Titration ppt

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This is useful to the chemical analysis persons. Tittration is one of the basic and standard method for quantitative chemical analysis. This describs the principles of titration, function of indicators, calculation of errors etc.

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Titration ppt

  1. 1. Principles of Titration and errors By Dr. A. Amsavel
  2. 2. IntroductionVolumetric analysis Simple and easy Fast and can be done on site Less expensive Estimation of content or Assay Precise and accurate  Depends on method and specificity
  3. 3. Requirements of a Titration Reaction Reaction must complete by 99.9 % so that < 0.1 % (or 1 ppt) remains unreacted Rxn must be rapid Titration needs to be performed in a reasonable time period The stoichiometry must be well defined, and known It can be predicted from equilibrium constants A method must be available to determine the equivalence point
  4. 4. Types of Titration1) Precipitation – A(aq) + B(aq) = AB(s)2) Acid-Base rxn – H+ + OH¯ = H2O (strong acids or bases) – HA + OH¯ = H2O + A¯ (weak acids) – A¯ + H+ = H2O + HA (weak bases)3) Complexation rxn – Zn2+ + 4NH3 = Zn(NH3)42+4) Redox rxn (oxidation-reduction) – Fe2+ + Ce4+ = Fe3+ + Ce3+
  5. 5. Standards• Measurements are made with reference to standards – The accuracy of a result is only as good as the quality and accuracy of the standards used – A standard is a reference material whose purity and composition are well known and well defined• Primary Standards – Used as titrants or used to standardize titrants – Requirements • Usually solid to make it easier to weigh • Easy to obtain, purify and store, and easy to dry • Inert in the atmosphere • High formula weight so that it can be weighed with high precision
  6. 6. Endpoint DetectionIt is critical, to know the completion of reaction / determination1) Visual indicators: • Observe a colour change or precipitation at the endpoint. – Rxn progress checked by addition of external or self indicator2) Photometry: • Use an instrument to follow the colour change or precipitation3) Electrochemistry: • Potentiometry - measure voltage change ( pH electrode) • Amperometry - measure change in current between electrodes in solution • Conductance – measure conductivity changes of solution Later two used for coloured, turbid, end point accurate
  7. 7. Acid-base titration Neutralization titration Neutralization Indicators Indicators & mixed indicators Neutralization curve Non-aqueous titration
  8. 8. Acids & BasesAcids:  Arrhenius acid: Any substance that, when dissolved in water, increases the concentration of hydronium ion (H3O+)  Bronsted-Lowry acid: A proton donor conjugate base  Lewis acid: An electron acceptorBases:  Arrhenius base: Any substance that, when dissolved in water, increases the concentration of hydroxide ion (OH-)  Bronsted-Lowery base: A proton acceptor conjugate acid  Lewis acid: An electron donor
  9. 9. Brønsted-Lowry Theory of Acids & BasesThe conjugate acid of a base is the base plus theattached proton and the conjugate base of an acid isthe acid minus the proton p. 507
  10. 10. Lewis Theory of Acids & Bases p. 506
  11. 11. pH calculationQ1: Calculate the pH of a solution if [H+] = 2.7 x 10-4 M pH = -log[H+] pH = -log(2.7 x 10-4) = 3.57Q2: Find the hydrogen ion concentration of a solution if its pH is 11.62. [H+] = 10-pH [H+] = 10-11.62 = 2.4 x 10-12MQ3: Find the pOH and the pH of a solution if its hydroxide ion concentration is 7.9 x 10-5M pOH = -log[OH-] pOH = -log(7.9 x 10-5) = 4.10 pH + pOH = 14 pH = 14 - 4.10 pH = 9.9
  12. 12. A solution with a pH of 1 has [H+] of 0.1 mol/L or 10-1A solution with a pH of 3 has [H+] of 0.001 mol/L or 10-3
  13. 13. pH of solutionsStomach juice: pH = 1.0 – 3.0 Human blood: pH = 7.3 – 7.5Lemon juice: pH = 2.2 – 2.4 Seawater: pH = 7.8 – 8.3Vinegar: pH = 2.4 – 3.4 Ammonia: pH = 10.5 – 11.5Carbonated drinks: pH = 2.0 – 4.0 0.1M Na2CO3: pH = 11.7Orange juice: pH = 3.0 – 4.0 1.0M NaOH: pH = 14.0
  14. 14. ENDPOINT = POINT OF NEUTRALIZATION = EQUIVALENCE POINT MOLES OF ACID = MOLES OF BASE
  15. 15. Ka and KbThe equilibrium constant for a Brønsted acid isrepresented by Ka, and base is represented by Kb.CH3COOH(aq) + H2O(l) H3O+(aq) + CH3COO–(aq) [H3O+][CH3COO–] Notice that H2O is not Ka = ––––––––––––––––– included in either [CH3COOH]equilibrium expression. NH3(aq) + H2O(l) NH4+(aq) + OH–(aq) [NH4+][OH–] Kb = ––––––––––––– pH of 1M AcoH =2.4 [NH3]
  16. 16. NaOH Titration curve: HCl Vs NaOH solution 1 M Sol 0.1M sol 14.0Vol ml pH pH 0.0 0.0 1.0 12.0 50.0 0.5 1.5 75.0 0.8 1.8 10.0 90.0 1.3 2.3 98.0 2.0 3.0 8.0 99.0 2.3 3.3 99.5 2.6 3.6 6.0 99.8 3.0 4.0 99.9 3.3 4.3 4.0 100.0 7.0 7.0 100.1 10.7 9.7 2.0 100.2 11.0 10.0 100.5 11.4 10.4 0.0 0.0 20.0 40.0 60.0 80.0 100.0 120.0 140.0 160.0 101.0 11.7 10.7 Series1 Series2 102.0 12.0 11.0 110.0 12.7 11.7 150.0 13.3 12.3
  17. 17. Titration curveNaOH 1 M Sol 0.1M solVol ml pH pH 98.0 2.0 3.0 99.0 2.3 3.3 99.5 2.6 3.6 99.8 3.0 4.0 99.9 3.3 4.3 100.0 7.0 7.0 100.1 10.7 9.7 100.2 11.0 10.0 100.5 11.4 10.4 101.0 11.7 10.7 102.0 12.0 11.0
  18. 18. Titration Curve: Strong Acid with Strong Base At the equivalence point in an acid–base titration, the acid and base have been brought together in precise stoichiometric proportions. (Endpoint) Bromphenol blue, bromthymol blue, and phenolphthalein all change color at very nearly 20.0 mL At about what volume would we see a color change if we used methyl violet as the indicator?
