Analytical Chem Review and Q Fall Quantitative chemical Analysis

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1 Analytical Chem Review and Q Fall 2017 Quantitative chemical Analysis 1

2 Lecture Summary Fall 17 I. Classic Methods and Statistics 1 The Analytical Process 2. Measurements 3. Tools of the Trade 4. Statistics 5. Quality Assurance and Calibration Methods 26. Gravimetric Analysis 27. Sample preparation II. Equilibrium and Solutions 6. Chemical Equilibrium 7. Activity and the Systematic Treatment of Equilibrium 8. Monoprotic Acid-Base Equilibria 9. Polyprotic Acid-Base Equilibria 10. Acid-Base Titrations 11. EDTA Titrations III. Electrochemistry 13.. Fundamental of Electrochemistry 14. Electrodes and Potentiometry 15. Redox Titration 16. Electroanalytical Techniques IV. Spectroscopy 17. Fundamentals of Spectrophotometry 18. Applications of Spectrophotometry 19. Spectrophotometers 20. Atomic Spectroscopy --. NMR Spectroscopy 21. Mass Spectroscopy V. Chromatography 22. Introduction to Analytical Separation 23. Gas Chromatography 24. High-Performance Liquid Chromatography 2 August 2010

3 Experiment / Activities Fall 14 Experimental Methods 00 Glassware Calibration 1. Gravimetric Analysis of Barium Sulfate 2. EDTA titration of Zinc 3. Iodometric titration of vitamin-c 4. Cyclic Voltammetry of Vitamin-C 5. Atomic Absorption of Metal 6. Fluorimetry of Vitamin-C 7. Spectroscopic analysis of organic unknowns 8. Gas Chromatography of Hydrocarbons 9. Determination of Ascorbic Acid in Juices 10. LCMS of an unknown mixture Activities Basic Chem Penny Statistics Equilibrium, Acid-Base Electrochemistry Spectroscopy Analysis Chromatography 3

4 Chapter 0 Analytical Process Steps in analytical process from sample to Interpretation of results Formulating Question Selection the Method Sampling / Prep Sampling Solutions Homogeneous Heterogeneous Decant Analytes Quantitative transfer Mother liquor Supernatant Sample preparation Analysis Qualitative Analysis Quantitative Analysis Calibration Curves Standards Interference Masking Reporting and Interpretation Conclusions 4

5 Chapter 1 Measurements Units Solution Chemistry Concentrations Molarity Molality Weight percent Volume percent pph, ppm, ppb Solution Preparation Titration Apparatus for Titration Titration Calculations Stoichiometry 5

6 Chapter 2 Tools of the Trade Safety Lab notebook Equipment Analytical Balance Buret Volumetric Glassware Pipets, Syringe Technique Filtration Drying Conditioning Calibration Data Processing Excel 6

7 Chapter 3 Experimental Error Significant Figures Determination of Rules in math operation Type of Error Systematic vs Random Error Precision and Accuracy Statistical treatment of errors 7

8 Chapter 4 Statistics Gaussian Distribution Confidence Intervals Student's t 3 Case example Comparison of Std deviation (F-Test) Q-Test (for bad data) Method of Least Square LINEST with Excel 8

9 Chapter 5 Quality Assurance and Calibration Methods Basic Quality Assurance Method Validation Standard Addition Internal Standard 9

10 Chapter 6 Chemical Equilibrium Equilibrium Constant Equilibrium and Thermodynamics Solubility Products Common Ion Effect Separation by precipitation Complex formation Acid Base Conjugate pairs ph Strength of Acid Base Strong Acid Base Weak Acid Base Polyprotic Acid Base Ka, Kb relationship Solving Equilibrium problems using ice Table 10

11 Chapter 8 Activity and the Systematic Treatment of Equilibrium Effect of Ionic Strength on Solubility Activity Coefficients Extended Debye-Huckel Equation Using chart How to Interpolate ph with activities Systematic treatment of Equilibrium 11

12 Chapter 9 Monoprotic Acid-Base Equilibria Strong Acids and Bases The problem and cure Weak Acids and Bases Equilibria with Weak Acids and Bases Using Appendix G Buffers 12

13 Chapter 10 Polyprotic Acid-Base Equilibria Diprotic Acid and Bases Diprotic Buffers Polyprotic Acids and Bases Fractional or percent dissociation (a) 13

14 Chapter 7/11 Acid-Base Titrations Titration of Strong Acid Base Titration curve (stoichiometry) Titration of Weak Acid with Strong Base Titration curve 4-type calculations Titration of Weak Base with Strong Acid Titration curve 4-type calculations Titration curve of Diprotic system Titration curve features 14

