Introduction to Volumetric Analysis for Basic Chemistry

Introduction to Volumetric Analysis for Basic Chemistry

1. VOLUMETRIC ANALYSIS

2. LESSON OUTCOMES Types of volumetric analysis Principles of titration: equivalence point and end point Back-titration Acid base titration curve

3. TYPES OF VOLUMETRIC ANALYSIS

4. VOLUMETRIC ANALYSIS 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 .

5. TYPES OF TITRATION Types Example Acid-Base Determine % acetic acid in vinegar Complexometric Water hardness (Determine Ca 2+ in water) Precipitation Determine of Cl - in water Redox Quantify hydrogen peroxide (H 2 O 2 )

6. Titration- Preparation Two solutions are used: The solution of unknown concentration. The solution of known concentration which is also known as the standard solution . Write a balanced equation for the reaction between your two chemicals. Clean all glassware to be used with distilled water. The pipettes and burettes will be rinsed with the solutions you are adding to them.

7. Titration Set up The burette is attached to a clamp stand above a conical flask. The burette is filled with one of the solutions (in this case a yellow standard solution). A pipette is used to measure an aliquot of the other solution (in this case a purple solution of unknown concentration) into the conical flask. Prepare a number of flasks for repeat tests. Last, an indicator is added to the conical flask.

8. The Titration Process Read the initial level of liquid in the burette Turn the tap to start pouring out liquid of the burette into the flask. Swirl the flask continuously. When the indicator begins to change colour slow the flow . When the colour changes permanently, stop the flow and read the final volume. The volume change needs to be calculated (and written down). This volume is called a titre. Repeat the titration with a new flask now that you know the ‘ rough ’ volume required. Repeat until you get precise results.

9. Acid-Base Titration When the reaction involves an acid and a base, the method is referred to as an acid-base titration. The aim of acid-base titration is to determine the volume of two solutions that will exactly react with one another A known volume of a solution is placed in a conical flask. The other solution is added from a burette until there is enough of it to completely react with the solution originally in the flask However, most acids, bases and their salts are colourless . Therefore an indicator is needed to determine the end-point and equivalence point of the titration. Acid-base indicator would change colour at the end-point of a titration

10. Procedure of Acid- Base Titration

11. Definition of terms Volumetric Analysis Measurement of volume of a solution of known concentration , which is used to determine the concentration of the analyte 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

12. What is a titration A procedure of carefully controlled addition of reagent (titrant) to an analyte until the reaction between the two is judged complete Known concentration which is called standard Usually in burette Usually in a conical flask

13. Successful Volumetric Titration must employed….. 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.

14. Direct and Back Titration 1. Direct Titration A titration process where a titrant (standard solution) is added to the analyte until the reaction is between analyte and titrant is judged to be complete. 2. Back Titration A titration process in which the excess standard solution used to react with analyte is determined by titration with a second standard solution .

15. 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 . Equivalence point The point at which an equivalent or stoichiometric amount of titrant is added to the analyte based on the stoichiometric equation What is End Point and Equivalence Point in Titration?

16. End Point VS Equivalence Point END POINT EQUIVALENCE POINT The point at which the reaction is observed to be completed The point at which an equivalent or stoichiometric amount of titrant is added to the analyte The end point signal frequently occurs at some point other than the equivalence point which tells the analyst to stop adding TITRANT and record the volume. The point at which the reaction is complete The selected indicator should change color very near to the equivalent point. Theoretically at the equivalence point we can calculate the amount of titrant that is required to react EXACTLY with the amount of analyte present.

17. Good End Point Bad End Point (Overshot)

18. 1. Direct Titration (Disadvantages) Sometime direct titration are not feasible due to….. Reaction kinetic or the reaction rate is slow . No suitable indicator in the direct titration. The color change is slow or delay. The end point is far from the equivalent point.

19. 2. Back Titration In a simple acid-base titration, a base (reagent) is added in a known quantity – greater than the amount required for acid neutralization. Acid and base reacts completely. The remaining base is titrated with a standard acid. The system has gone from being ACID , past the equivalent point to the BASIC (excess base), and back to the equivalence point again. The final titration to the equivalence point is called BACK TITRATION.

