Abstract: acid-base titration is a technique commonly used to determine the moles of acid in a sample by adding a known volume of strong base of a known concentration. The strong base provides the hydroxide ion, to react quantitatively with the acid. The point at which the acid is completely and exactly consumed the known quantity of base is called the equivalence end point and is signalled by a colour change in the solution (end point). This colour change is created by an indicator dye which is extremely sensitive to the presence of even a small excess of aqueous.
From the stoichiometry of the balanced chemical reaction, the number of moles of the unknown acid solution can be determined. If the number of grams unknown acid is measured, the molecular weight can be calculated. Titration is the process of the gradual addition of a standard solution to a second solution until all of the solute in the second solution has completely reacted. Introduction: in the experiment you will be determining the molarity of an unknown acid solution by measuring the volume of a sodium hydroxide solution of known concentration required to neutralize a measured volume of the unknown acid solution.
The indicator is added because there is no observation change that occurs when the neutralization reaction is complete. The indicator used in this experiment, phenolphthalein, is colourless in acidic solution but turns pink when there is an excess of base present, therefore you will know when the reaction is complete when you see the first faint of pink colour in the reaction mixture. Because the pink colour occurs when there is an excess of base present, the lighter the pink colours is at the end the better.
Safety; when handling acids, bases and chemical indicators, wear safety glasses. The skin should be immediately washed with water for several minutes if exposed to any chemical. During the titration maintain a slow, steady stream of titrant. Avoid sudden stops and starts. If the glassware breaks, use a dustpan and brush to clean up. Equipment: Burette Pipette Sodium hydroxide Hydrochloric acid Phenolphthalein Conical flask Boss Clamp Clamp stand White tile Distilled water Funnel Beaker Safety glasses Measuring cylinder Method: (stage 1) Pour 0.1mol dm-3 Sodium hydroxide into a burette, filling it to the 0 mark Using a pipette, accurately measure 25cm30. 1 mol dm-3 hydrochloric acid.
Add phenolphthalein as an indicator to the flask Drop wise add sodium hydroxide to the flask by slowly opening the top of the burette, keep swirling the flask. The reaction is complete when you see the indicator turn from colourless to red. Use a table to record your final volume of sodium hydroxide used, repeat 5 times. Work out the average of best titration results. Remember you can ignore the outliers Results: Initial reading cm3 4 3 3. 3 3. 2 Final burette reading cm3 27 26. 50 26. 80 26. 10 Titre 23 23. 50 23. 80 23. 9 Average: 23. 5cm3 Second method: (stage 2) .
Work out the average of the best titration results- remember to ignore any outliers Weigh the evaporating dish Pipette 25cm3 of 0. 1mol dm-3 hydrochloric acid into the evaporating dish Add the correct amount 0. 1mol dm-3 acid; that is the average of the best titration results. Reduce the volume of the solution by about half by evaporating it on a steam bath as instructed and leave until the following week to crystalize.
Evaporating dish: 50. 94 grams Final mass of evaporating dish: 51. 08 grams Calculating the percentage yield and atom economy: Calculating the percentage yield: How efficiently did we make our copper (ii) sulphate? Actual number of moles= measured mass (g) (51. 08-50. 94) 249. 61 = 2. 395619439 x10-3 We can calculate the percentage yield; % yield= actual number of moles Expected number of moles 100 23. 5/ 1000 0. 1 Prepare a solid organic product Abstract: in this experiment you will synthesize aspirin, purify it and determine the percentage yield.
The purity of the product will be confirmed by qualitative analysis and by measuring its melting point range. Introduction: aspirin is a synthetic organic derived from salicylic acid. Salicylic acid is a natural product found in the bark of the willow tree and was used by the ancient Greek and Native Americans, among others, to counter fever and pain. However it can be bitter to the stomach. The reaction that is used for the synthesis is shown below. In this reaction an excess of acetic anhydride is added to a measured mass of salicylic acid in the presence of a catalyst, sulphuric acid.
The mixture is heated to from the acetylsalicylic acid and acetic acid. After the reaction takes place, water is added to destroy the excess acetic anhydride and cause the product to crystallize. The aspirin is then collected, purified by recrystallization, and its melting temperature measured. Safety: make sure that goggles are worn throughout the entire experiment, this experiment uses salicylic acid, anhydride. The salicylic acid and aspirin may cause irritation to your skin or eyes. Dispose of any excess in a fume cupboard as this can be harmful to the environment.
If you spill any on the benches or floor be sure to wipe it up immediately, and discard of the cloth used to do so. Make sure that if you get any of the chemicals on your skin to wash off immediately as they can irritate the skin and cause burning. Make sure that you wear gloves when completing the experiment to ensure Equipment: (part 1) Salicylic acid 100cm3 conical flask 10cm3 measuring cylinder Ethanoic anhydride Concentrated sulphuric acid in a dripping bottle 400cm3beaker Tripod, gauze and Bunsen burner Thermometer 250cm3 beaker Reduced pressure filtration apparatus.
Filter paper Glass stirring rod Distilled water Spatula Part 2: 25cm3 measuring cylinder Boiling tube Ethanol Thermometer Distilled water 250cm3 beaker 100cm3 conical flask Glass stirring rod A kettle Preparation (part 1) Weigh out approximately 6. 00g of salicylic acid directly into a 100cm3 conical flask Record the mass of salicylic acid used Using a 10cm3 measuring cylinder, add 10cm3 of ethanoic anhydride to the flask and swirl the contents. Add 5 drops of concentrated sulphuric acid to the flask and swirl the mixture in the flask for a few minutes to ensure thorough mixing.
Warm the flask for twenty minutes in a 400cm3 beaker of hot water at approximately 600C. The temperature of the flask should not be allowed to rise above 650C. Allow the flask to cool and pour its contents into a 75cm3 beaker of water, stirring well to precipitate the solid. Filter off the aspirin under reduced pressure, avoiding skin contact Collect the crude aspirin on a double thickness of filter paper and allow it to dry. Part 2: Using a 25cm3 measuring cylinder, measure out 15cm3 of ethanol into a boiling tube Prepare a beaker half-filled with hot water at a temperature of a approximately 750c.
the safest way to do this is to use a kettle of boiling water and add water from the kettle to cold water in the beaker until the temperature is at approximately 750C Use a spatula to add the crude aspirin to the boiling tube and place the tube in the beaker of hot water Stir the contents of the boiling tube until all of the aspirin dissolves into the ethanol Pour the hot solution containing dissolved aspirin into approximately 40cm3 of water in a 100cm3 conical flask.
If a solid separates at this stage, gently warm the contents of the flask in the water bath until solution is complete. Allow the conical flask to cool slowly and white needles of aspirin should separate If no crystals have formed after the solution has cooled to room temperature, you may need to use an ice bath and to scratch the insides of the flask with a glass stirring rod to obtain crystals. Filter off the purified solid under reduced pressure and allow it to dry on filter paper Record the mass of the dry purified solid.