15 Chapter 12 EDTA Titration Metal Chelate Complexes Ligand Metal coordinated bond Monodentate Ligands Polydentate Ligands Naming (Nomenclature) 15

16 Chapter 13 Advanced Topics in Equilibrium General Approach to Acid-Base Systems Activity approach Solubility and ph 16

17 Chapter 14 Fundamental of Electrochemistry Oxidation Reduction Reactions Lose electrons oxidation Gain electrons reduction Oxidizing and Reducing Agents Balancing redox reactions Redox Titrations Voltage and Free Energy Galvanic Cells (Battery) Components of a cell Standard Potentials Nernst Equation (relationship to Keq and Q) Working with activity. 17

18 Chapter 15 Electrodes and Potentiometry Reference Electrodes Indicator Electrodes Calculating ph from an Electrode. 18

19 Chapter 16 Redox Titration (See chapter 14) Balance redox chem eqn: Solve problem using stoichiometric strategy. Q: g Fe (Fe +2 ) ore requires ml of M KMnO 4. How pure is the ore sample? When iron ore is titrated with KMnO 4. The equivalent point results when: KMnO 4 (purple) g Mn 2+ (pink) Mn (+7) Mn(+2) Rxn: Fe +2 + MnO 4 - g Fe +3 + Mn 2+ Bal. rxn: 5 Fe 2+ + MnO H + g 5 Fe 3+ + Mn H 2 O Note Fe 2+ g 5 Fe 3+ : Oxidized Lose e- : Reducing Agent Mol of MnO 4 - = ml 0.180(mol/L) = mmol MnO 4 - Amt of Fe: = mmol 5 mol Fe g = g 1 mol MnO 4-1 mol Fe 2+ % Fe = ( g / g) 100 = 18.6 % 19

20 Chapter 17 Electroanalytical Techniques Electrolysis Predicting reaction Electrogravimetric Coulometry & Amperometry Voltammetry Basic Technique Electrode profile Cyclic Voltammetry Signal Analysis I pa, I pc, E pa, E pc, E pc/2, E pa/2, E 1/2, DE 20

Chapter 18 Fundamentals of Spectrophotometry Absorption of Light Measuring Absorbance Beer s Law Spectrophotometric Titration Luminescience Absorption Process is best described by Beer-Lambert Law

21 Chapter 18 Fundamentals of Spectrophotometry Absorption of Light Measuring Absorbance Beer s Law Spectrophotometric Titration Luminescience Absorption Process is best described by Beer-Lambert Law Attenuation (decrease of photon to sample) of a bean of radiation by an absorbing solution. The incoming radiation has a higher radiant power than the radiant transmitted by the solution. The path length of the absorbing solution is b and the concentration is c. A = -log T = log (Po/P) = e b c c = molarity (M) e = molar absorptivity L mol -1 cm -1 Optimal A:

22 Chapter 19 Applications of Spectrophotometry Analysis of Mixture Measuring Equilibrium Constant Method of Continuous Variation 22

23 Spectrophotometers Chapter 20 Lamps and Laser: Sources of Light Monochromators Detector Optical Sensor Fourier Transformed Infrared Spectroscopy Signal and Noise 23

24 Atomic Spectroscopy Chapter 21 Overviews Atomization: Flames, Furnaces and Plasmas Temperature influence Instrumentation Interference ICP, MS Atomic absorption spectroscopy (AAS) determines the presence of metals in liquid samples. It also measures the concentrations of metals in the samples, with concentrations range in the low mg/l range (ppm). In their elemental form, metals will absorb ultraviolet light when they are excited by heat. Each metal has a characteristic wavelength that will be absorbed. The AAS instrument looks for a particular metal by focusing a beam of UV light at a specific wavelength through a flame and into a detector. The sample of interest is aspirated into the flame. If that metal is present in the sample, it will absorb some of the light, thus reducing its intensity. The instrument measures the change in intensity. A computer data system converts the change in intensity into an absorbance. 24

