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Istanbul Bilgi university Engeniring
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The following basic safety rules should be observed at all times in the laboratory:
Beaker Erlenmeyer Flask Volumetric Flask Pipettes Buret Pipette Filler Graduated cylinder Ring stand
Funnel Filter Paper Watch glass Spatulas Stirring rod Dropper or Pasteur pipette Analytical Balance Thermometer Polystyrene cup
PART A. Prepare 100 mL of a 0.15 M aqueous solution of sodium chloride (NaCl).
PART C. The Use of the Buret. !!Empty and clean your buret if there is any solution in it. !!Fill your buret with 0.1 M CuSO 4 solution. !!Adjust the liquid level in the buret to be at 0-mL mark. !!Record the initial buret reading !!Dispense 10.0 mL of CuSO 4 solution into each of four erlenmeyer flasks !!Record the final buret reading each time !!Empty the excess CuSO 4 solution into the graduated cyclinder !!Number your erlenmeyer flasks !!Fill your buret with distilled water and adjust the liquid level to be at 0-mL !!Dispense 10 mL of water into 2nd erlenmeyer flask, 20 mL into 3rd erlenmeyer flask and 30 mL into 4th erlenmeyer flask !!Calculate the molarity of each CuSO 4 solution
3 3
Purpose : To understand the principles of stoichiometry and the concept of theoretical and actual yield. Theory : Mole ratios reactants, products, their physical states etc. for any chemical reaction can be obtained from the balanced equation. 2 Ag 2 O (s) 4 Ag (s) + O 2 (g) AgNO 3 (aq) + NaCl (aq) AgCl (s) + NaNO 3 (aq) For all reactions the reactants are combined in given ratio, and if one of them is in excess, it will remain as unreacted in the medium. The reactant which is completely used up in the reaction is called the limiting reactant , the one which remains is called the excess reactant. The amount of the product should be calculated according to the limiting reactant. The amount of the product calculated from the balanced equation of the reaction is the theoretical yield of that reaction. Most reactions, however, do not go to completion. In practice we always get less. The amount of the product produced at the end of the reaction is the actual yield. In this experiment the following reaction will be carried out and the percent yield will be calculated. FeCl 3 (aq) + 3 K 2 C 2 O 4. H 2 O(aq) 3 KCl (aq) + K 3 Fe (C 2 O 4 ) 3. 3H 2 O (s) Ferric chloride + potassium oxalate potassium chloride + potassium ferritrioxalate trihydrate The complex anion is photosensitive, which means that under the influence of light, Fe (C 2 O 4 ) -^3 undergoes a reduction in which ferric ion, Fe+3, is reduced to ferrous , Fe+2^ , and oxalate ion, C 2 O 4 -^2 , is oxidized to CO 2. 2 Fe (C 2 O 4 ) -^3 2 Fe+2^ + 5 C 2 O 4 -^2 + 2 CO 2 Fe+2^ produced can be detected by adding a solution of potassium ferricyanide, K 3 Fe(CN) 6 , which (^) produces a complex, deep blue in color. Fe+2^ + Fe(CN) 6 -^3 Fe[Fe(CN) 6 ] deep blue color
Procedure : Synthesis of K 3 Fe(C 2 O 4 ) 3. 3H 2 O !Weigh the filter paper (it should be clean and dry) and record its mass. ! !Weigh out 3 g of K 2 C 2 O 4 .H 20 (Mw :184.23 g/mol ) into the beaker and add 9 mL of distilled (^) water and mix until all of the compound dissolves. ! !Measure exactly 4 mL of FeCl 3 (Mw :162.2 g/mol) solution by using your pipet, and add into (^) the beaker. !Put beaker in the ice bath and wait for complete precipitation (approx. 10 minutes). !Put the funnel onto erlenmeyer flask and place the filter paper into the funnel correctly. ! Filter the contents of beaker, lift out the filter paper and try to excess water from it by using a piece of paper towel. ! Write your name on your watch glass, place your filter paper on it and hand it to your instructor. !Complete your calculations while you are waiting for your product to dry. !Weigh the dry product and calculate the percent yield of your reaction. Safety: FeCl 3 : Corrosive, it can cause burns on the skin. Harmful is swallowed or inhaled.
