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Compulsory textbook Fundamentals of Analytical Chemistry, D. A. Skoog, D. M. West, F. J. Holler and S. R. Crouch, Brooks/Cole 2004 Brooks/Cole, 2004, 8th Edition Recommended textbooks Principles of Instrumental Analysis, D. A. Skoog, F. J. Holler and T. A. Nieman, Saunders College Publishing 1999 Publishing, 1999. Quantitative Analytical Chemsitry, J. S. Fritz and G. H. Schenk, Allin and Bacon, 1987. Instrumental Methods of Analysis, H. H. Willard et al., Wadsworth Publ. Co., 1988. Topics of the first semester 1. 2 2. 3. 4. 5. 6. 7 7. 8. General introduction (1) Fundamental concepts in analytical chemsitry (1) Gravimetric methods in the analysis (1) Titrimetry – general principles and concepts (2) Precipitation titrations (argentometry) (1) Neutralization titrations (acidi-alkalimetry) (2) Complexometric titrations (chelatometry) (1) Redox titrations (oxidi-reductometry) (3) 1 Analytical chemistry involves separating, identifying and d t determining i i the th relative l ti amounts t of f the th components (analytes) of the sample Qualitative analysis what is present? – chemical identity of the species in the sample (preceeds quant. anal.) Quantitative analysis how much is present? – percentage or mass of the analyte in the sample Separation techniques (chromatographies) different components may interfere one with another Analysis types 1. Complete analysis – each constituent is analysed y – each element is 2. Ultimate ((elemental)) analysis analysed 3. Partial analysis – the amount of selected compounds/atoms/components Examples water analysis blood sample analysis N,S,P,C-content in foodstuffs serial analysis of a pharmaceutical product household gas analysis air analysis etc., etc., etc. 2 Methods of analytical chemistry Classical gravimetry i t volumetric methods or titrations Instrumental electroanalysis spectrometric analysis magnetic methods thermal methods miscellaneous methods To be considered 1. accuracy & reliability required vs. economics 2. no. of samples to be analysed 3. complexity of the samples 3 representative sampling – when the sample truly represents the object to be analysed grinding (homogeneity) drying (deliquescence) homogeneous sample: its constituents can be d distinguished h d visually ll or with the aid of a light microscope ( heterogeneous sample) 4 replicate samples: portions of the material of (approximately) the same size that are carried through the analytical procedure weighing (by an analytical balance – measurement of f mass)) pipetting (by a pipette – measurement of volume) preparing aqueous solutions solubilization (digestion) 5 interference: species other than the analyte, which interferes with the results of the measurement, i.e., causes errors The measured property, X has to vary in a known and reproducible way with the concenctration of the analyte, cA Ideally cA = k×X X – the signal k – characteristic to the method, usually s ll unknown k (except ( x t gravimetry and coulometry) calibration – the process of determining k 6 For the calculations 1. experimental data 2. stoichiometry 3. instrumental data are required uncertainties associated with the measurements must be known – errors in the chemical analysis 7 Chapter 2. Chemicals, Apparatus and Unit Operations of Analytical Chemsitry - dealt with in p practical (compulsory) ( mp y) Chapter 3. Using Spreadsheets in Analytical Chemistry - dealt with in practical (optional) Calculations used in analytical chemistry Atom –the smallest particle of an element Molecule - the smallest particle of a compound Compounds are combination of elements – molecules are made up of atoms The important thing for an (analytical) chemist is the number of atoms reacting (and not the mass) Atomic mass (Ar): relative masses based on the 12C isotope Molecular mass (Mr): the sum of the atomic masses of the atoms that make up the molecule The chemical mass unit: the mole (1 mole = 6.022×1023 atoms of an element or molecules of a compound) Number of moles (n) = grams of material (m) _____________________________ formula mass (Ar or Mr) 8 Expressing concentration of solutions 1. Molar concentration (molarity) the number of moles of solute present in 1 L of solution c= number of moles of the solute volume of solution unit: mole/litre or mole/dm3 or M unit (equal to mmol/mL!!!) Expressing concentration