20. 2. Back Titration Example, the titration of insoluble organic acid with NaOH is not practical because the reaction is slow . To overcome it add NaOH in excess and allow the reaction to reach completion and then titrate the excess NaOH with a standard solution of HCl . The system has gone from being ACID , past the equivalence point to the BASIC side (excess base), and then back to the equivalence point. The final titration to the equivalence point is called a BACK TITRATION.

21. Why we need to do back titration? The analyte may be in solid form The analyte reacts slowly with the titrant in direct/forward titration The analyte may contain impurities which may interfere with direct titration

22. 150.0 mL of 0.2105 M nitric acid was added in excess to 1.3415 g calcium carbonate. The excess acid was back titrated with 0.1055 M sodium hydroxide. It required 75.5 mL of the base to reach the end point. Calculate the percentage (w/w) of calcium carbonate in the sample. Examples On Back Titration Method

23. First write a balance equation for the above reactions. 2HNO 3 + CaCO 3  Ca(NO 3 ) 2 + CO 2 + H 2 O ------ 1 HNO 3 + NaOH  NaNO 3 + H 2 O ------- 2 From Equations above: 2 mole HNO 3 required 1 mole CaCO 3 1 mole HNO 3 required 1 mole NaOH Initial mole of acid = Molarity x volume (L) = 0.2105 mol/L x 0.150 L = 0.031575 mol

24. Remaining/excess acid during back titration. Mole of excess acid = molarity x volume (L) = 0.1055 mol /L x 0.0755 L = 0.007965 mol acid. Mole of CaCO 3 in sample= ½ x mmole acid = ½ x 0.02361 = 0.011805 mol Mole of acid reacted with CaCO 3 = Initial – remaining/excess = ( 0.031575 – 0.007.965 ) mol = 0.02361 mol

25. Mass of CaCO 3 in sample = mole x molar mass = 0.011805 mol x 100 g/mol = 1.1805 g. % (w/w) of CaCO 3 in sample = 87.99 % (w/w)

26. Indicator and Choice of Indicator Acid –Base Indicators : The acid-base indicator function by changing colour just after the equivalence point of a titration; this colour change is called the end point . The end point is most often detected visually . Acid-base indicator are usually weak organic acids or bases that dissociate partially in water with the undissociated molecules have different color from their ions. ( eg : Undissociated molecule of phenolphthalein is colourless & it anion is pink) Indicators can be monoprotic ( HIn ) or diprotic (H 2 In) acids.

27. The acid form of an indicator is usually coloured ; when it loses a proton resulting in anion (In-) , or base form of the indicator, exhibiting different colour .

28. Acid Base Indicators Common Name Transition range Colour Change ACID BASE Crystal violet 0.1 – 1.5 Yellow Blue Thymol blue 1.2 – 2.8 Red Yellow Methyl yellow 2.9 – 4.0 Red Yellow Methyl orange 3.1 – 4.4 Red Orange Bromocresol green 3.8 – 5.4 Yellow Blue Methyl red 4.2 – 6.3 Red Yellow Chlorophenyl red 4.5 – 6.4 Yellow Red Bromothymol blue 6.2 – 7.6 Yellow Blue Phenol red 6.8 – 8.4 Yellow Red Thymol blue 8.0 – 9.6 Yellow Blue Phenolpthalein 8.3 – 10.0 Colourless Red Alizarin yellow 10.0 – 12.0 Colourless Yellow

29. Choosing a Titrant In theory, any strong acid or strong base can be used as titrant . The reason for this is that most reaction involving a strong acid or a strong base is quantitative . Strong Acid Titrant Weak Acid Titrant Hydrochloric acid ( HCl ) Nitric acid (HNO 3 ) Perchloric acid (HClO 4 ) Phosphoric acid (H 3 PO 4 ) Acetic acid (CH 3 COOH) Ammonium ion (NH 4- ) Hydrogen fluoride (HF) Carbonic acid (H 2 CO 3 ) Nitrous acid (HNO 2 ) Hydrogen sulphide (H 2 S) Hydrogen cyanide (HCN Acid Titrant

