Analytical Chem Review and Q Fall...

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Analytical ChemReview and Q

Fall 2017

Quantitative chemical Analysis

2 August 2010

Lecture Summary Fall 17

I. Classic Methods and Statistics1 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 Solutions6. 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. Electrochemistry13.. Fundamental of Electrochemistry

14. Electrodes and Potentiometry

15. Redox Titration

16. Electroanalytical Techniques

IV. Spectroscopy17. Fundamentals of Spectrophotometry

18. Applications of Spectrophotometry

19. Spectrophotometers

20. Atomic Spectroscopy

--. NMR Spectroscopy

21. Mass Spectroscopy

V. Chromatography22. Introduction to Analytical Separation

23. Gas Chromatography

24. High-Performance Liquid Chromatography

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Experiment / Activities Fall 14

Experimental Methods00 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

ActivitiesBasic Chem

Penny Statistics

Equilibrium, Acid-Base

Electrochemistry

Spectroscopy Analysis

Chromatography

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Chapter 0Analytical Process

• Steps in analytical process from sample to Interpretation of results

Formulating

Question

Selection the

MethodSampling / Prep

• Sampling

• Solutions

HomogeneousHeterogeneous

• Decant• Analytes

• Quantitative transfer• Mother liquor

• Supernatant• Sample preparation

Analysis

• Qualitative Analysis

• Quantitative Analysis

• Calibration Curves

• Standards

• Interference

• Masking

Reporting and Interpretation

Conclusions

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Chapter 1Measurements

• Units

• Solution ChemistryConcentrations

MolarityMolalityWeight percentVolume percentpph, ppm, ppb

Solution Preparation• Titration

Apparatus for TitrationTitration Calculations

• Stoichiometry

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Chapter 2Tools of the Trade

• Safety

• Lab notebook

• EquipmentAnalytical BalanceBuretVolumetric GlasswarePipets, Syringe

• TechniqueFiltrationDryingConditioningCalibration

• Data ProcessingExcel

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Chapter 3Experimental Error

• Significant FiguresDetermination ofRules in math operation

• Type of ErrorSystematic vs Random ErrorPrecision and AccuracyStatistical treatment of errors

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Chapter 4Statistics

• Gaussian Distribution

•Confidence Intervals

• Student's t3 Case example

• Comparison of Std deviation (F-Test)

•Q-Test (for bad data)

• Method of Least SquareLINEST with Excel

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Chapter 5Quality Assurance and Calibration Methods

• Basic Quality Assurance

• Method Validation

• Standard Addition

• Internal Standard

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Chapter 6Chemical Equilibrium

• Equilibrium Constant

• Equilibrium and Thermodynamics

• Solubility ProductsCommon Ion EffectSeparation by precipitation

• Complex formation

• Acid BaseConjugate pairspHStrength of Acid Base

Strong Acid – BaseWeak Acid – Base

Polyprotic Acid BaseKa, Kb relationship

• Solving Equilibrium problems using iCe Table

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Chapter 8Activity and the Systematic Treatment of Equilibrium

• Effect of Ionic Strength on Solubility

• Activity CoefficientsExtended Debye-Huckel EquationUsing chart

How to Interpolate

• pH with activities

•Systematic treatment of Equilibrium

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Chapter 9Monoprotic Acid-Base Equilibria

• Strong Acids and BasesThe problem and cure

• Weak Acids and BasesEquilibria with Weak Acids and Bases

Using Appendix G

• Buffers

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Chapter 10Polyprotic Acid-Base Equilibria

•Diprotic Acid and Bases

•Diprotic Buffers

•Polyprotic Acids and Bases

•Fractional or percent dissociation (a)

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Chapter 7/11Acid-Base Titrations

•Titration of Strong Acid BaseTitration curve (stoichiometry)

• Titration of Weak Acid with Strong BaseTitration curve4-type calculations

• Titration of Weak Base with Strong AcidTitration curve4-type calculations

Titration curve of Diprotic systemTitration curve features

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Chapter 12EDTA Titration

• Metal Chelate ComplexesLigand – Metal coordinated bond

Monodentate LigandsPolydentate Ligands

• Naming (Nomenclature)

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Chapter 13Advanced Topics in Equilibrium

• General Approach to Acid-Base SystemsActivity approachSolubility and pH

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Chapter 14Fundamental of Electrochemistry

• Oxidation Reduction ReactionsLose electrons oxidationGain electrons reductionOxidizing 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.

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Chapter 15Electrodes and Potentiometry

• Reference Electrodes

• Indicator Electrodes

• Calculating pH from an Electrode.

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Chapter 16Redox Titration (See chapter 14)

Balance redox chem eqn: Solve problem using stoichiometric strategy.

Q: 1.225 g Fe (Fe+2) ore requires 45.30 ml of 0.0180 M KMnO4. How pure is the ore sample?