25 Basic NMR Basic Theory Instrumentation Experiments NMR The interpretation of a 1H spectra depends on three features: chemical shifts, multiplicities and integrated peak area. Note the presence or absence of saturated structures, most of which gives resonances between 0 and 5 d ppm. Note the presence or absence of unsaturated structures in the region between 5 & 9 d ppm. Alkene protons resonate between 5 and 7 d ppm and aromatic protons between 7 and 9 d ppm. Note that alkyne protons resonance upfield around 1.5 d ppm. Note any very low field resonance (downfield) between 9 and 16 d ppm, which are associated with aldehydic and acidic protons, especially those involving in H- bonding. Measure the integrals, if recorded, and calculate the number protons in each resonance signal. Check for spin-spin splitting patterns given by adjacent alkyl group according to the n+1 rule and Pascal s triangle. Note that the position of the lower field multiplet of the two is very sensitive to the proximity of electronegative elements and groups such as O, CO, COO, OH, X, NH 2, etc.) Examine the splitting pattern given by aromatic protons, which couple around the ring and are often complex due to second order effects. 1,4 and 1,2-disubstituted rings give complex but symmetrical looking patterns of peaks, whereas mono- 1,3-and tri-substituted rings give more complex asymmetric patterns. Note any broad single resonance, which are evidence of liable protons from alcohols, phenols, acids and amines that can undergo slow exchange with other labile protons. 25

26 Mass Spectroscopy Chapter 22 Overviews What is Mass Spcctroscopy Type of Mass Spectroscopy Instrumentation Chromatography Mass Spectroscopy 26

27 Chapter 23 Introduction to Analytical Separation Solvent Extraction What is Chromatography? Plumber s View Efficiency of Separation Band spread 27

28 Chapter 24 Gas Chromatography Separation process in GC Techniques Sample Injection Detectors Sample Preparation Method Development 28

29 Chapter 25 High-Performance Liquid Chromatography Chromatography process Instrumentation Injection and Detection in HPLC Method Development and Reverse-Phase Separation Gradient Separation 29

30 Chapter 26 Skip This Section Chromatographic Methods & Capillary Electrophoresis Ion-Exchange Chromatography Ion Chromatography Molecular Exclusion Chromatography Affinity Chromatography Hydrophobic Interaction Chromatography Principle of Capillary Electrophoresis Conducting Capillary Electrophoresis 30

31 Glassware Calibration Glassware Calibration Eppendorf Calibration Reading volumes Calibration of Pipet Buoyancy Correction Buret calibration 31

32 Gravimetric Analysis Gravimetric Analysis Calculate Concentration of ions in an unknown solution 1 molarity 2 molality 3 ppm of Ba +2 in unknown solution. 4 % concentration of solution 32

33 Iodometric Titration Iodometric Titration Redox equations and stoichiometry Molarity of Vitamin-C based on titration curve g/l calculation % Vitamin-C in unknown 33

34 Zn Analysis in Supplement Standard Addition Concentration Calculations Dilution equation and dilution Factor Standard Addition equation Molarity of metal based on standard additon equation g/l in stoick solution calculation % metal in sample Std Addition Eqn: Conc Analyte (unknown) Conc Analyte (+ std in mixure) = Signal (unknown) Signal (mixture) [X] i [X] f +[S] f = I X I S+X 34

35 Cyclic Voltammetry Cyclic Voltammetry of Ferrocene Signal Analysis of voltammogram 35

36 Sample Questions 1. A serum sample is analyzed for iron by UV-Vis spectroscopy using standard addition. Two mL aliquots are added to 5.000mL portion of water. One of these portion contains 10.0ml of M Fe +2. The total volumes of both solutions are 5.500ml. The net absorbances are and respectively. What is the iron concentration in the serum. 2. HCN is a weak acid with a Ka = 4.9 e-10. A ml sample of a 0.25 M solution of this acid is titrated with a 0.50 M NaOH solution. What is the ph of the solution: a. after 15 ml of NaOH is added? b. at the equivalence point? 3 Calculate the % relative error in silubility by using concentrations instead of activities for CuCl (a Cu+ = 0.3nm) in M KNO3. 4 i) Agree Disagree: In the reaction: Br2 + H2O well as an oxidizing agent. BrO3- + Br, Br 2 can act as a reducing agent as ii) Agree Disagree: The pkb of H 2 PO 4 - is iii) Agree Disagree: The solubility of lead(ii) hydroxide will increase with the addition of ammonia (NH 3 ). iv) Agree Disagree: The ionic strength of 0.40M iron(ii) sulfate is higher than 0.20M (NH4) 2 CrO 4. 36