of solutions 2. Molal concentration (molality or Raoult’s- concentration) – m= number of moles of solute mass of solvent unit: mole/kg Advantage of m over c: m is independent of temperature 9 Expressing concentration of solutions 3. Mole fraction – X= number of moles of solute the moles of solvent + the moles of solute unit: Grams per volume – the mass of the solute divided by the volume of solution mass of solute volume of solution unit: g/L Expressing concentration of solutions 4. ppm – the mass of the solute in mg divided by the volume of the solution in litre mass of solute in mg concentration in ppm = volume of solution in litre The mass of 1 litre of water equals to 1000g Unit: ppm – part(s) per million, ppb – the mass of the solute in µg divided by the volume of the solution in litre mass of solute in µg concentration in ppb = volume of solution in litre Unit: ppb, part(s) per billion 10 Expressing concentration of solutions 5. Mass percent – the mass of solute divided by the mass of solution mass percent = mass of solute mass of solution x100 Unit : g/100g or m/m% Volume percent - the volume of solute divided by the volume of solvent vol% = volume l of f solute l t volume of solution x100 Unit : mL/100mL or V/V% Expressing concentration of solutions 6. Analytical molarity – the total number of moles of solute present in a given volume of solution (it says nothing about the actual state of the solute, whether it ionizes or not, etc.) S b l c or cT Symbol: Equilibrium molarity – the concentration of ions or molecules actually present in solution, taking into account the possible dissociation of the solute into ions Symbol: […] The analytical concentration is equal to the sum of the equilibrium concentrations of the various forms of the solute Example: HAc H+ + Ac- cHAc = [Ac-] + [HAc] – this is a mass balance equation 11 Errors in chemical analysis – how certain can we be about the results we obtain? Obtained values (results) for a given quantity from N replicates: x1, x2, x3, …, xN Mean (median, arithmetic mean, average), x x= 1 N N ∑x i =1 i Precision – the reproducibility of the measurements, or the closeness of results that have been obtained exactly in the same way; can be obtained b i db by repeating i the h measurements (MEMO-technique: pre=rep) Accuracy – the closeness of our measurements to the true or accepted values; cannot be obtained by repeating the measurements; expressed in terms of the absolute error: E = xtrue - xi 12 Types of errors Random (indeterminate) error – causes data to be scattered symmetrically y y around the mean value; associated with the precision (or the reproducibility) of the measurement Systematic (determinate) error – (for example instrumental, method or personal error); causes the mean of the data set to differ from the true value; associated with the accuracy of the measurement Gross error – they occur occasionally and lead to outliers Characterization of random errors, i.e., the precision of the measurement Gaussian (normal error) curve shows the symmetrical distribution of data around the mean of an infinite set of data 13 Sample standard deviation, s – the measure of precision of a measurement N s= ∑ (x − x ) 2 i i =1 N −1 di = x – xi – the deviation of the i-th result, xi from the mean; N - the total number of the measurements; N-1 N 1 – number of degrees of freedom s – standard deviation s2 – sample variance Significant figures The significant figures in a number are all the certain digits plus the first uncertain digit – most important example: reporting a burette reading l t us say, that let th t th the b burette tt iis of f 25 mL L capacity it smallest division is 0.1 mL 12.24 mL the first three digits are certain the last digit is estimated, i.e., uncertain if you report 12 mL – rounding off error (ca. 1.8 %) if you report 12.24478 mL – nobody believes you (rightly so) if you try to read < 0.1 mL on the same burette – meaningless result - another very important example: reading an analytical balance 0,9668 g the sample weighed must be at least 100 mg (to keep error at ≤1% level) 14 Sampling The analytical method of choice depends on the sample size and constituent type sample size > 0.1 g 0.01-0.1 g 0.0001 g – 0.01 g < 0.0001 g type of analysis macro semimicro micro ultramicro analyte l level l l 1% - 100% 0.01%(100 ppm) -1% 100 ppm – 1 ppb < 1 ppb type of f constituent major minor trace ultratrace Minimizing errors in analytical procedures 1. 