Chemistry, A molecular approach

Acid-Base Titration
When performing quantitative chemical analysis, it is important to know the quantities
and concentrations of your analytical reagents very accurately. Standardization is the
process of determining a very accurate value for the concentration (molarity) of a solution.
The most common method of standardization of aqueous acids and bases is titration.
Titration is a procedure by which a solution with a known concentration of one reactant (the
titrant) is gradually added to a solution with an unknown concentration of another reactant
(the analyte). The titrant is added to the analyte until all of the analyte is consumed, at which
point (called the endpoint) the titration is halted and the volume of titrant delivered is
recorded. The concentration of the analyte can then be easily calculated using this volume.
The reaction between a strong acid and strong base in aqueous solution is essentially
quantitative (complete). The net ionic equation for the reaction is:
H+(aq) + OH-(aq)
Thus, when you use an acidic titrant to standardize a basic analyte, at the endpoint
the number of moles of H+ in the added titrant will be the same number of moles of OHoriginally present in the analyte. If you know the molarity of H+ the titrant, then the molarity of
the analyte can be determined using stoichiometry of that reaction.
Because there is no visually apparent change in the solution at the endpoint, another
reagent called an indicator is used. An acid-base indicator will change color dramatically
when a certain concentration of H+ ions is reached. Different indicators change color at
different H+ concentrations (i.e. different pH). The indicator we will be using in this
experiment is phenolphthalein. It is colorless in acidic solutions (pH < 7), but turns bright pink or fuchsia at about pH 8.2. Although in a titration of a strong base with a strong acid the endpoint will be reached at pH 7, the next drop of acidic solution added will drastically change pH. Therefore, when we see the phenolphthalein change color, we know that the endpoint was reached when the previous drop of titrant was delivered. Phenolphthalein is an excellent indicator for titrations involving strong acids and strong bases, but it may not be the best choice of indicator for titrations involving weak acids or weak bases. If a strong base is titrated with a weak acid, the endpoint will usually be at pH 9 or higher, well past the point that phenolphthalein changes color. A better choice of indicator for this application would be thymolphthalein, which changes from colorless to blue at about pH 9.3. If a weak base is titrated with a strong acid, the endpoint will be reached at a pH below 7, well before phenolphthalein changes color. Better indicators for this type of titration would be bromothymol blue or methyl red, which change color at pH 6.0 and 4.4 respectively. It is best to choose an indicator so that the color change occurs very near to the pH at which the endpoint would be reached in a given acid-base titration. Page 1 of 8 Name: Date: In this experiment you will standardize (accurately determine the concentration of) a NaOH solution by titrating the solution using a previously standardized solution of HCl. Not only is HCl a strong acid, meaning it dissociates completely into H+ and Cl- ions when dissolved in water, it is also monoprotic, meaning that every mole of HCl provides 1 mole of H + ions to the solution. NaOH is a strong base and also monobasic, meaning is supplies 1 mole of OH- for every mole of NaOH in aqueous solution. In the second part of the experiment you will use the NaOH solution you standardized to determine the molar mass of an unknown acid by titration. Some of the unknown acids available to you will be monoprotic while others will be diprotic and triprotic. Diprotic acids produce 2 moles of H+ ions per mole of dissolved acid so it takes 2 moles of OH- to neutralize 1 mole of the acid. Triprotic acids produce 3 moles of H+ ions per mole of dissolved acid so it takes 3 moles of OH- to fully neutralize 1 mole of the acid. Since it is impossible to know if