When iron ore is titrated with KMnO4 . The equivalent point results when:KMnO4 (purple) g Mn2+ (pink)

Mn (+7) Mn(+2)

Rxn: Fe+2 + MnO4- g Fe+3 + Mn2+

Bal. rxn: 5 Fe2+ + MnO4- + 8 H+ g 5 Fe3+ + Mn2+ + 4 H2O

Note Fe2+ g 5 Fe3+ : Oxidized Lose e- : Reducing Agent

Mol of MnO4- = 45.30 ml • 0.180(mol/L) = 0.8154 mmol MnO4-

Amt of Fe: = 0.8154 mmol • 5 mol Fe+2 • 55.8 g = 0.2275 g1 mol MnO4- 1 mol Fe2+

% Fe = (0.2275 g / 1.225 g) • 100 = 18.6 %

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Chapter 17Electroanalytical Techniques

•ElectrolysisPredicting reaction

•Electrogravimetric

• Coulometry & Amperometry

• VoltammetryBasic TechniqueElectrode profileCyclic Voltammetry

Signal Analysis Ipa, Ipc, Epa, Epc, Epc/2, Epa/2, E1/2, DE

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Chapter 18Fundamentals of Spectrophotometry•Absorption of Light•Measuring Absorbance

Beer’s LawSpectrophotometric 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: 0.4 - 0.9

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Chapter 19Applications of Spectrophotometry

•Analysis of Mixture

•Measuring Equilibrium ConstantMethod of Continuous Variation

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Chapter 20Spectrophotometers

•Lamps and Laser: Sources of Light

• Monochromators

• Detector

• Optical Sensor

• Fourier Transformed Infrared Spectroscopy

• Signal and Noise

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Chapter 21Atomic Spectroscopy•Overviews • Atomization: Flames, Furnaces and Plasmas

Temperature influenceInstrumentationInterference

• 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.

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NMRBasic NMR

• Basic Theory• Instrumentation• Experiments

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, NH2, 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.

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Chapter 22Mass Spectroscopy•Overviews • What is Mass Spcctroscopy

Type of Mass SpectroscopyInstrumentationChromatography

Mass Spectroscopy

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Chapter 23Introduction to Analytical Separation

•Solvent Extraction

• What is Chromatography?Plumber’s ViewEfficiency of SeparationBand spread

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Chapter 24Gas Chromatography

•Separation process in GC

• TechniquesSample InjectionDetectorsSample PreparationMethod Development

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Chapter 25High-Performance Liquid Chromatography

•Chromatography process

• InstrumentationInjection and Detection in HPLC

• Method Development and Reverse-Phase SeparationGradient Separation

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Chapter 26Skip 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

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Glassware Calibration

Glassware Calibration

• Eppendorf Calibration

• Reading volumes

• Calibration of Pipet

• Buoyancy Correction

• Buret calibration

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Gravimetric AnalysisGravimetric Analysis

• Calculate Concentration of ions in an unknown solution1 molarity2 molality3 ppm of Ba+2 in unknown solution.4 % concentration of solution

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

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Zn Analysis in SupplementStandard Addition•Concentration Calculations

Dilution equation and dilution Factor•Standard Addition equation

Molarity of metal based on standard additon equationg/L in stoick solution calculation% metal in sample

€

Std Addition Eqn:ConcAnalyte (unknown)

ConcAnalyte (+ std in mixure)

=Signal (unknown)

Signal (mixture)

[X] i

[X] f + [S] f

=I X

I S+X

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Cyclic Voltammetry

Cyclic Voltammetry of Ferrocene

•Signal Analysis of voltammogram

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Sample Questions1. A serum sample is analyzed for iron by UV-Vis spectroscopy using standard addition. Two 0.500-mL aliquots are added to 5.000mL portion of water. One of these portion contains 10.0ml of 0.0500M Fe +2 . The total volumes of both solutions are 5.500ml. The net absorbances are 0.321 and 0.581 respectively. What is the iron concentration in the serum.

2. HCN is a weak acid with a Ka = 4.9 • e-10. A 50.00 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 0.0500M KNO3.

4i) __ Agree __ Disagree:

In the reaction: Br2 + H2O ® BrO3- + Br – , Br2 can act as a reducing agent as well as an oxidizing agent.

ii) __ Agree __ Disagree:The pKb of H2PO4- is 11.88.

iii) __ Agree __ Disagree:The solubility of lead(II) hydroxide will increase with the addition of ammonia (NH3).

iv) __ Agree __ Disagree:The ionic strength of 0.40M iron(II) sulfate is higher than 0.20M (NH4)2 CrO4.

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Sample Questions1. A serum sample is analyzed for iron by UV-Vis spectroscopy using standard addition. Two 0.500-mL aliquots are added to 5.000mL portion of water. One of these

portion contains 10.0µl of 0.0500M Fe+2. The total volumes of both solutions are 5.500ml. The net absorbances are 0.321 and 0.581 respectively. What is the iron

concentration in the serum.

