Quantitative Analysis MD KHAJA ARIF Asst. Professor of Chemistry MVS Government Arts & Science College Mahabubnagar Mobile:

Quantitative Analysis MD KHAJA ARIF Asst. Professor of Chemistry MVS Government Arts & Science College Mahabubnagar E-Mail: khajahcu@gmail.com Mobile: +91 9052143116 1

Analytical chemistry Analytical chemistry is the science of obtaining, processing, and communicating information about the composition and structure of matter. In other words, it is the art and science of determining what matter is and how much of it exists. Analytical Chemistry deals with methods for Identification - of molecules / functional groups Structural Determination - determination of structure of a molecule Quantification - Amountpresent in a sample/mixture Qualitative analysis –What is present/Identity of species in a impure sample/mixture? Separation – Separation of mixtures–Chromatographic Techniques. 2

A nalytical M ethodology 1.Plan: Qualitative or quantitative or both; what kind of information have; which technique is suitable etc. 2.Sampling: Accuracy depends on proper sampling, characteristic of sample is very important, required good representative sample (from top, middle and bottom and mix up and take average sample). 3.Sample preparation: Depends on analytical technique. 4.Analytical measurement: 5.DataAnalysis:Whetherthedatamake sense or not. 3

Analytical Techniques Chemical Methods (Classical Method) Qualitative Analysis Quantitative Analysis Physical Methods (Instrumental Method) Spectroscopic Technique Electrochemical Technique Separation Techqnique Classification of Analytical Methods 4

I NTRODUCTION Qualitative Analysis (identification) provides information about the identity of species or functional groups in the sample (an analyte can be identified). Quantitative Analysis (Quantify) provides numerical information of analyte (quantitate the exact amount or concentration). 5

Qualitative analysis Macro analysis > 100 mg Semi micro analysis 10-100 mg Micro analysis < 10 mg 6

Quantitative Analysis Quantitative analysis Gravimetric Method Volumetric Method Acid Base titration Redox method Precipitation titrations Complexometric titrations 7

Spectroscopic Technique Absorption of EMR UV- Visible IR NMR X- Ray Spectroscopy Emissions of EMR Fluorescence Phosphoresces Flame Photometry Scattering Or Refraction of EMR Nepholometry Terbidometry Raman Spectroscopy Refrectometry Rotation of EMR Polarimetery 8

E lectro C hemical M ethods PotentiometrypH metryConductometryVoltametryPolarographyAmparometeryCoulombmetry 9

Separation Techniques Separation Technique Paper Chromatography Thin layer Chromatography (TLC) Column Chromatography HPLCGas Chromatography 10

C hemical M ethod (C lassical M ethod ) Gravimetric Method Volumetric Method Acid Base titrations (Neutralization method) Oxidation reduction titrations (Redox method) Precipitation titrations Complexometric titrations 11

Various Terms in Volumetric Analysis 12

Volumetric Analysis It is a general term for a method in quantitative chemical analysis in which the amount of a substance is determined by the measurement of the volume that the substance occupies. It is commonly used to determine the unknown concentration of a known reactant. Volumetric analysis is often referred to as titration, a laboratory technique in which one substance of known concentration and volume is used to react with another substance of unknown concentration. 13

Definitions 1.Volumetric Analysis: Involves the preparations, storage, and measurement of volume of chemicals for analysis. 2.Volumetric Titrimetry: Quantitative chemical analysis which determines volume of a solution of accurately known concentration required to react quantitatively with the analyte (whose concentration to be determined). The volume of titrant required to just completely react with the analyte is the TITRE. 14

3. Titration: A process in which a standard reagent is added to a solution of analyte until the reaction between the two is judged complete. 4. Standardization: A process to determine the concentration of a solution of known concentration by titrating with a primary standard 5. Standard solution: A reagent solution of accurately known concentration is called a standard solution. 15

Ideal standard solution: Conc. Does not change with time Quick to react – minimum time between additions Selective reaction with analyte Completely reacts Simple balanced equation  Primary standard solution: Compound of sufficient purity from which std. solution can be made by direct weighting followed by dilution.  Requirements: - Purity of at least 99.9% - Stable and easy to purify if necessary - Convenient to weigh - Well soluble and stable in solution - non-toxic, non-expensive  Ex: Anhy. Sodium carbonate( Na 2 CO 3 ), Pot. Dichromate (K 2 Cr 2 O 7 ), Mohr’s salt (FeSO 4.(NH 4 )2SO 4.6H 2 O, crystalline oxalic acid ( H 2 C 2 O 4.2H 2 O) 16

 Secondary standard solution: Substance from which std, solution cannot be prepared directly but whose solution can be determined accurately using a standard solution.  Process – standardization.  NaOH, H 2 SO 4, HCl, HNO 3, KMNO 4 etc. 17

6. End point: The point at which the reaction is observed to be completed is the end point The end point in volumetric method of analysis is the signal that tells the analyst to stop adding reagent and make the final reading on the burette. Endpoint is observed with the help of indicator 7. Equivalent point: The point at which an equivalent or stoichiometric amount of titrant is added to the analyte based on the stoichiometric equation. 18

Endpoint Detection Endpoint Detection is critical; it is to know the completion of reaction and accuracy of analysis; 1 Visual indicators: Observe a colour change or precipitation at the endpoint. Reaction completion is identified by addition of external or self indicator 2Photometry: Use an instrument to find out the colour change or precipitation 3Electrochemistry: Potentiometry :Measure the potential change ( pH electrode) Amperometry :Measure the change in current between electrodes in Reaction solution Conductance:Measure the conductivity changes of solution Later two method can be used for coloured, turbid solution and accurate end point 19

Equipment Volumetric analysis involves a few pieces of equipment: Pipette – for measuring accurate and precise volumes of solutions Burette – for pouring measured volumes of solutions Conical flask – for mixing two solutions Wash bottles – these contain distilled water for cleaning equipment Funnel – for transfer of liquids without spilling Volumetric flasks – a flask used to make up accurate volumes for solutions of known concentration 21