  19. 19. Titration Curve: Weak Acid with Strong Base The equivalence-point pH is NOT 7.00 here. Why not?? Bromphenol blue was ok for the strong acid/strong base titration, but it changes color far too early to be useful here.
  20. 20. Acid–Base Indicators An acid–base indicator is a weak acid or base. The acid form (HA) of the indicator has one color, the conjugate base (A–) has a different color. One of the “colors” may be colorless. In an acidic solution, [H3O+] is high. Because H3O+ is a common ion, it suppresses the ionization of the indicator acid, and we see the color of HA. In a basic solution, [OH–] is high, and it reacts with HA, forming the color of A–.
  21. 21. Function of IndicatorsExample: phenolphthalein Near pH 8, Indicator dissociates and gives red base Human eye can detect it as a pink tinge at that pH Indicators must be carefully chosen so that their colour changes take place at the pH values expected for an aqueous solution of the salt produced in the titration.
  22. 22. Basis of Indicator selectionIndicator colour change, from acid pH pKind example of titration useto alkali range weak base - strong acidMethyl orange, (red ==> yellow) 3.7 3.1-4.4 titration e.g. ammonia titrated with hydrochloric acidBromophenol blue, (yellow ==> weak base - strong acid 4.0 2.8-4.6blue) titration weak base - strong acidMethyl red, (red ==> yellow) 5.1 4.2-6.3 titration strong acid - strong baseBromothymol blue, (yellow ==> 7.0 6.0-7.6 titration e.g. hydrochloric acidblue) <=> sodium hydroxide titration strong acid - strong basePhenol red, (yellow ==> red) 7.9 6.8-8.4 titration e.g. hydrochloric acid <=> sodium hydroxide titrationThymol blue (base form), (yellow weak/strong acid - strong base 8.9 8.0-9.6==> blue) titration weak acid - strong basePhenolphthalein, (colourless ==> 8.3- 9.3 titration e.g. ethanoic acidpinky-red) 10.0 titrated with sodium hydroxide
  23. 23. Colours of indicator at different pH
  24. 24. Indicators: Color changes against pH
  25. 25. Non-Aqueous Titration Theory is same as acid-Base titration Reaction carry out in non-aqueous medium Applied where  Material which are not soluble in water  Week acid and bases are titrated  Poor end point in water mediumPrinciple based on Brønsted-Lowry Theory
  26. 26. Brønsted-Lowry TheoryThe conjugate acid of a base is the base plus theattached proton and the conjugate base of an acid isthe acid minus the proton p. 507
  27. 27. Solvents used in NATSolvents used can be classified as four types: Aprotic solvents: Chemically neutral  Eg. Toluene, carbon tetrachloride Protogenic solvents: Acidic nature readily donate protons,  Eg. Anhyd. HF, H2SO4 Amphiprotic solvent: Which are sly ionize and donate and accept protons,  Eg Alcohols, weak organic acids.  Acetic acid makes weak acid into storing base Protophilc solvents: Posses high affinity for protons.  Eg. Liq ammonia, Amine, Ketones  Increases the acidic strength
  28. 28. Selection of Solvents for NATAcetic acid used for titration of weak bases, Nitrogen containing compoundsAcetonitrile / with ACOH: Metal ethanoatesAlcohols (IPA, nBA) : Soaps and salts of organic acids,DMF: Benzoic acid, amides etc
  29. 29. Titrants for NAT Perchloric acid in acetic acid  Amines, amine salts, amino acids, salts of acids Potassium Methoxide in Toluene- Methanol Quan ammonium hydroxide in Acetonitrile- pyridine  Acids, enols, imides & sulphonamides
  30. 30. Indicators for NAT Principle is similar to acid base titrationIndicators: Crystal violet, Methyl red, Thymol blue, & 1-Naphthaol benzein
  31. 31. Calculationo Normality: Eq.wt/1000ml or meq/mLo Morality: Mole/1000mlo V1 N1 = V2N2o N1 = V2N2/V1 Normality = Wt of sample x 1000 / Eq. Wt x V Wt of sample (mg) = V x N x Eq. wt Assay = Qty estimated in sample x 100/ wt of sample Assay = V x N x Eq. wt x 100/ wt of sample x 1000
  32. 32.  1 ml of 1N HCl = 0.04g of NaOH (40/1000) 1 ml of 0.1N HCl = 0.004g of NaOH
  33. 33. Titration ErrorError in methods: The endpoint method may not show a change exactly at the equivalence point due to the reactions involved Titration Error = Vol at endpoint - Vol at equivalence point• Negative error & Positive error means endpoint is early - before equiv point or late after equiv point