30. Strong Base Titrant Weak Base Titrant Sodium hydroxide, NaOH Potassium hydroxide, KOH Magnesium hydroxide, Mg(OH) 2 Barium hydroxide , Ba(OH) 2 Ammonium hydroxide, NH 4 OH Amine acetate. Carbonate, CO 3 2- Fluoride ion,F - Sodium carbonate,NaCO 3 Base Titrant

31. Acid- Base Titration Curve

32. The Titration (or pH) Curve The change in pH during an acid-base titration can be followed by measuring the pH of the mixture using a pH meter The change is then plotted against the volume of base (or acid) added from the burette These titration curves allow us to choose the most suitable indicator for the particular titration

33. The Relationship Between pH & pOH pH = -log [H + ] [H + ] = [OH - ] [H + ] > [OH - ] [H + ] < [OH - ] Solution Is neutral acidic basic [H + ] = 1 x 10 -7 [H + ] > 1 x 10 -7 [H + ] < 1 x 10 -7 pH = 7 pH < 7 pH > 7 At 25 C pH [H + ] [H + ][OH - ]= K w =1.0 x 10 -14 -log [H + ] – log [OH - ] = 14.00 pH + pOH = 14.00

34. Titration curve for the titration of 25.00 mL of 0.20 M HCl with the 0.20 M NaOH Sudden sharp changes of pH Salt formed: NaCl Neutral salt: So pH is 7 Slow change of pH Strong acid with strong base Buffer Region NaOH ( aq ) + HCl ( aq ) H 2 O ( l ) + NaCl ( aq )

35. Before addition of NaOH , the pH=1. As NaOH is added, the pH increases gradually There is a sharp increase in pH (from 4 to 11) slightly before and after equivalence point (pH=7) Beyond equivalence point, pH increases gradually as more NaOH added The suitable indicators are: Indicator pH Range Colour Change Methyl red 4.2 – 6.3 Red- Yellow Chlorophenyl red 4.8 – 6.4 Yellow- Red Bromothymol blue 6.0 – 7.6 Yellow- Blue Phenol red 6.4 – 8.4 Yellow -Red Cresol red 7.2-8.8 Yellow -Red Phenolpthalein 8.3- 10.00 Colourless - pink

36. HCl ( aq ) + NH 3 ( aq ) NH 4 Cl ( aq ) NH 4 + ( aq ) + H 2 O ( l ) NH 3 ( aq ) + H + ( aq ) At equivalence point (pH < 7): Weak base with Strong acid Sharp decrease in pH pH is less than 7 (usually around 5-6) Salt formed is acidic

37. Before addition of HCl , the pH=11. When HCl is added, the pH increases gradually as more NH 3 neutralised and the solution becomes more acidic. There is sharp decrease in pH (from 7 to 3) slightly before and after equivalence point. Beyond equivalence point, pH decreases gradually as more HCl is added The suitable indicators are: Indicator pH Range Colour Change Methyl Orange 3.1-4.4 Red-Yellow Methyl Red 4.2-6.3 Red-Yellow Chlorophenol blue 4.8-6.4 Yellow-Red

38. 41 CH 3 COOH ( aq ) + NaOH ( aq ) CH 3 COONa ( aq ) + H 2 O ( l ) CH 3 COO - ( aq ) + H 2 O ( l ) OH - ( aq ) + CH 3 COOH ( aq ) At equivalence point (pH > 7): Weak acid with Strong base pH changes rapidly pH is higher than 7 (usually around 8-9) Salt formed is alkaline

39. Before addition of NaOH , pH≈ 2.9 ( ethanoic acid is a weak acid). When NaOH is added, the pH increases gradually as more CH 3 COOH is neutralised and the solution becomes less acidic There is a sharp increase in pH (from 7 to 11) slightly before and after equivalence point Beyond the equivalence point, the pH increases gradually as more NaOH is added The suitable indicators are: Indicator pH Range Colour Change Cresol Red 7.2-8.8 Yellow-Red Phenolpthalein 8.3- 10.00 Colourless - pink

40. Weak Acid- Weak Base Titrations involving a weak acid with a weak base are not normally done This is because the equivalence point cannot be accurately observed