Reaction must be stoichiometric, well defined reaction between titrant and analyte. Reaction should be rapid. Reaction should have no side reaction, no interference from other foreign substances. Must have some indication of end of reaction, such as color change, sudden increase in pH, zero conductivity, etc. Known relationship between endpoint and equivalence point. Requirement for Volumetric Analysis 22

Titration procedure 23

Volumetric analysis methods There depending on the nature of the reaction between the analyte are two different ways of using a volumetric determination, and the titrant. 1.Direct titration: titrant is added to the analyte until the end point is reached. 2.Back titration *: a measured excess amount of a standard reagent is added to the analyte, followed by titration with a second standard reagent to determine the amount of unreacted first reagent. *Back titrations are useful whenever there is not a suitable indicator or the kinetics or equilibrium constant are not extremely favorable for the direct titration. 24

Acid and Base 25

The pH Scale The acidity/basicity of a solution is often expressed as pH. pH is defined as the negative of the common logarithm of the hydronium ion concentration. pH = −log [H 3 O + ] pH 7 is basic, pH = 7 is neutral. 26

The pOH Scale Another way of expressing the acidity/basicity of a solution is pOH. pOH is defined as the negative of the common logarithm of the hydroxide ion concentration. pOH = −log [OH - ] pOH 7 is acidic, pOH = 7 is neutral 27

General Knowledge: pH of various solutions Stomach juice: pH = 1.0 – 3.0 Lemon juice: pH = 2.2 – 2.4 Vinegar: pH = 2.4 – 3.4 Carbonated drinks: pH = 2.0 – 4.0 Orange juice: pH = 3.0 – 4.0 Human blood: pH = 7.3 – 7.5 Seawater: pH = 7.8 – 8.3 Ammonia: pH = 10.5 – 11.5 0.1M Na 2 CO 3 : pH = 11.7 1.0M NaOH: pH = 14.0 28

pH of Some Common Substances 29

Indicators Indicators are organic chemicals that change colour as the pH of a solution changes Above a certain pH value they have one colour, but have another colour when the pH is below that value Allow for the accurate determination of end points of acid/base reactions 30

Colours of indicator at different pH 31

Indicators: Color changes against pH 32

Indicators pH > 7 pink pH < 7 colourless 33

Phenolphthalein Phenolphthalein is an organic compound often used as an acid-base indicator. Phenolphthalein is colorless in acidic solutions, but turns pink when the pH is greater than 8.3. 34

Volumetric Method Acid Base titrations (Neutralization method) Oxidation reduction titrations (Redox method) Precipitation titrations Complexometric titrations 36

Acid Base titrations 37

Acid-Base Titration Titration is a method for determining the concentration of a solution by reacting a known volume of that solution with a solution of known concentration. The analyte is a measured volume of an acid or base of unknown concentration. The standard solution (titrant) is an acid or base solution whose concentration is known. Standard solution (titrant) analyte 38

Titration Procedure In a titration procedure, a measured volume of an analyte is placed in a beaker or flask, and initial pH recorded. The standard solution (titrant) is filled in a burette. A couple of drops of an acid-base indicator are added to the flask. The standard solution is slowly added to the unknown solution in the flask. As the two solutions are mixed the acid and the base are neutralized. 39

Titration Procedure (cont’d) As the base is added to the acid, H + reacts with OH – to form water. But there is still excess acid present so the color does not change. Once enough base has been added to neutralize all the acid, the indicator changes color. The difficulty is determining when there has been just enough titrant added to complete the reaction…without going over! 40

The Equivalence Point End point: The point at which an indicator changes color. Equivalence point: The point at which the moles of acid added equals the moles of base that you started with (should be the same as the end point.) An abrupt change in pH occurs at the equivalence point. 41

Titration curve:HCl Vs NaOH solution pH curve close to end point: 100ml of HCL titrated against NaOH of same normality NaOH Vol ml 1 M Sol pH 0.1M sol pH 98.02.03.0 99.02.33.3 99.52.63.6 99.83.04.0 99.93.34.3 100.0 7.07.07.07.0 100.1 10.79.79.7 100.211.010.0 100.511.410.4 101.011.710.7 102.012.011.0 42

Titration of Strong Acid with Strong Base Since the salt produced is neutral, the solution at the equivalence point has a pH of 7. the pH starts off low and increases as you add more of the base. The pH doesn't change very much until you get close to the equivalence point. Then it surges upwards very steeply 43

Titration Curve: Strong Acid Vs 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? 44

Titration of Strong Base with Strong Acid This curve is very similar to the previous one for the titration of a strong acid with strong base. The main difference is that the curve starts basic and then turns acidic after the equivalence point (rather than vice-versa). 45

Titration of Weak Acid with Strong Base The salt is basic, so equivalence point is at a pH > 7. Before the equivalence point, the solution acts as a buffer. The start of the graph shows a relatively rapid rise in pH but this slows down due to the buffering effect. 46

Titration Curve: Weak Acid Vs 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. 47

Titration of Weak Base with Strong Acid Salt formed is acidic, hence, equivalence point comes at a pH < 7. This curve is very similar to the titration of a weak acid with a strong base. The main differences are that the curve starts basic and has an acidic equivalence point. 48

Titration of a Polyprotic Acid A polyprotic acid titration will have more than one equivalence point. The first equivalence point represents the titration of the first proton, while the second equivalence point represents the titration of the second proton. 49

Redox Titration 50

Definitions 54

The main redox titration types S.No Redox titrationTitrant 1 IodometryIodine (I 2 ) 2 BromatometryBromine (Br 2 ) 3 CerimetryCerium(IV) salts 4 PermanganometryPotassium permanganate 5 DichrometryPotassium dichromate 56

Direction of End Points in Redox Reaction 64

Redox indicator 68

Complexometric Titration 70

Complexometric titration is a form of volumetric titration in which the formation of a colored complex is used to indicate the end point of a titration. The complexes are formed by the reaction of a metal ion ( an acceptor, a central atom or a cation) with an anion, a neutral molecule or very rarely a positive ion. 71

Complexometric titrations are particularly useful for the determination of a mixture of different metal ions in solution. An indicator capable of producing an distinct color change is usually used to detect the end point of the titration. 72