your acid is monoprotic, diprotic, or triprotic beforehand, you will be told which type of acid your unknown is. Procedure Part A: Standardization of NaOH solution You are provided a stock solution of 6 M NaOH. Using a small graduated cylinder measure out about 7 mL of this solution, place it into an 800 mL beaker, and dilute using about 400 mL of deionized water. Mix the solution thoroughly using a stirring rod. Thoroughly clean and dry a 125 mL Erlenmeyer flask. Into this flask, draw about 75 mL of the standardized HCl solution. This amount should be sufficient for all your titration needs, so only take the amount you need and do not waste it. The concentration will be about 0.1 M, but it will have been determined to 3 significant figures prior to the experiment, so be sure to remember that in your calculations. Prepare a 50 mL buret by first checking it for cleanliness. If it is dirty, clean it with detergent, rinse with tap water, and finally rinse with deionized water. A clean buret will leave an unbroken film of water along the interior without beading or droplets. When you are satisfied that your buret is clean, rinse it with your titrant solution to be used. To do this, add a few milliliters of the titrant to the buret and swirl it around, making sure it touches all of the interior surface, then drain it into a beaker. Repeat the rinse process 2 more times. Then, with the stopcock closed, fill the buret with titrant solution to a level above the top graduation (the 0 mL mark). Open the stopcock carefully and release the solution dropwise until you are sure that the tip of the buret is filled completely with your titrant and there are no air bubbles or leaks. If necessary, refill the buret to over the 0 mL line and then carefully open the stopcock, releasing solution until the bottom of the meniscus is on the 0 mL line. If you have difficulty seeing the meniscus or the graduation markings, it is helpful to place a piece of paper or index card behind the buret. It is important to prepare the buret this way to make sure that the concentration of titrant is the exact same in the buret as it is in the stock bottle. Since the HCl solution has been standardized and the concentration was determined very accurately, we want to maintain this level of accuracy throughout the experiment. The addition of even 1 mL of extra water will change the concentration of titrant enough to alter your results. Page 2 of 8 Name: Date: Use the above procedure to prepare two 50 mL burets, one to hold and deliver HCl solution and the other to hold and deliver NaOH solution. Label the acid buret A and the base buret B. Place buret A in the left side of your buret clamp and place buret B on the right side. Read and record the levels of the two burets to an accuracy of 0.01 mL. If you are just starting out, it is best for the initial level to be 0.00 mL. Use buret A to deliver about 25 mL of the HCl solution into a clean 250 mL Erlenmeyer flask, then add about 25 mL of deionized water along with 2 – 3 drops of phenolphthalein solution. Place the Erlenmeyer flask under buret B, being careful not to touch the tip of the buret to the flask. A piece of white paper placed under the flask will help to see any subtle color changes when you start to titrate. While carefully and continuously swirling the Erlenmeyer flask, begin adding the NaOH solution in buret B to the HCl solution intermittently, a few mL at a time. Notice the pink color that quickly appears then disappears as the two solutions meet. As you get closer to the endpoint the pink color will persist for longer and longer. When you think you are about to reach the endpoint, add the NaOH solution drop by drop, swirling the flask after each drop. You have reached the endpoint when the solution is a faint pink that persists for at least 20 seconds. Record the final levels of both burets A and B to 0.01 mL If you have an intense pink solution, then you have overshot the endpoint. In this case, add a few drops of the HCl solution until the color disappears. Then, carefully add NaOH drop by drop as you did before until you achieve a proper endpoint. Record the