2. 3. 4. 5. 6. 7. Choosing the correct blank solution Application of separation techniques – elimination of i t f interferences s Saturation – deliberate addition of large amount of interfering components to all samples and standards (this may degrade sensitivity and detectability) Matrix modification – a non-interfering component is added to modify the response, to make it independent of the presence of the interfering species Adding g of a masking g agent g – it selectively y reacts with the interfering component and makes it „invisible” Dilution method Matrix matching method – for example, synthetic seawater 15 Classical methods of chemical analysis includes gravimetry titrimetry argentometry acidi-alkalimetry complexometry redox titrations Gravimetric methods are quantitative methods that are based on determining the mass of the pure compound to which the analyte is chemically related The mass is always measured on an (accurate) analytical balance Types T p of f gravimetric im t i m methods th ds precipitation gravimetry volatilazation gravimetry electrogravimetry thermogravimetry gravimetric titrimetry Steps of precipitation gravimetry: 1 an excess precipitating reagent added to the sample, 1. sample thus the analyte converted into sparingly soluble product (precipitate) 2. precipitate is filtered 3. precipitate is washed from impurities 4. precipitate is dried or ignited (to convert it to a product with known composition) 5. precipitate is weighed 16 Precipitation gravimetry A successfull gravimetric determination meets the following criteria 1 1. The h analyite l must be b completely l l (quantitatively) ( l ) precipitated d 2. The precipitating agent reacts selectively or, at least, specifically with the analyte • Selective reagent reacts only with a single chemical species (rare) • Specific reagent reacts with several, but limited number of chemical species (more common) 3 The precipitate must easily filtered and washed free from 3. contaminants 4. Must be of sufficiently low solubility (to avoid loss of the analyte) 5. Must be unreactive with constituents of the atmosphere 6. Its weighed form must be of known composition (gravimetric factors) Solubility of precipitates Precipitate – it is formed from a solution which is supersaturated with respect to the solute; when no more precipitate is ible to form, the remaining solution is called saturated solution Types of electrolyte solutions: 1. Non-saturated (or undersaturated) 2. Saturated 3. Supersaturated 1 2 3 17 Solubility of precipitates Solubility (or equilibrium solubility,S): the concentration of a saturated solution in molarity at a given temperature; characteristic to the given salt (depends on solvent and temperature) MA M+ + A- MxAy xMy+ + yAx- S = [My+ ]/x = [Ax-]/y Solubility y product p (L, Ksp) : the equilibrium q constant for the components of the precipitate in a saturated solution (i.e., in a solution, which contains some precipitate) Ksp = [M+ ][A-] Ksp = [My+ ]x[yAx-]y Solubility and solubility product [ K sp = M y + ] [A ] x S = x+ y x− y = ( Sx) x ( Sy ) y K sp xx y y E Examples: l calculate l l t solubility l bilit f for AgCl in water, at 25 oC Ksp = 1.0.10-10 Ag2CrO4 in water, at 25 oC Ksp = 1.1.10-12 Bi2S3 in water, at 25 oC Ksp = 1.0.10-72 18 Factors influencing the solubility of a precipitate 1. Common ion effect – common ion will reduce the concentration of the other ion (and therefore the solubility) of ppt (unless the common ion forms complex compound with the ppt) 2. Effect of pH – • if the anion gets protonated, decrease of pH increases solubility • if the cation hydrolyses, increase of pH increases solubility 3. Effect of complexation – complexation always increases solubility 4. Effect of foreign ions – foreign ions in small quantities increase, while in large quantities decrease solubility (latter is called salting-out) …now back to gravimetry… Steps of precipitation gravimetry: 1. Precipitation: an excess precipitating reagent added to the sample, sample thus the analyte converted into precipitate 2. Filtration: precipitate is separated from the solution via filtration 3. Washing: precipitate is washed from impurities 4. Drying: precipitate is dried or ignited (to convert it to a product with known composition) 5. Weighing: precipitate