Complexometry : is the type of volumetric analysis involving the formation of complexes which are slightly ionized in solution, like weak electrolyte and sparingly soluble salt. Complex is formed by the reaction of metal ion (M n+ ) with either an anion e.g. [Ag(CN) 2 ] - or neutral molecule e.g. [Ag(NH 3 ) 2 ] + The metal ion is known as Central metal atom. The anion or neutral molecule is known as Ligand (L) 73

M + + LML Ag + Cu 2+ Ag + + 2 CN - + 4 CN - + 2 NH 3 [Ag(CN) 2 ] - [Cu(CN) 4 ] 2- [Ag(NH 3 ) 2 ] + Central metal atom = acts as Lewis acid (electron acceptor) Ligand = acts as Lewis base (electron donor) Coordinate bond (dative) = The bond formed between central metal atom (ion) (acceptor) and the Ligand (donor) 74

Dative bond is similar to covalent bond (formed of two electrons) But in dative bond the electrons pair are donated from one atom to the other.The atom gives electron pair is known as donor, while the atom accept electron pair is known as acceptor.The bond is represented by an arrow (  ) from donorto acceptor. NH 3  NH 3  Cu  NH 3  NH 3 75

* Coordination number = The no. of coordinate bonds formed to a metal ion by its ligands. * Characters of coordination number * 1- It is even number:2e.g. Ag +, 4e.g. Ni 2+, Cu 2+, 6 e.g. Fe 3+, Cr 3+ 2- It is usually double the charge of the metal. The charge of a complex is the algebraic sum of the charges of the central ion and ligand.. e.g. [Ag(CN) 2 ] -   Ag + +2 CN - 1 (+ve) + 2 (-ve) = 1 (-ve) e.g. [Fe(CN) 6 ] 3-   Fe 3+ +6 CN - 3 (+ve)+6 (-ve)=3 (-ve) The higher the valence of metal ion the more stable the complex e.g.Ferricyanide is more stable than Ferrocyanide 76

Types of complexing agents (Classification of ligands according to the no. of sites of attachment to the metal ion) Unidentate (Monodentate) Ligand or "Simple Ligand" The ligand attached to metal at one site e.g. H 2 O, NH 3, CN -, Cl -, I -, Br -, (i.e. forming one coordinate bond, or capable of donating one unshared pair of electrons) 77

H2CH2C H2CH2C NH 2 H2CH2C H2C H2C Cu CH 2 H2NH2N H2NH2N + Cu 2+ 2 Bidentate Ligand The ligand attached to metal at two sites. Ethylene diamine 78

Tridentate Ligand: The Ligand attached to metal at 3 sites Tetradentate Ligand: The Ligand attached to metal at 4 sites Diethylene triamine Triethylene tetramine 79

Chelation Chelate : It is a complex formed between the ligand containing two or more donor groups and metal to form ring structure. (heterocyclic rings or chelate rings). Chelating agents: organic molecules containing two or more donor groups which combine with metal to form complex having ring structure. Chelates are usually insoluble in water but soluble in organic solvent. Sequestering agent : Ligands which form water soluble chelates e.g. EDTA. 80

Classification of Complexometric Titrations 1)Direct Titration 2)Back titration 3)Replacement Titration 4)Alkalimetric titration of metals 81

DirectTitration In this type of titrations, the sample solution of metal ion, in the presence of a suitable buffer, is titrated against standard disodium edetate solution. M-EDTA complex must be more stable than M- Ind. complex in buffered medium. The compound to be determined is water soluble. The reaction between EDTA and metal must be rapid. If the reaction is slow it must be catalyzed. M n+ should not be ppt. at the pH of titration.If M n+ is ppt. as MOH, auxiliary reagent must be added to prevent pptn. of M n+. 82

Direct determination of water hardness Water hardness is due to the presence of Ca 2+ & Mg 2+ salts. EDTA forms complex with Ca 2+ & Mg 2+, Ca-EDTA complex is more stable than Mg-EDTA complex. At pH 12 EDTA forms complex with Ca 2+ only. Total Ca 2+ & Mg 2+ : Total Ca 2+ and Mg 2+ determined by titration with EDTA at pH 10 using ammonia buffer and EBT (Eriochrome Black T) as ind. Upon titration with EDTA, Ca 2+ will be chelated first, then Mg 2+. For Ca 2+ only: Direct titration with EDTA at pH 12 using 8% NaOH and Murexide. Mg 2+ is pptd. as Mg(OH) 2 leaving Ca 2+ which is titrated with EDTA For Mg 2+ : Total – Ca 2+ =Mg 2+ 83

Back Titration Addition of known excess of st. EDTA to the sample The medium is buffered. Exss. EDTA is titrated with standard soln. of another metal ion e.g. Mg 2+ or Zn 2+ It is used in the following cases: Insoluble substances e.g. BaSO 4, Ca(C 2 O 4 ) 2, PbSO 4, Mg 3 (PO 4 ) 2 … etc. Usually soluble in hot EDTA. The reaction between M n+ & EDTA is slow (incomplete) e.g. Fe 3+, Al 3+, Cr 3+, Th 4, … etc. The M n+ is pptd. at the pH suitable for titration e.g. Al(OH) 3. The colour change at the end point: From free ind. colour  to M-Ind. complex (opposite that direct titration) 84

Determination of Aluminium salts: Sample of Al 3+ is heated with known xss. of st. EDTA at pH 7-8. The soln. is then adjusted to pH=10 using ammonia buffer. The residual EDTA is titrated against st. Zn 2+ using EBT ( Eriochrome Black T) indicator. The colour change from blue to wine red. pH 7-8 +H 2 Y 2-  AlY - +2 H + Al 3+ Boil Zn 2+ pH 10 +H 2 Y 2-  ZnY 2- +2 H + Zn 2+ + H Ind. 2-  Zn-Ind. - + Bluewine red H+H+ 85

ReplacementTitration When both back titration and direct titration is not possible due to the end point not being sharp enough. Then the replacement titration is a method of choice. In this method of titration determination of metal ion is done by displacing magnesium or zinc ions from EDTA complex with an equivalent amount of metal ion and liberated Mg or Zn ions are then titrated with standard EDTA solution. Mordant black used as indicator. 86

Calcium salt is determined in this way. In this, add standard volume of Mg-EDTA solution to Ca salt in the presence of buffer. Calcium displaces Mg ion and forms a stable complex with EDTA as Ca-EDTA complex. The displaced and liberated Mg ions are then titrated Standard EDTA solution using Mordant black as indicator. Ca 2+ Mg 2+ +Mg-EDTA  Ca-EDTA+Mg 2+ +EDTA 2-  Mg-EDTA Cadmium, Lead and Mercury can also be determined by this titration. 87

Alkalimetric titration of metals Metal-EDTA complex formation reaction explains that proton are liberated from disodium edetate leading to formation of acid. M+M+ +H 4 Y  MY+4H + The acid that is formed can be titrated against a standard alkali but in in an unbuffered solution. End point detection can be done by using acid base visual indicator or potentiometric method of detecting end point. 88

 [M(CN) 2 ] - M+M+M+M+ +2 CN - +4 CN -  [M(CN) 4 ] 3- (B)- Triethanolamine : N CH 2 CH 2 OH CH 2 CH 2 OH CH 2 CH 2 OH Masking and demasking agent Masking agents: are reagents which prevent interfering ion from reaction without physical separation. These reagents form complexes with interfering ions which are more stable than complexes formed with ind. & EDTA. Examples of masking agent: (A)- KCN It is used as masking agent for Ag +, Cu 2+, Cd 2+, Co 2+, Ni 2+, Zn 2+, … etc. - It is used as masking agent for Fe 3+, Al 3+ and Sn 2+ (C) Fluoride (e.g. NH 4 F): - It is used as masking agent for Fe 3+ and Al 3+ to give hexafluoro complex [FeF 6 ] 3- and [AlF 6 ] 3- (D)- Iodide (KI): - It is used as masking agent for Hg 2+ to give tetraiodo complex (HgI 4 ) 89

Demasking agent :is the process in which masking substance reveres back to its ability to take part in the reaction. - are reagents which regain the ability of masked ion to enter the reaction with ind. and EDTA. Example: - The masking by CN – can be removed by: -mixture of formaldehyde – acetic acid -on addition of demasking agent to [Zn(CN) 4 ] 2-, Zn is liberated and titrated. [Zn(CN) 4 ] 2- +4 HCHO + 4 CH 3 COOH (less stable)  Zn 2+ CN + 4 CH 2 +4 CH 3 COO - OH Cyanohydrin (more stable) 90

Precipitation titration 91

Precipitation titration The titration in which precipitation reactions takes place is called precipitation titration. for eg: Titration of AgN0 3 with halide ions such as Cl or Br.  Solubility product The product of molar concentration of ions raised to the power equal to its stoichiometric coefficient presents in the ionic equation In saturated solution at a fixed tempreature is called solubility product. 92

It is donated by Ksp SOLUBILITY : The molar concentration of solute dissolved in solution at saturated condition at any temperature is called solubility. for eg: out of 5 moles of solute, suppose 2moles of solute is ionized in a one liter solution at any fixed temperature. Therefore, here solubility of solute is equal to 2moles/lit. 93

Principle of precipitation The product of concentrations of ions presents in a solution at any fixed temperature is called ionic product of the salt It is denoted by Q. If the solution is saturated solution then the ionic product is called solubility product of the salt “if the ionic product of the salt is greater than that its solubility product value then the preciptation of the salt takes place otherwise not,this is called the principle of precipitation“. 94

Depending upon the values of ionic product, the solutions can be classified into three different categories as follows: i.Q=Ksp,the solution is just saturated and no precipitation takes place. ii.If Q>Ksp,the solution is supersaturated and precipitation takes place. iii.If Q<Ksp,the solution is unsaturated and more of the solute can dissolve.so no precipitation takes place at this condition. 95

Factors affecting the solubility of precipitate a)Effect of tempreature: With increased temperature solubility of precipitate increases b)Effect of solvent : Solubility of inorganic salt is reduced by addition of organic solvent such as ethanol, methanol, propanol, and acetone but in presence of only water, hydration of ions of salt increases due to the high dipole moment of water molecule. 96

This hydration produces energy called hydration energy which is sufficient to overcome the attractive force between ion of solid lattice. The ions in crystals do not have so large an attraction for oraganic solvents,and hence the solubilities are usually less than in water. 97

c)Effect of acid: The solubilityof the salt of weak acid is affected by the addition of acid, hydrogen ion of added acid combines with the anions of the salt and forms weak acid thereby increasing the solubility of the salt. d)Formation of complex ions: The increase in solubility of a precipitate upon adding excess precipitating agent is frequently due to the formation of complex ion. 98

Argentometric titration (Argentometry): Titration involving precipitation with a standard solution of sliver nitrate is called argentometric titration or argentometry. conditions required for argentometry :  Precipitates should be practically insoulble.  Precipitation reactions should be rapid and quantitative. 99

Precipitate should not interfere in the sharp detection of the end point. Titration results should not differ apprecially due to adorptions on the precipitate. Different methods involved in argentometry: a)Mohr’s method b)volhard’s method c)Fajan’s method 100

Mohr’s method A precipitaion titration in which sliver ion is used as titrant and chromate ion as indicator is called mohr method. This method is appplicable for the quantitative analysis of halide ions (Cl, Br and I). THEORY: Solubility product of ppt of sliver halide is greater than that of sliver chromate so that solubility of sliver halide becomes less than that of sliver chromate. There fore, ppt of sliver halide is formed first than that of silver chromate. 2 Ag + (aq) + CrO 2– (aq) → Ag CrO (s) 424 101

For example,in titration of AgNO 3 with Nacl in presence of indicaor The chromate ion, CrO 4 2– Here initially AgNO 3 reacts with Nacl and produce white ppt of AgCl. NaCl + AgNO 3 AgCl + NaNO 3 white ppt 102

Therefore concentration of Ag 2 CrO 4 should be 0.014 M at end point of titrations but this gives the intense yellow colour which interfere the colour of end point there fore generally 0.003M- 0.005M concentration of Ag 2 CrO 4 is maintained at end point. 104