Equipment and Supplies:
an unknown concentration of acetic acid (the ingredient that gives vinegar its sour taste), using phenolphthalein as the indicator in the titration. Titration. In a titration, a buret is used to dispense measured increments of one solution into a known volume of another solution. The object of the titration is the detection of the equivalence point , that point in the procedure where chemically equivalent amounts of the reactants have been mixed. Whether or not the equivalence point comes when equimolar amounts of reactants have been mixed depends on the stoichiometry of the reaction. In the reaction of acetic acid, HC 2 H 3 O 2 , and NaOH, the equivalence point does occur when 1 mole of HC 2 H 3 O 2 has reacted with 1 mole of NaOH. However, in the reaction of H 2 SO 4 and NaOH, the equivalence point occurs when 2 moles of NaOH have reacted with 1 mole of H 2 SO 4. The titration technique can be applied to many types of reactions, including oxidation-reduction, precipitation, complexation, and acid-base neutralization reactions. Although a variety of instrumental methods for detecting equivalence points are now available, it is frequently more convenient to add an indicator to the reaction mixture. An indicator is a substance that undergoes a distinct color change at or near the equivalence point. The point in the titration at which the color change occurs is called the end point. Obviously, the titration will be accurate only if the end point and the equivalence point coincide fairly closely. For this reason, the indicator used in a titration must be selected carefully. Fortunately, a large number of indicators are commercially available and finding the right one for a particular titration is not a difficult task. Acids and Bases. Although several definitions of acids and bases may be given, the classical Arrhenius concept will suffice for this experiment. According to this concept, an acid is a substance that dissociates in water to produce hydrogen ions; a base is a substance that dissociates in water to produce hydroxide ions. The classical Arrhenius acid-base reaction is one in which an acid reacts with a base to form water (from the combination of hydrogen ions and hydroxide ions) and a salt. Such a reaction is called a neutralization reaction. The neutralization reaction of acetic acid, the primary ingredient of vinegar, and sodium hydroxide is as follows: HC 2 H 3 O 2 ( aq ) + NaOH ( aq ) à!H 2 O ( l ) + NaC 2 H 3 O 2 ( aq ) Because the mole ratio of acetic acid to sodium hydroxide is 1:1, the number of moles of acid present in the sample is equal to the number of moles of base that must be added to reach the equivalence point of the titration. This is formulated as:
!Record the initial buret reading of NaOH solution. ! !Titrate the KHP solution with NaOH solution (add NaOH from buret to Erlenmeyer flask slowly dropwise and mix them at the same time until you observe a faint pink color. !When the faint pink color is permanent, stop and record the final buret reading. !Calculate the molarity of NaOH solution. !Repeat the procedure and average the results. Titration of an Unknown. !Deliver 10 ml sample of undiluted acetic acid solution to 100-ml volumetric flask. !Add small amount of distilled water to volumetric flask and swirl the flask to mix its contents. !Dilute the solution to the 100-ml mark. !Use your pipet to deliver 10 ml of diluted acetic acid solution to a clean Erlenmeyer flask. !Add 3 drops of phenolphthalein indicator and swirl the flask. !Titrate the diluted acetic acid solution with your NaOH solution until a faint pink color is obtained. !Calculate the molarity of diluted acetic acid solution. !Repeat the titration steps and average your concentration results. !Calculate the molarity of undiluted acetic acid solution.
Name-Surname and ID No: ………………………...... Date: ………………… Partner’s Name : ………………................................... Section: ………………
Standardization of NaOH Solution Trial 1 Trial 2 Mass of KHP Moles of KHP Initial buret reading, NaOH Final buret reading, NaOH Volume used, NaOH Molarity of NaOH solution Average molarity of NaOH Titration of Diluted Vinegar Volume of undiluted vinegar delivered to 100-mL flask Volume of diluted vinegar to be titrated Trial 1 Trial 2 Initial buret reading, NaOH Final buret reading, NaOH Volume used, NaOH Molarity of acetic acid solution Average molarity of acetic acid solution Average molarity of undiluted acetic acid IMPORTANT! All of the equipment and work desk used in the lab were left clean.
Heats of reactions are usually reported in kJ/mol. Calorimeters are used in determining the heats of reactions. A simple calorimeter is composed of an insulated vessel, a thermometer and a stirrer. Based on the conservation of energy (the first law of thermodynamics), the following equation is used to determine the heats of reactions, Heat absorbed Heat lost Heat lost or released by = or gained by + or gained by the reaction the solution the calorimeter q = – mc soln( T f – T i) + C cal( T f– T i) where m is the mass of the final solution, c soln is the specific heat capacity of the solution, T i is the initial temperature of the reacting medium, T f is the final temperature, C cal is the heat capacity of the calorimeter (vessel, thermometer and stirrer). Procedure: Part A. Heat of Solution for Solid NH 4 Cl !Place the polystyrene cup in a 250 mL beaker to help prevent spills. !Measure 100 ml of distilled water and pour it to the polystyrene cup. !Record the temperature of water ( T i). !Weigh 10 g of NH 4 Cl (Mw: 53.49 g/mol) into an empty, clean and dry beaker. ! !Add the salt and stir to dissolve the compound, more importantly observe the temperature until it remains constant for about 15-20 seconds. !Record the temperature of solution ( T f). ! !Calculate heat of dissolution of NH 4 Cl, standard enthalpy change of this reaction and % error of (^) your experiment.
Part B. Heat of Neutralization for NaOH and HCl !Clean and place the polystyrene cup in a 250 mL beaker to help prevent spills. !Measure 50 mL of 2 M NaOH with a graduated cylinder and place it in the polystyrene cup. !Record the temperature of the solution ( T i). !Measure 50 mL of 2 M HCl with a graduated cylinder and add this solution carefully to the NaOH solution in the cup. !Constantly observe the temperature until it remains constant for about 15-20 seconds. !Record the final temperature of solution ( T f). ! !Calculate heat of neutralization reaction, standard enthalpy change of this reaction and % error of your experiment. FOR YOUR REPORT: For the purpose of this experiment, you may assume that heat loss/gain to the calorimeter is negligible. Part A. Sample Calculations:
