+ as a titrant
Ag+(aq) + I (aq) = AgI(s)
the titration curve may be a plot of pAg or pI as a function of the titrant's volume. As we have done with previous titrations, we first show how to calculate the titration curve.
Calculating the Titration Curve
As an example, let's calculate the titration curve for the titration of 50 mL of 0 M Cl- with 0 M Ag+. The reaction in this case is
Ag+(aq) + Cl- (aq) = AgCl(s)
Analytical Chemistry Lectures Notes ............................................................. 20
The equilibrium constant for the reaction is
K= (Ksp)-1 = (1 X 10-10)-1 = 5 X 10 9
Since the equilibrium constant is large, we may assume that Ag+ and Cl- react completely.
By now you are familiar with our approach to calculating titration curves. The first task is to calculate the volume of Ag+ needed to reach the equivalence point. The stoichiometry of the reaction requires that
Moles Ag+ = moles Cl- or MAgVAg = MClVCl
Solving for the volume of Ag+
MClVCl (0 M)(50 mL) VAg = ---------------- = --------------------------------- = 25 mL MAg (0 M)
shows that we need 25 mL of Ag+ to reach the equivalence point.
Before the equivalence point Cl- is in excess. The concentration of unreacted Cl- after adding 10 mL of Ag+, for example, is
moles excess Cl- MClVCl - MAgVAg [Cl-] = ----------------------------- = ----------------------- total volume VCl + VAg
(0 M)(50 mL) - (0 M)(10 mL) = ---------------------------------------------------------- = 2 x I0-2 M 50 mL + 10 mL
If the titration curve follows the change in concentration for Ch, then we calculate pCl as
pCl = -log[Cl-] = -log(2 x 10-2) = 1.
However, if we wish to follow the change in concentration for Ag+ then we must first calculate its concentration. To do so we use the Ksp expression for AgCl
Ksp=[Ag+][Cl-] = 1 x l0-
Solving for the concentration of Ag+
Ksp 1 x l0- [Ag+] = ---------- = ----------------- = 7 x 10-9 M [Cl-] 2 x 10-
Analytical Chemistry Lecture Notes
Course: Medical Technology (BSMT1)
University: Emilio Aguinaldo College
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