Precaution : The titration must be carried out in pH range of 6.5-9 (i.e neutral or slightly in alkaline) In more alkaline medium sliver hydroxide may be precipitated out in acidic solution conccentration of chromate ion decreases greatly as: 105

In acidic medium dueto the decrease in concentration of brick red color can not seen at the end point of titration. In actual practice the concentration of chromate produces an intense yellow color to such and extent that the end point is masked. Therefore, normally 0.003-0.005M concentration should be employed in analytical procedures. 106

Volhard’smethod A precipitation titration in which Ag+ ion is precipitated by SCN - (thiocyanate ions) in presence of Fe(lll) ions indicator in acidic medium is called volhard method. i.e initially Ag + (aq) + SCN – (aq) → AgSCN(s) 107

At end point Fe 3+ (aq) + SCN – (aq) → [FeSCN] 2+ reddish brown Here initially thiocyanate react with silver ions and forms precipitate at end point excess of thiocyanate (SCN-) react with Fe(lll) and forms reddish brown complex which indicate the end point of reaction. 108

This volhard method is used to determine the concentration of Ag+ ions or concentration of halide ions (i.e. Cl -, Br - and I) indirectly i.e. by back titration. precautions:  During the titration in resulting solution concentration of nitric acid should be 0.5M- 1M. Here, strong nitric acid is not used because it prevents the formation of thiocyanate iron(lll) complex. Here acidic medium prevents the hydrolysis of Fe(lll) as Fe(0H) 3. 109

 The solution must be free from nitrous acid otherwise it gives the instantred colouration forming the thiocynic acid.  Temperature of the solution must be maintained below 250 o C sinceat an elevated temperature the red colour of the ferric thiocyanate complex fades away rapidly. 110

When the excess of silver has reacted, the thiocyanate may react with the silver chloride, since thiocyanate is less soluble than silver chloride. i.e AgCl+SCN – → AgSCN + Cl – This will take place before reaction occurs with the iron (lll) ions in the solution, so there will be considerable titration error. 111

Fajan’s method (indicator adsorption method) The precipitation titration in which silver ions is titratedwith halide or thiocyanate ions in presence of adsorption indicator is called fajan’s method. Since the adsorption of indicator takes place at end point the method is also called indicator adsorption method. The indicator, which is a dye, exists in solution as the ionized form, usually an anion. 112

The method is generally used for the quantitative analysis of halide ions or thiocyanate ions. For eg: Titration of Cl – ions with AgN0 3 in presence of adsorption indicator here AgN0 3 is kept in burette and the Cl – ion solution with indicator is taken in titration flask. 113

a)Before the equivalence point: Before the equivalent point, colloidal particles of AgCl are negatively charged due to the adsorption of Cl – from the solution. The adsorbed Cl – from the primary layer which attract the positively charged Na + ions from the solution to form a more loosely held secondary layer as shown in figure. 114

b)After the equivalent point, excess of silver ions Ag+ displace the Cl – ions and form the primary layer which attract the negatively charged nitrate ions N0 3 –. c)At the end point, anion of indicator in (weak organic acid or base) replace the negatively charged ion N0 3 – from the second layer and give the intense color. -This intense color gives the end point of titration. 116

Precautions:  The AgCl should not be allowed to coagulate into large particles at the equivalent point since this will greatly decrease the surface available for adsoption of the indicator.  Aprotective colloid such as dextrin should be added to keep the precipitate highly dispersed in thepresence of dextrin the color change is reversible and if the end point is overrun, one can back titrate with a standard chloride solution. 118

 The adsorption of indicator should start just before the equivalent point and increase rapidly at the equivalent point some unsuitable indicators are so strongly adsorbed that they actually displace the primary adsorbed ion well before the equivalent point is reached.  The pH of the titration medium must be controlled to ensure a sufficient concentration of the ion of the weak acid or baseindicator. 119

fluorescein for example has a (Ka=10-7) and in solutions more than acidic than pH 7 the concentration of Fluoroscence ions is so small that no color change is observed fluoresceion can be used only in the P H range of about 7 to 10. It is preferable that the indicator ion be opposite in charge to the ion added as the titrant adsorption of the indicator will then not occur untill excess titrant is present. 120

Adsorption indicator 121

Comparison of silver titration methods 122

Gravimetric Analysis 123

Jons Jacob Berzelius (1779 - 1848), considered the leading chemist of his time, developed much of the apparatus and many of the techniques of 19th century analytical chemistry.Examples include the use of ashless filter paper in gravimetry, the use of hydrofluoric acid to decompose silicates, and the use of the metric system in weight determinations. He performed thousands of analyses of pure compounds to determine the atomic weights of most of the elements known then. Berzelius also developed our present system of symbols for elements and compounds. A History of Gravimetric Analysis 124

A History of Gravimetric Analysis Theodore W. Richards (1868 -1928) …and his graduate students at Harvard developed or refined many of the techniques of gravimetric analysis of silver and chlorine. These techniques were used to determine the atomic weights of 25 of the elements by preparing pure samples of the chlorides of the elements, decomposing known weights of the compounds, and determining the chloride content by gravimetric methods. From this work Richards became the first American to receive the Nobel Prize in Chemistry in 1914. 125

Gravimetric Analysis: is based upon the measurement of mass Gravimetric Analysis generalized into two types: precipitation and volatilization (i ) A technique in which the amount of an analyte in a sample is determined by converting the analyte to some product  Mass of product can be easily measured (ii) Analyte: the compound or species to be analyzed in a sample Advantages -  requires minimal equipment Disadvantage –  requires skilled operator,  slow. 126

The quantitative determination of a substance by the precipitation method of gravimetric analysis involves isolation of an ion in solution by: 1.precipitation reaction, 2.filtering, 3.washing the precipitate free of contaminants, conversion of the precipitate to a product of known composition, 4.drying 5.weighing the precipitate and determining its mass by difference. 127