is weighed 19 1. Precipitation • Particle size and filterability/washability – the larger the better • The factor determining the particle size: relative supersaturation = (Q-S)/S (where Q is the concentration of the supersaturated solution) • Nucleation • Particle growth • At large relative supersaturation the rate of nucleation is large – large naumber of small particles are formed • At small relative supersaturation the particle growth dominates, large particles are formed ☺ • In practice: elevate temperature to increase solubility, use dilute solution (to decrease Q) and add the precipitating agent slowly and under vigorous stirring Filtration and washing • Filtration may happen on paper filter or on glass filter • Mother liquor – is the liquid from which the precipitate is formed • Washing liquids – distilled water or water saturated with the precipitate • Peptization – is a process by which the precipitate returns to it dispersed state (behaves as a solution again) • Coprecipitation – soluble components other than the analyte are removed from the solution together g with the precipitate p p – surface adsorption – mixed crystal formation – occlusion and mechanical entrapment • Precipitation from homogeneous solution 20 Drying and weighing • Drying/ignition is necessary to obtain constant mass for the precipitate • Drying/ignition leads to the weighing form – the form of the analyte with accurately known composition (or stoichiometry) • Drying: t < 200 oC • Ignition: t = 6-800 oC • If the filtration is done with filter paper, ignition can be done, if glass filter is used, only drying is allowed • Weighing is always done by using an analytical balance • The weighed mass must always be larger than 100 mg (to have accuracy better than 1%) • Examples: SO42- ions in the form of BaSO4 Ca2+ ions in the form of Ca(COO)2.H2O Fundamentals of titrimetry A chemical reaction between the titrant solution and the analyte is suitable for titrimetry, titrimetry if 1. it takes place according to one kind of known stoichiometry 2. it is quantitative (conversion is > 99.9%, no excess of reactant is needed) 3. it isreasonably fast 4. completion of the reaction can be indicated 21 Terms used in titrimetry Standard solution – is a reagent of exactly known concentration that is used in the titrimetric analysis Titration – is a process in which the standard solution is added to the analyte until the reaction between the analyte and the reagent is complete Equivalence point – the point in the titration, when the amount of reagennt added to thge solution is exactly equivalent to the amount of the analyte (theoretical value) End point – the point in the titration, when a physical change occurs that is associated with the chemical equivalence q (practical value, this is what we obsrerve) Titration error – Et = Vep – Veq, where Vep is the actual volume of reagent required to reach the end point and Veq is the theoretical volume to reach the equivalence point Perfect titration - Vep = Veq, Terms used in titrimetry Titration curves – plot the reagent volume on the horizontal axis and some function of the analyte on the vertical axis; the equivalence point can be read off the titration curve; it can either be sigmoidal or linear segment curve. Indicators – they are added to the analyte solution to produce a visually observable physical change (usually colour change) at or very near to the equivalence point Primary standard – is an ultrapure compound that serves as a reference material for titrimetric method of analysis high purity atmospheric stability f hydrate y water absence of reasonable cost reasonable solubility large molar mass Secondary standard – a compound, whose purity has been established by chemical analysis and serves as a reference material for titrimetric method of analysis 22 Terms used in titrimetry Preparation of standard solutions: 1. Direct method – includes (i) accurate weighing of a 1 primary standard which is then (ii) dissolved in a suitable solvent and (iii) diluted to exactly known volume in a volumetric flask 2. Indirect method – includes preparation of the tatrant solution by approximate weighing and dilution to an approximately known volume, followed by standardization, which means (i) titrating a weighed quantity of a primary or a secondary standard or (ii) titrating a known volume of a standard solution 23