€

Sample1 VT = 5.500mL A1 = 0.321Sample2 VT = 5.500mL A2 = 0.581Sample2 Contains, 10•10-6L of 0.0500M Fe+2

Iron Standard, Cstd = 10•10-6L *0.0500 molL

* 1.0055L

= 9.0909 *10−5 M

A =εbc → A1

A2

=C1

C2

= C1

C1 + Cstd

, Rearranging -

A1

A2

= C1

C1 + Cstd

→ A1(C1 + Cstd ) = C1(A2 )

C1A1 + CstdA1 = C1A2 → CstdA1 = C1A2 - C1A1 → C1 (A2 - A1 ) = A1Cstd → C1 = A1Cstd

(A2 - A1 )

C1 = A1Cstd

(A2 - A1 ) = .321 •9.0909 *10−5 M

(.581 - .321)= 2.9182 *10−5 M

0.260 = 1.1224 *10−4 M

Conc. Fe+2 in serum; C1V1 = CserumV0.500ml Aliquot → Cserum = C1V1

V0.500ml Aliquot

=1.1224 *10−4 M•0.00550L

0.000500L

Cserum = 1.1224 *10−4 M•0.00550L0.000500L

= 1.2346 *10−3 M= 1.23 *10−3 M

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Sample Questions2. HCN is a weak acid with a Ka = 4.9 • 10-10. A 50.00 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) MolarityHCN = 0.25M, KaHCN =4.9 ⋅ 10-10, 50.00 mL The HCN will neutralize with 0.50 M NaOH HCN + OH- ! H2O + CN-

s 12.50mmol 7.5mmol 1•10-7 0R - 7.5mmol - 7.5mmol +7.5mmol +7.5mmol f 5 mmol 0 +7.5mmol +7.5mmol

Buffer :

pH = pKa + log Cb

Ca

= 9.31 + log 7.5 mmol5 mmol

pH = 9.486

€

b) At equivalent point HCN + OH- ! H2O + CN-

s 12.50 mmol 12.50 mmol 1•10-7 0R -12.50mmol -12.50 mmol +12.5 mmol + 12.5 mmol f 0 mmol 0 mmol +12.5 mmol + 12.5 mmol

VolTot = 50 + 25 mL = 75 mL Kb = 2.041•10-5

CN- + H2O ! OH- + HCNi 0.1667M - 1•10-7 0R - x - +x + x f 0.1667M - x 0 mmol +x + x

[OH-] = 2.041•10-5(0.1667) = 1.845•10-3 M

pOH = 2.734 pH = 11.266

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Sample Questions3 Calculate the % relative error in solubility by using concentrations instead of

activities for CuCl (a Cu+ = 0.3nm) in 0.0500M KNO3.

€

%Error, CuCl Ksp = 1.9 ⋅10-7

MolarSolubility (no activities)− Ksp = 1.9 ⋅10-7 = (s)2

s = 1.9 ⋅10-7 = 4.3589-04

%Error = 4.358910-04 − 4.9116 ⋅10-04

4.9116 ⋅10-04•100

%Error = 11.25 %

€

%Error, CuCl Ksp = 1.9 ⋅10-7, αCu+ = 0.30

with activity : µ = 12

[.05 (1)2 + .05 (1)2] = .05

log γ = -0.51 (1)2 .05

1 + (300 .05 / 305= -1.1404 ⋅10-2

1 + .2199= −9.3479 ⋅10-3

γCu+ = 0.9787, γCl- = 0.805

Molar Solubility (with activities)− Ksp = 1.9 ⋅10-7 = [αCu+ ] [αCl- ] = [Cu+] γCu+ [Cl-] γCl-

1.9 ⋅10-7

γCu+γCl-= 1.9 ⋅10-7

(.9784) (.805) = [Cu+][Cl-] = s2

s = 2.4124 ⋅10−7 = 4.9116 ⋅10-4

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Sample Questions4i)__ Agree __ Disagree:

In the reaction: Br2 + H2O ® BrO3- + Br – , Br2 can act as a reducing agent as well as an oxidizing agent.Agree, Br2 is oxidixed in BrO3

- to a +5 oxidation state and then Br2 is reduced to a -1 oxidation state in Br-.

ii) __ Agree __ Disagree:The pKb of H2PO4- is 11.88.

The pKa1 for H3PO4 equals to 7.11 e-3. That means Kb3 = Kw / Kaq = 1e-14 / 7.11e-3 = 1.406 e-12 or pKb = 11.85 ~ 11.88

iii) __ Agree __ Disagree:The solubility of lead(II) hydroxide will increase with the addition of ammonia (NH3).

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 CrO4.Ionic strength µ = 0.5 [ Conc1 • M^2 + Conc2 • X^2]FeSO4, µ = 0.5 [ 0.40 (2)2 + 0.40 (2)2] = 1.6(NH4)2CrO4 µ = 0.5 [ 0.40 (1)2 + 0.20 (2)2] = 0.6True, the ionic strength of iron(II)sulfate is higher than ammonium chromate.