S.NoMechanism of Action Example 1Hydration ANHYDRONE® (Magnesium Perchlorate anhydrous), CaCl 2, MgO, MgSO 4, K 2 CO 3, KOH, Drierite, Na 2 SO 4 (anhydrous), H 2 SO 4, ZnCl 2 2Absorption and Adsorption BaO, CaSO 4, Molecular Sieve, H 3 PO 4, NaOH Pellets 3ChemisorptionCaO, P 2 O 5 Common Desiccants 129

Silica gel goes from blue to pink as it absorbs moisture Can be regenerated in oven. Anhydrous sodium sulfate gets clumpy as it absorbs water. 130

Precipitation Reaction of potassium iodide solution and lead (II) nitrate solution. Determination of lead (Pb +2 ) in water 132

Determination of lead (Pb +2 ) in water Pb + + 2Cl -  PbCl 2 (s)  By adding excess Cl - to the sample, essentially all of the Pb +2 will precipitate as PbCl 2.  Mass of PbCl 2 is then determined. used to calculate the amount of Pb +2 in original solution Reagent Analyte Solid Product 133

Mechanism of precipitation 1.Induction period time between mixing and visual appearance of a precipitate called the induction period 2.Nucleation is the formation, in a super saturation solution, of the smallest aggregate of molecules capable of growing into a large precipitate particle. 134

3.Crystal growth Once a nucleation aggregate has formed, it begins to grow as ions or molecules from the solution deposit on the surface in a regular, geometric pattern. 4.Aggregate growth Natural cohesive forces exist between particles having the same composition and, as a result, most precipitate to consist of a relatively few large aggregate of crystals. Crystal Growth 135

Why colloids occur?? 136

How colloidal suspension is prevented???? Keep the volume of the counter ion layer small: Keep the charge of the primary adsorption layer small by avoiding an excess of the precipitating ion. Increase the ionic strength by adding soluble electrolyte. Heating and stirring reduces the volume of the electrical double layer. This will allow the ions in the counter ion layer be closer to the opposite charged ions of the primary layer. 137

Desired Properties of Solid Product  Should be very insoluble  Easily filterable (i.e., large crystals)  Very Pure  Known and constant composition Few precipitates have all of these properties, but in most cases appropriate techniques can help optimize these qualities 138

Solubility: The solubility of a precipitate can be decreased by:  Decreasing temperature of solution  Using a different solvent - usually a less polar or organic solvent (likes dissolves likes) Solubility vs. pH Solubility vs. Temperature Solubility vs. Common Ion Effect 139

Solubility depends pH Influence the solubility of the analytical precipitate. Possibility of interference from other substance. Example 01 : During Calcium oxalate precipitation REASONREASON 140

Solution of precipitating reagent added to Sample solution Precipitate formed Eg: AgNO 3 + NaCl →AgCl + NaNO 3 Precipitation process involves heterogeneous equilibrium as of steps Solubility product Precipitation Process Actually precipitate occurs in series of steps 141

SUPERSATURATION Increase super saturation -Increase nucleation Should avoid Super saturation & nucleation Super saturation is one of the important parameter in precipitation process It ‘s play vital role in gravimetry. 142

Nucleation- Individual ions/atoms/molecules coalescence to form “nuclei (small particles come together to form colloid particles) 143

Nucleation Effects of Increase nucleation Initial nucleus will grow by the deposition of their precipitate particles To form a crystal of certain geometric shape. Again the greater super saturation More rapid crystal growth rate & colloidal precipitate formed Increase the growth rate increased chance of imperfection in the crystal & surface area of precipitate increase this leads to easy trapping of impurities. Trapping of impurities 144

Relationship between supersaturation & Nucleation High relative supersaturation Low relative supersaturation 145

Filterability: product be large enough to collect on filter: Doesn’t clog filter Doesn’t pass through filter Best Case: Pure Crystals Worst Case: Colloidal suspension Difficult to filter due to small size Tend to stay in solution indefinitely  suspended by Brownian motion usually 1-100 nm in size 148

Conditions for analytical precipitation An analytical precipitate for gravimetric analysis should consist of perfect crystals large enough to be easily washed and filtered. The perfect crystal would be large and free from impurities. The precipitate should also be "insoluble". Colloidal suspension Crystal formation Want to Convert to 149

It has been shown that the particle size of precipitates is inversely proportional to the relative supersaturation of the solution during precipitation; Relative supersaturation = S = concentration of solute in solution at equilibrium Q = actual concentration of solute added to solution 150

Methods used to improve particle size and filterabil ity 1. Precipitation from hot solution The solubility S of precipitates increases with temperature and so an increase in S decreases the supersaturation. 2.Precipitation from dilute solution This keeps Q low. Slow addition of precipitating reagent with effective stirring. This also keeps Q low; stirring prevents local high concentrations of the precipitating agent. 3.Precipitation at a pH near the acidic end of the pH range Many precipitates are more soluble at the lower (more acidic) pH values and so the rate of precipitation is slower. 4.Digestion of the precipitate. Heating the precipitate in the precipitating solution, a process called digestion, results in larger and purer particles by giving the crystal a chance to dissolve and reprecipitate. 151

Impurities in Precipitates Impurities can be incorporated into a precipitate during its formation, called co-precipitation, or after its formation while still in contact with the precipitating solution, called postpricipitation Co-precipitation a) Surface adsorption 152

b) Occlusion Impurities absorbed or trapped within pockets in the crystal c) Inclusion Impurities placed in the crystal instead of analyte 153

Occlusion & Inclusion OCCLUSION Trapping of impurities with in the crystal. Example : Water may trapped in pockets when silver nitrate crystals are formed INCLUSION Trapping of impurities with in the crystal lattice. Example : K + in NH MgPO 4 Impurities Occlusion & Inclusion impurities are difficult to remove. Digestion may help some but this is not completely effective. Purification by dissolving & reprecipitating may be helpful 154

The degree of contamination of a ppt. by Co- precipitation depends on specific adsorption, specific surface, decreasing solubility of the contaminant, increasing concentration of the contaminant, higher valence, and smaller size of the foreign ion, lower tempareture and greater no. of particles/ unit mass of the precipitation. 155