final level of both burets to 0.01 mL. Now add about 10 mL more of the standardized HCl solution from buret A to the Erlenmeyer flask containing the titrated solution. The pink color should disappear. Titrate this solution with NaOH from buret B as you did before to reach the endpoint. Again record the final level of both burets. Repeat this process once more. You have now completed 3 titrations of NaOH solution with a standardized HCl solution. The total volume of HCl solution delivered should be about 25, 35, and 45 mL respectively. You can now calculate the concentration of your NaOH solution for each of the three titrations using the stoichiometry for that reaction. For each calculation, be sure to use the total volumes of acid and base that were added at that point. If two of these concentrations agree to within 1 – 2 %, proceed to the next part of the experiment. If not, you may need to repeat the titration until 3 measurements do agree. Page 3 of 8 Name: Date: Part B: Determination of the Molar Mass of a Solid Acid Weigh approximately 0.200 g of unknown solid acid. Carefully place the sample in a clean, but not necessarily dry, 250 mL Erlenmeyer flask. Add 50 mL of deionized water and 2 – 3 drops of phenolphthalein then swirl the solution until the solid acid has dissolved. Fill the sodium hydroxide buret. Read the level of the NaOH solution to the nearest 0.01 mL and record that value as the initial volume of NaOH. Titrate the acid solution with NaOH to the first permanent faint pink color. Record the level of NaOH as the final volume. Weigh a new sample of solid acid, place it in a clean Erlenmeyer flask and repeat the titration. The unknown solid acid can be a monoprotic, diprotic, or triprotic; that is, there can be 1, 2, or 3 moles of ionizable H+ ions per mol of solid acid. Ask your instructor for the number of moles of ionizable H+ ions per mole of your unknown solid acid and record that number below. To calculate the molar mass of your unknown, note that the number of moles of sodium hydroxide required per mole of solid acid to reach the endpoint in the titration depends upon the number of moles of ionizable H+ ions in the solid. 1 ionizable H+ ion 2 ionizable H+ ions 3 ionizable H+ ions HA + NaOH H2B + 2 NaOH H3C + 3 NaOH NaA + H2O Na2B + 2 H2O Na3C + 3 H2O Page 4 of 8 Name: Date: Acid-Base Titration Data and Calculations Name Date Part A: Standardization of NaOH solution. Data: Trial 1 Trial 2 Initial level, HCl buret mL Final level, HCl buret mL Initial level, NaOH buret mL Final level, NaOH buret mL Trial 3 mL mL mL mL Calculations: Trial 1 Total volume HCl solution Trial 2 mL Molarity of standardized HCl Trial 3 mL mL M Total volume NaOH solution mL mL mL Molarity of NaOH solution M M M Average Molarity of NaOH solution M Page 5 of 8 Name: Date: Part B: Determination of the Molar Mass of a Solid Acid Number of moles of ionizable H+ ions per mole of solid acid Unknown Number MNaOH (The average molarity of NaOH from Part A.) Data: Trial 1 Trial 2 Mass of sample g g Initial volume, NaOH solution mL mL Final volume, NaOH solution mL mL Calculations: Trial 1 Trial 2 Volume of NaOH solution used mL mL Moles of NaOH (OH-) used mol mol Moles of H+ in sample mol mol Moles of solid acid in sample mol mol Molar mass of solid acid g/mol g/mol Average Molar Mass g/mol Page 6 of 8 Name: Date: Acid-Base Titration Pre-lab Assignment Page 7 of 8 Name: __________________________________ CH131 Chapter 6 – Bomb Calorimetry A bomb calorimeter measures changes in internal energy (∆E) for combustion reactions. Hint: you can always ignore the heat capacity of the small sample of hydrocarbon because it is negligible compared to the heat capacity of the calorimeter. 