37 Sample Questions 1. A serum sample is analyzed for iron by UV-Vis spectroscopy using standard addition. Two mL aliquots are added to 5.000mL portion of water. One of these portion contains 10.0µl of M Fe+2. The total volumes of both solutions are 5.500ml. The net absorbances are and respectively. What is the iron concentration in the serum. Sample 1 V T = 5.500mL A 1 = Sample 2 V T = 5.500mL A 2 = Sample 2 Contains, L of M Fe +2 Iron Standard, C std = L * mol L * L = * 10 5 M A =εbc A 1 A 2 = C 1 C 2 = C 1 C 1 + C std, Rearranging - A 1 A 2 = C 1 C 1 + C std A 1 (C 1 + C std ) = C 1 (A 2 ) C 1 A 1 + C std A 1 = C 1 A 2 C std A 1 = C 1 A 2 - C 1 A 1 C 1 (A 2 - A 1 ) = A 1 C std C 1 = A 1 C std (A 2 - A 1 ) C 1 = A 1 C std (A 2 - A 1 ) = * 10 5 M = * 10 5 M ( ) = * 10 4 M Conc. Fe +2 in serum; C 1 V 1 = C serum V 0.500ml Aliquot C serum = C 1 V * 10 4 M L = V 0.500ml Aliquot L C serum = * 10 4 M L L = * 10 3 M= 1.23 * 10 3 M 37

38 Sample Questions 2. HCN is a weak acid with a Ka = A ml sample of a 0.25 M solution of this acid is titrated with a 0.50 M NaOH solution. What is the ph of the solution: a. after 15 ml of NaOH is added? b. at the equivalence point? a) Molarity HCN = 0.25M, Ka HCN = , ml The HCN will neutralize with 0.50 M NaOH HCN + OH -! H 2 O + CN - s 12.50mmol 7.5mmol R - 7.5mmol - 7.5mmol + 7.5mmol + 7.5mmol f 5 mmol mmol + 7.5mmol Buffer : ph = pka + log C b C a ph = = log 7.5 mmol 5 mmol b) At equivalent point HCN + OH -! H 2 O + CN - s mmol mmol R mmol mmol mmol mmol f 0 mmol 0 mmol mmol mmol Vol Tot = ml = 75 ml K b = CN - + H 2 O! OH - + HCN i M R - x - + x + x f M - x 0 mmol + x + x [OH - ] = (0.1667) = M poh = ph =

39 Sample Questions 3 Calculate the % relative error in solubility by using concentrations instead of activities for CuCl (a Cu+ = 0.3nm) in M KNO 3. %Error, CuCl K sp = , α Cu+ = 0.30 with activity : µ = 1 2 [.05 (1) (1) 2 ] = (1) 2.05 log γ = = 1 + ( / 305 γ Cu+ = , γ Cl- = = Molar Solubility (with activities) K sp = = [α Cu + ] [α Cl -] = [Cu + ] γ Cu + [Cl - ] γ Cl γ Cu +γ Cl - = (.9784) (.805) = [Cu + ][Cl - ] = s 2 s = = %Error, CuCl K sp = MolarSolubility (no activities) K sp = = (s) 2 s = = %Error = %Error = % 39

40 Sample Questions 4 i) Agree Disagree: In the reaction: Br 2 + H 2 O BrO Br, Br 2 can act as a reducing agent as well as an oxidizing agent. Agree, Br 2 is oxidixed in BrO 3- to a +5 oxidation state and then Br 2 is reduced to a -1 oxidation state in Br-. ii) Agree Disagree: The pkb of H 2 PO 4 - is The pka1 for H3PO4 equals to 7.11 e-3. That means K b3 = Kw / Kaq = 1e-14 / 7.11e-3 = e-12 or pk b = ~ iii) Agree Disagree: The solubility of lead(ii) hydroxide will increase with the addition of ammonia (NH 3 ). Pb(OH)2 dissociates to Pb+2 and 2 OH-. NH3 reacts with water to form NH4+ + OH-. The common ion will decrease the solubility of lead(ii) hydroxide. iv) Agree Disagree: The ionic strength of 0.40M iron(ii) sulfate is higher than 0.20M (NH4) 2 CrO 4. Ionic strength µ = 0.5 [ Conc1 M^2 + Conc2 X^2] FeSO4, µ = 0.5 [ 0.40 (2) (2)2] = 1.6 (NH4)2CrO4 µ = 0.5 [ 0.40 (1) (2)2] = 0.6 True, the ionic strength of iron(ii)sulfate is higher than ammonium chromate. 40

Analytical Topics to Consider in preparation for the MFAT/GRE

Analytical Topics to Consider in preparation for the MFAT/GRE 1. Solutions and Measurement: (CHEM 222) Describe the steps in a chemical analysis Know the meaning and use of the following solution concentrations.