Co-precipitation can be minimised by the following techniques By precipitation from hot solution so that initially a few particles are formed and relative super saturation is increased. For precipitations having appreciable solubility, however the solution must be cooled and the relative super saturation is increased by the control of pH or any other device. By precipitation from dilute solution to decrease the concentration of foreign ions. By adding reagent as a solution, slowly with stirring in order to reduce the local concentration (Q). By prior removal of the foreign ion as precipitation by rn with a different. Reagent or by conversion of the foreign ion to a different valence state or sometimes to a complex. 156

Fe +3, Ca +2 HClHCl Fe(OH) 3.xH 2 O Included as co-ppt. Fe +3, Ca +2 (less) OH - Istppt. 2 nd ppt. +Soln (Ca +2 ) +2+2 Fe(OH) 3.xH 2 O (Ca)+Ca +2 in soln i)By allowing the precipitation to satnd (digestion and aging) for perfect growth of crystals ii)By thorough and proper washing of the precipitation. iii)By double precipitation. 157

Three examples of impurities that may form during precipitation. The cubic frame represents the precipitate and the blue marks are impurities: (a) inclusions, (b) occlusions, and (c) surface adsorbates. Inclusions are randomly distributed throughout the precipitate. Occlusions are localized within the interior of the precipitate and surface adsorbates are localized on the precipitate’s exterior. For ease of viewing, in (c) adsorption is shown on only one surface. 158

Post precipitation Post –precipitation involves incorporation of impurities when the supernatant liquid is super saturated w.r.t some phase which crystalline slowly. Eg. thus in case of metal sulfide precipitation from 0.1N acid medium (HCl), ZnS is not normally precipitated. A solution saturated with H 2 S, however, remains supersaturated with ZnS, which slowly separates as a precipitation which is accelerted in presence of the fine particles of HgS or CuS. This is known as post precipitation of ZnS. Similar situation arises when Mg remains in a solution containing large excess of oxalate after Ca-Oxalate precipitation. Mg 2 C 2 O 4 precipitate slowly from the supernatant liquid by post-precipitation. 159

Post precipitation difference from Co- precipitation in the following respects Contamination increases with aging or digestion (in case of post precipitation), but the reverse is true in case of co-precipitation. Contamination increases with faster agitation or stirring of the supernatant liquid in case of post precipitation, but the reverse is true for co-precipitation. Degree of contamination or its magnitude is usually much greater due to post precipitation in comparison to co-precipitation. To avoid post –precipitation a shorter period of digestion particularly under warm condition is to be carried out and the ppt. should be filtered as quickly as possible so that the precipitation does not remain in contact with a mother–liquor for a long time. 160

There are several requirements that must be met to make precipitation reliable: The precipitate must have a very low solubility in water; i.e. its K sp must be very small number It must precipitate in a high state of purity or be capable of reprecipitation for further purification. It must be capable of drying or of ignition. It should not be hydroscopic at room temperature. 161

Ageing &digestion The precipitate should be in contact with the solution from which the precipitate is formed. Warm the solution that contains the precipitate for some time to obtain complete precipitation in a form which can be readily filtered. 162

During the process of ageing and digestion, two changes occur: After precipitation has occurred, the very small particles, which have a greater solubility than the greater ones, tend to pass into solution and will redeposit upon the larger particles. Thus co precipitation on the minute particles is eliminated. The rapidly formed crystals are irregular. Thus on ageing they will become regular and the surface area is reduced, so adsorption will be reduced. The net result of digestion is usually to reduce the extent of co precipitation and to increase the size of the particles, rendering filtration easier. 163

Filtration A precipitate may be separated by filtering it through paper, sintered glass, or sintered porcelain. The choice depends on the nature of the precipitate and on the temperature to which it will be heated after filtering. 164

Washing The precipitate and filter must be washed with suitable electrolyte to remove dissolved solids that remain in the precipitate and wetted filter. Problems with coprecipitation and surface adsorption may be reduced by careful washing of the precipitate. With many precipitates, peptization occurs during washing. 165

Precipitates from ionic compounds - need electrolyte in wash solution TO keep precipitate from breaking up and redissolving (peptization) Electrolyte should be volatile removed by drying - HNO 3, HCl, NH 4, NO 3, etc. Example: AgCl(s) should not be washed with H 2 O, instead wash with dilute HNO 3 166

Drying/Igniting Precipitates  Precipitates are dried at about 120 o C for accurate, stable mass measurements 167

 True or correct or actual value are known only when the count object or when a quantity is assigned a particular value such as atomic weight. Otherwise the true value of a quantity is never known.  Standard value is observed value given by the expert using a suitable method and good quality apparatus and chemicals.  Observed value is the result obtained.  Error: The difference between true value or standard value is called error 168

 The Accuracy of an analytical procedure expresses the closeness of agreement between the value, which is accepted either as a conventional true value or an accepted reference value and the value observed (individual observation or mean of measurements).  The Precision (VARIABILITY) of an analytical procedure is nearness between several measurements of the same quantity usually expressed as the standard deviation (S), variance (S 2 ), or coefficient of variation (= relative standard deviation, R.S.D.) of a series of measurements. 169

Difference between Precision & Accuracy Two analysis (I & II) of substance whose true value is 100% is given below; Analysis I 98.80, 98.82, 98.84, 98.82 Average valueis 98.82% Error is (100.00-98.82) = 1.18% The accuracy is poor but precision is better Analysis II 100.00, 99.60, 99.70, 99.10 Average valueis 99.60% Error is (100.00-99.60) = 0.4% The precision is poor but accuracy is good 171

Coagulation, agglomeration 172

 The term error is used to show the difference between measured and true value.  Since the true value are never known one has to make use standard value.  The standard value can be obtained by  Absolute Method: sample is synthesized using know quantities to obtain a primary standard.  Comparative method: standard data is obtained. 173

 DETERMINATE OR SYSTEATIC ERRORS  Personal error  Instrumental or reagent error  Method errors  Additive error  Proportional error  INDTERMINATE OR RANDOM ERRORS  The cause of a random error may not be known 174

 Propositional error: the magnitude of error depends upon sample size  Additive error: The value of error is constant is independent of amount of sample taken for analysis  Endpoint error 175