1. When 1.010 grams of sucrose (C12H22O11) undergoes combustion in a bomb calorimeter, the temperature rises from 24.92°C to 28.33°C. Find ∆Erxn for the combustion of sucrose in kJ/mole sucrose. The heat capacity of the bomb calorimeter is 4.90 kJ/°C. 2. When 1.550 grams of liquid hexane (C6H14) undergoes combustion in a bomb calorimeter, the temperature rises from 25.87°C to 38.13°C. Find ∆Erxn for the reaction in kJ/mole hexane. The heat capacity of the bomb calorimeter is 5.73 kJ/°C. 3. The combustion of toluene (C7H8) has a ∆Erxn of -3.91 × 103 kJ/mole. When 1.55 grams of toluene undergoes combustion in a bomb calorimeter, the temperature rises from 23.12°C to 37.57°C. Find the heat capacity f the bomb calorimeter. Chapter 6 –Calorimetry 1. Are the following processes exothermic or endothermic? Is energy, in the form of heat, absorbed or released by the reaction? What is the sign of ∆H for each process? a. When solid KBr is dissolved in water, the solution gets colder. b. Natural gas is burned in a furnace c. When concentrated H2SO4 is added to water, the solution gets very hot. d. sweat evaporating from the skin 2. A coffee-cup calorimeter initially contains 125 grams of water at 24.2°C. Then 10.5 grams of KBr, also at 24.2°C, is added to the water. After the KBr dissolves, the final temperature is 21.1°C. The specific heat of the solution is 4.18 J/g°C. a. Calculate q for this reaction in J and kJ. b. Calculate the number of moles of KBr involved in this reaction. c. Calculate the enthalpy change (∆H) for this reaction in kJ/mole of KBr. d. Is this reaction exothermic or endothermic? How do you know? 3. When 9.55 grams of solid NaOH dissolves in 100.0 grams of water in a coffee-cup calorimeter, the temperature rises from 23.6°C to 47.4°C. The specific heat of the solution is 4.18 J/g°C. a. Calculate q for this reaction in J and kJ. b. Calculate the number of moles of NaOH involved in this reaction. c. Calculate the enthalpy change (∆H) for this reaction in kJ/mole of NaOH. d. Is this reaction exothermic or endothermic? How do you know? 4. In a coffee-cup calorimeter, 50.0 mL of 0.100 M AgNO3 at 22.6°C are mixed with 50.0 mL of 0.100 M HCl, also at 22.6°C. The final temperature of the resulting solution is 27.9°C. The final solution has a combined mass of 100.0 grams and a specific heat capacity of 4.18 J/g°C. a. Write a balanced equation for this reaction. b. Write the net ionic equation for this reaction. b. How many moles of solid product are formed? c. Calculate the enthalpy change that accompanies this reaction (∆H) in kJ/mole of solid product. d. Is this reaction exothermic or endothermic? How do you know? 5. How many joules are needed to change the temperature of 10.0 g of Al from 20.0oC to 45.0oC? The specific heat of q = mass x Cs x ∆T Al is 0.90 J/g oC. Heat Transfer 6. 10.0 g of a metal powder were heated to 90.0oC and then poured into a beaker containing 50.0 g of water at 20.0oC. The temperature of the water – metal mixture rose to 22.5oC. Assuming no loss of heat to the surroundings from the metal and water, calculate the specific heat of the metal. The specific heat of water is 4.18 J/g oC. (The metal does not dissolve in the water.) Calorimetry Using the q = mass x Cs x ∆T formula to calculate the ∆H of a reaction 7. When a 3.88 g sample of solid ammonium nitrate dissolves in water to produce 60.0 mL of solution in a coffee cup calorimeter, the temperature drops from 23.0 oC to 18.4 oC. Calculate ∆H (in kJ/mole NH4NO3) for the solution process. (The molar mass of NH4NO3 = 80.1 g/mole) The density of the solution is 1.02 g/mL. Assume that the specific heat of the solution is the same as that of pure water. NH4NO3 (s)  NH4+ (aq) + NO3- (aq) Using a bomb calorimeter to calculate the ∆H of a reaction qcal = Ccal x ∆T 8. HC2H3O2 (l) + 2 O2(g)  2 CO2(g) + 2 H2O (l) When 10.0 g of HC2H3O2 were combusted in a bomb calorimeter in the presence of oxygen, the temperature of the bomb calorimeter rose from 25.00oC to 35.81oC. If the heat capacity of the calorimeter and its contents is 13.43 kJ/oC, what is the enthalpy change of the reaction? (The molar mass of HC2H3O2 is 60.0 g/mole) Chapter 6 – Enthalpy of Formation Using enthalpies of formation, calculation ∆Hf° (in kJ) for each of the following three reactions: + H2O (g) → NH3 (g) + O2 (g) 1) NO (g) 2) The combustion of isopropanol (hint: you may want to write a balanced equation first!) Is the reaction balanced? 