  34. 34. Errors in Volume and weight: 10 ml titre volume = 100 %  If difference is 0.1ml error is 1%; 0.2ml = 2% 5ml titre volume = 100 %  0.1ml = 2% error Optimum level is about 25ml 25 ml titre volume = 100 %  0.1ml = 0.4% error
  35. 35. Volumetric apparatusUSP: Burette selection:  NLT 30% nominal volume (15ml in 50ml burette)  Micro burette for < 10ml Limit of error:  Volumetric flask: 25ml, 50ml, 100ml is 0.03, 0.05& 0.08ml  Pipets:5, 10, 25 ml is 0.01, 0.02 &0.03ml  Burets:10, 25, 50ml is 0.02, 0.1&0.1ml Tips: out flow NMT 500uL per second
  36. 36. Operational & personal errorList several of the variables involved in correctly using a 10mL volumetric pipette. drain time; possible beads on the inner surface temperature; bringing meniscus to the proper level; angle of drain; touching off last drop; rinsing of the pipet with the solution used; Pipet calibration; etc.
  37. 37. Error in weighing can occur.• Misreading of the balance,• Balance not level,• Not cleaning the surface of the balance first,• Touching the weighed object with moist hands,• Leaving the balance doors open during weighing,• Using a miscalibrated balance,• Not cooling the sample down to near room temperature,• Not removing a static charge from the sample,• Excess vibration or air currents from people or nearby equipment, and• Prolonged time sample left on pan adds/loses moisture.
  38. 38. Possible contaminationAn analyst could contaminate a sample during weighing by placing a contaminated spatula placing the sample on or into a contaminated holder during weighing, dropping some lint/hair/skin or sneeze into the sample while weighing, opening up a bottle of chemicals near the sample being weighed. When performing trace analysis, it is possible for just a microgram even massive fingerprint!
  39. 39. Units of measurementName Defining UnitsMolarity moles of solute/liter (solutions), or(e.g. 0.1200 M) millimoles/milliliter (solutions) (grams of substance/grams of sample) xPercent 100%, or(e.g. 23.45 %) centigrams/gram (seldom used) milligrams/liter (solutions), orParts per million micrograms/milliliter (solutions)(e.g 2.34 ppm, 2.34 milligrams/kilogram (solids), ormg/L) micrograms/gram (solids)Parts per billion micrograms/liter (solutions), or(e.g. 0.45 ppb, 0.45 nanograms/gram (solids)ug/L)
  40. 40. Oxidation- reduction titration Oxidation-reduction reaction Reduction potential is calculated by  Nernst equation • E1= E’ + 0.591/n log (ox)/(red) • E=(E1+E2)/2 Equivalence point by redox potential Vs Volume Indicator selection
  41. 41. Precipitation titration Reagents used id based on Solubility products of precipitate Titration curve: -log Conc. Of ion Vs Volume Concentration of ions  Eg. Ksol(Agcl) = Ag + Cl Indicator:  Formation of coloured compound (ppt/complex)  Adsorption indicators
  42. 42. Complexation titration M + EDTA M(EDTA)  Complex formation depend on Stability constant, pH, titration curve pM Vs, Vol of EDTA Indicators (Metal / metal ion indicators):  M-ln + EDTA M(EDTA) + In
  43. 43. Types of Complexation titration Back titration Masking Selective de-masking Separation by ppt and solvent extraction Application, almost metals,
  44. 44. An Equation for Buffer Solutions In certain applications, there is a need to repeat the calculations of the pH of buffer solutions many times. This can be done with a single, simple equation, but there are some limitations. The Henderson–Hasselbalch equation: [conjugate base] pH = pKa + log –––––––––––––– [weak acid]• To use this equation, the ratio [conjugate base]/[weak acid] must have a value between 0.10–10 and both concentrations must exceed Ka by a factor of 100 or more.
  45. 45. The Common Ion Effect Consider a solution of acetic acid. If we add acetate ion as a second solute (i.e., sodium acetate), the pH of the solution increases: LeChâtelier’s principle: What happens to [H3O+] when the equilibrium shifts to the left?

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