 Proper calibration of apparatus  Running a blank determination  Carrying out a control determination (use of standard and reference)  Use of independent method. (results of two different methods are compared)  Reparative determinationand statistical evaluation 176

 Average is measure of central tendency  It is the arithmetic mean of different values obtained by measuring the same quantity several time.  Arithmetic mean =sum of different value Number of times determination is made 177

 Deviation (d): difference between the measured value and average valued =(x 1 -x - )  Average or mean deviation (d-) is arithmetic mean of different deviation observed d - = d 1 +d 2 +d n n  Relative Mean deviation is= Mean deviationx 100 Mean The positve or negative sing of individual deviation is ignored 178

 The standard deviation is square root of the sum of the squared individual deviation divided by (n-1) 23n23n s = √ d 1 2 + d 2 + d 2 + d 2 (n-1)  The square of standard deviation is called Variance and coefficient of variation (C.V) (also known as relative Std dev.) is defined as C.V. = s x 100 X - 179

 The number of significant figures in a result is simply the number of figures that are known with some degree of reliability. The number 13.2 is said to have 3 significant figures. The number 13.20 is said to have 4 significant figures 180

 All measurements are approximations--no measuring device can give perfect measurements without experimental uncertainty. By convention, a mass measured to 13.2 g is said to have an absolute uncertainty of 0.1 g and is said to have been measured to the nearest 0.1 g.  In other words, we are somewhat uncertain about that last digit —it could be a "2"; then again, it could be a "1" or a "3". A mass of 13.20 g indicates an absolute uncertainty of 0.01 g 181

Rules for deciding the number of significant figures in a measured quantity 1.All nonzero digits are significant: 1.234 g has 4 significant figures, 1.2 g has 2 significant figures. 2. Zeroes between nonzero digits are significant: 1002 kg has 4 significant figures, 3.07 mL has 3 significant figures. 3. Leading zeros to the left of the first nonzero digits are not significant; such zeroes merely indicate the position of the decimal point: 0.001 oC has only 1 significant figure, 0.012 g has 2 significant figures. 4. Trailing zeroes that are also to the right of a decimal point in a number are significant: 0.0230 mL has 3 significant figures, 0.20 g has 2 significant figures. 182

5. When a number ends in zeroes that are not to the right of a decimal point, the zeroes are not necessarily significant: 190 miles may be 2 or 3 significant figures, 50,600 calories may be 3, 4, or 5 significant figures. The potential ambiguity in the last rule can be avoided by the use of standard exponential, or "scientific," notation. For example, depending on whether the number of significant figures is 3, 4, or 5, we would write 50,600 calories as: 5.06 × 104 calories (3 significant figures) 5.060 × 104 calories (4 significant figures), or 5.0600 × 104 calories (5 significant figures). 183

In carrying out calculations, the general rule is that the accuracy of a calculated result is limited by the least accurate measurement involved in the calculation. 1. In addition and subtraction, the result is rounded off to the last common digit occurring furthest to the right in all components. Another way to state this rules, it that, in addition and subtraction, the result is rounded off so that it has the same number of decimal places as the measurement having the fewest decimal places. For example, 100 (assume 3 significant figures) + 23.643 (5 significant figures) = 123.643, which should be rounded to 124 (3 significant figures). 2. In multiplication and division, the result should be rounded off so as to have the same number of significant figures as in the component with the least number of significant figures. For example, 3.0 (2 significant figures ) × 12.60 (4 significant figures) = 37.8000 which should be rounded off to 38 (2 significant figures) 184

1. If the digit to be dropped is greater than 5, the last retained digit is increased by one. For example, 12.6 is rounded to 13. 2. If the digit to be dropped is less than 5, the last remaining digit is left as it is. For example, 12.4 is rounded to 12. 3. If the digit to be dropped is 5, and if any digit following it is not zero, the last remaining digit is increased by one. For example, 12.51 is rounded to 13. 4. If the digit to be dropped is 5 and is followed only by zeroes, the last remaining digit is increased by one if it is odd, but left as it is if even. For example, 11.5 is rounded to 12, 12.5 is rounded to 12. This rule means that if the digit to be dropped is 5 followed only by zeroes, the result is always rounded to the even digit. The rationale is to avoid bias in rounding: half of the time we round up, half the time we round down. 185

1. 37.76 + 3.907 + 226.4 = 2. 319.15 - 32.614 = 3. 104.630 + 27.08362 + 0.61 = 4. 125 - 0.23 + 4.109 = 5. 2.02 × 2.5 = 6. 600.0 / 5.2302 = 7. 0.0032 × 273 = 8. (5.5) 3 = 9. 0.556 × (40 - 32.5) = 10. 45 × 3.00 = 11. 3.00 x 105 - 1.5 x 102 = (Give the exact numerical result, then express it the correct number of significant figures). 12. What is the average of 0.1707, 0.1713, 0.1720, 0.1704, and 0.1715? 13. Calculate the sum of the squares of the deviations from the mean for the five numbers gives in Question 12 above, in two different ways: 186

1. 37.76 + 3.907 + 226.4 = 268.1 2. 319.15 - 32.614 = 286.54 3. 104.630 + 27.08362 + 0.61 = 132.32 4. 125 - 0.23 + 4.109 = 129 5. 2.02 × 2.5 = 5.0 6. 600.0 / 5.2302 = 114.7 7. 0.0032 × 273 = 0.87 8. (5.5)3 = 1.7 x 102 9. 0.556 × (40 - 32.5) = 4 10. 45 × 3.00 = 1.4 x 102 11. 3.00 x 105 - 1.5 x 102 = (Give the exact numerical result, then express it the correct number of significant figures). 12. What is the average of 0.1707, 0.1713, 0.1720, 0.1704, and 0.1715? Answer = 0.1712 13. Calculate the sum of the squares of the deviations from the mean for the five numbers given in Question 12 above, in two different ways: (a) carrying all digits through all the calculations; (b) Round all intermediate result to 2 figures 187

Quantitative Analysis MD KHAJA ARIF Asst. Professor of Chemistry MVS Government Arts & Science College Mahabubnagar Mobile:

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