2 3) SO3 (g) + H2O (l) → H2SO4 (l) Is the reaction balanced? Why do the following two reactions have different values of ∆H? H2 (g) + ½ O2 (g) → H2O (l) ∆Hf° = -286 kJ/mole 2 H2 (g) + O2 (g) → H2O (l) ∆H = -572 kJ What is the difference between ∆Hf° and ∆H? For which of the following reactions does ∆H = ∆Hf°? Be sure to explain your answer(s). C(s) + O2 (g) → CO2 (g) ∆H = = -394 kJ ½ N2 (g) + O2 (g) → NO2 (g) ∆H = = 34 kJ 2 Fe (s) + 3/2 O2 (g) → Fe2O3 (s) ∆H = = -824 kJ H2 (g) + Cl2 (g) → 2 HCl (g) ∆H = = -92 kJ 3 Chapter 6 – Heat and Heat Transfer 1. How much energy (heat), in kJ, is required to warm 1.50 kg of sand from 25.0°C to 95.5°C? 2. a. Calculate the energy, in kJ, needed to raise the temperature of a 850 gram block of aluminum from 22.8°C to 94.6°C. b. Calculate the molar heat capacity of aluminum. (hint: think about the units of molar heat capacity) 3. Suppose that 23.4 grams of gold is initially at 27.1°C. What is the final temperature of the gold after it absorbs 2.35 kJ of heat? 4. A kilogram of water and a kilogram of aluminum metal are each warmed to 75.0°C and placed in two identical (and separate) containers. One hour later, the two containers are opened and the temperature of each substance is measured. The aluminum metal cooled to 35.0°C while the water cooled only to 66.0°C. Explain this difference (no calculations are needed) 5. If two objects, A and B, of different temperature come in direct contact, what is the relationship between the heat lost by one object and the heat gained by the other? 6. Suppose 1.0 gram of substance A at 20°C is added to 1.0 gram of substance B at 40°C. Will the final temperature of the mixture be 30°C? Explain why or why not. 7. A 32.5 gram iron rod, initially at 22.7°C, is submerged into an unknown mass of water at 63.2°C. The final temperature of the mixture upon reaching thermal equilibrium is 59.5°C. What is the mass of the water? 8. A 15.0 gram sample of silver metal is heated to 96.7°C and then dropped into 55.0 grams of water at 23.0°C. Assuming that all the heat lost by the silver is absorbed by the water, calculate the final temperature of the silver and water. 4 Chapter 6 Internal Energy, Work, and Heat 1) Which of the following signs on q and w represent a system that is doing work on the surroundings, as well as losing heat to the surroundings? A) q = - , w = - B) q = +, w = + C) q = -, w = + D) q = +, w = - E) None of these. 2) Calculate the amount of heat (in kJ) required to raise the temperature of a 79.0 g sample of ethanol from 298.0 K to 385.0 K. The specific heat capacity of ethanol is 2.42 J/g°C. A) 57.0 kJ B) 16.6 kJ C) 73.6 kJ D) 28.4 kJ E) 12.9 kJ 3) Determine the specific heat capacity of an alloy that requires 59.3 kJ to raise the temperature of 150.0 g alloy from 298 K to 398 K. A) 4.38 J/g°C B) 2.29 J/g°C C) 3.95 J/g°C D) 2.53 J/g°C E) 1.87 J/g°C 4) A sample of copper absorbs 43.6 kJ of heat, resulting in a temperature rise of 75.0 °C, determine the mass (in kg) of the copper sample if the specific heat capacity of copper is 0.385 J/g°C. A) 1.51 kg B) 6.62 kg C) 1.26 kg D) 7.94 kg E) 3.64 kg 5) A 4.98 g sample of aniline (C6H5NH2, molar mass = 93.13 g/mol) was combusted in a bomb calorimeter with a heat capacity of 4.25 kJ/°C. If the temperature rose from 29.5°C to 69.8°C, determine the value of ΔH°comb for aniline in kJ/mol. A) +7.81 × 103 B) -3.20 × 103 C) +1.71 × 103 D) -1.71 × 103 E) -7.81 × 103 6) A piece of iron (mass = 25.0 g) at 398 K is placed in a styrofoam coffee cup containing 25.0 mL of water at 298 K. Assuming that no heat is lost to the cup or the surroundings, what will the final temperature of the water be? The specific heat capacity of iron = 0.449 J/g°C and water = 4.18 J/g°C. A) 348 K B) 308 K C) 287 K D) 325 K E) 388 K 7) A 6.50-g sample of copper metal at 25.0°C is heated by the addition of 84.0 J of energy. Calculate the final temperature of the copper. The specific heat capacity of copper is 0.38 J/g.K. A) 29.9 °C B) 25.0 °C C) 9.0 °C D) 59.0 °C E) 34.0 °C 8) The specific heat of copper is 0.385 J/(g ∙ °C). If 34.2 g of copper, initially at 24.0°C, absorbs 4.689 kJ, what will be the final temperature of the copper? A) 24.4°C B) 26.8°C C) 356°C D) 380°C E) 297oC 5

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