CH 104 HEATS OF REACTION A calorimeter is

CH 104: HEATS OF REACTION • A calorimeter is used to measure the amount of heat absorbed or released during a chemical reaction. • In today’s experiment you will measure the heat of a reaction in a calorimeter that is made out of 2 Styrofoam coffee cups. • The inner cup holds an aqueous reaction mixture. The outer cup gives additional thermal insulation from the surrounding environment. The thermometer is used to measure the initial and final temperatures of the reaction mixture.

CALORIMETRY • The first law of thermodynamics says that energy, in all its forms, is conserved in all processes. Or the heat lost (qlost) by a system equals the heat gained (qgained) by the surroundings. qlost + qgained = 0 • The fundamental equation of calorimetry says the heat lost by a reaction (qlost = qreaction) is gained by the surrounding water and calorimeter (qgained = qwater + qcalorimeter). qreaction + qwater + qcalorimeter = 0 • q is negative if heat is lost. • q is positive if heat is gained. • q is measured in joules (J).

CALORIMETRY • Again, the fundamental equation of calorimetry is qreaction + qwater + qcalorimeter = 0 • 4. 184 joules (J) are needed to heat 1 gram (g) of water 1 degree Kelvin (K). This is the specific heat of water. Therefore, the heat gained by the water is • And the heat gained by the calorimeter is qcalorimeter = (heat capacity of the calorimeter) x (Tfinal – Tinitial) Where • The mass of water is in grams. • The heat capacity of each calorimeter is unique. Therefore, in today’s experiment you will measure the heat capacity of your calorimeter. • Tfinal is the final temperature of the water in either Kelvin or Celsius. • Tinitial is the initial temperature of the water in either Kelvin or Celsius. • Why can these temperatures be measured in either Kelvin or Celsius? • A Kelvin degree is the same size as a Celsius degree. Therefore, the change in temperature is the same if it is measured in either Kelvin or Celsius.

CALORIMETRY • Again, the fundamental equation of calorimetry is qreaction + qwater + qcalorimeter = 0 • Or qreaction = –(qwater + qcalorimeter) • Then substituting

MEASURING THE HEAT CAPACITY OF YOUR CALORIMETER • In today’s experiment you will measure the heat capacity of your calorimeter. • A student puts 75. 0 g water in a calorimeter. The temperature of this water and the calorimeter is 21. 4° C. Then he adds 74. 5 g of water at 58. 0° C. The final temperature of this mixture is 37. 6° C. The specific heat of water is 4. 184 Jg-1 K-1. • How much heat was lost by the hot water? qwater = 4. 184 Jg-1 K-1 x 74. 5 g x (37. 6° C – 58. 0° C) = – 6. 36 x 103 J • How much heat was gained by the cold water? qwater = 4. 184 Jg-1 K-1 x 75. 0 g x (37. 6° C – 21. 4° C) = 5. 08 x 103 J

MEASURING THE HEAT CAPACITY OF YOUR CALORIMETER • How much heat was gained by the calorimeter? qlost + qgained = 0 qlost by water + qgained by calorimeter = 0 qgained by calorimeter = –(qlost by water + qgained by water) qgained by calorimeter = –(– 6. 36 x 103 J + 5. 08 x 103 J) = 1. 28 x 103 J • What is the heat capacity of the calorimeter? Heat Capacity of the Calorimeter = qgained by calorimeter / (Tfinal – Tinitial) = 1. 28 x 103 J / (37. 6° C – 21. 4° C) = 78. 7 JK-1 Heat capacity MUST be positive. If the calculated heat capacity of your calorimeter is negative, it is wrong.

HEAT OF REACTION • For a solution, the fundamental equation of calorimetry is qreaction = –(qsolution + qcalorimeter) • Or • qreaction is called the heat of reaction. If qreaction is measured at constant pressure, like in our calorimeters which are at atmospheric pressure, then qreaction is also called the change in enthalpy (ΔH). • The change in enthalpy per mole reaction is sometimes written as.

MEASURING THE HEAT OF A REACTION • In today’s experiment you will measure the change in enthalpy ( the following reaction. ) for Mg(s) + 2 HCl(aq) → H 2(g) + Mg. Cl 2(aq) • A student puts 100. m. L of 2. 0 M HCl(aq) (an excess) in a calorimeter. The temperature of this aqueous solution and the calorimeter is 21. 4° C. Then he adds 0. 252 g of Mg(s). The final mass of this solution is 100. g. The final temperature of this solution is 31. 1° C. The specific heat of this solution is 3. 62 Jg-1 K-1. Assume the heat capacity of the calorimeter is 78. 7 JK-1. • How much heat was gained by the solution? qsolution = 3. 62 Jg-1 K-1 x 100. g x (31. 1° C – 21. 4° C) = 3. 5 x 103 J

MEASURING THE HEAT OF A REACTION • A student puts 100. m. L of 2. 0 M HCl(aq) (an excess) in a calorimeter. The temperature of this aqueous solution and the calorimeter is 21. 4° C. Then he adds 0. 252 g of Mg(s). The final mass of this solution is 100. g. The final temperature of this solution is 31. 1° C. The specific heat of this solution is 3. 62 Jg-1 K-1. Assume the heat capacity of the calorimeter is 78. 7 JK-1. Mg(s) + 2 HCl(aq) → H 2(g) + Mg. Cl 2(aq) • How much heat was gained by the calorimeter? qcalorimeter = (heat capacity of the calorimeter) x (Tfinal – Tinitial) qcalorimeter = 78. 7 JK-1 x (31. 1° C – 21. 4° C) = 7. 6 x 102 J

MEASURING THE HEAT OF A REACTION • A student puts 100. m. L of 2. 0 M HCl(aq) (an excess) in a calorimeter. The temperature of this aqueous solution and the calorimeter is 21. 4° C. Then he adds 0. 252 g of Mg(s). The final mass of this solution is 100. g. The final temperature of this solution is 31. 1° C. The specific heat of this solution is 3. 62 Jg-1 K-1. Assume the heat capacity of the calorimeter is 78. 7 JK-1. Mg(s) + 2 HCl(aq) → H 2(g) + Mg. Cl 2(aq) • How much heat was evolved by the reaction? qreaction = –(qsolution + qcalorimeter) qreaction = –(3. 5 x 103 J + 7. 6 x 102 J) = – 4. 3 x 103 J

MEASURING THE HEAT OF A REACTION • A student puts 100. m. L of 2. 0 M HCl(aq) (an excess) in a calorimeter. The temperature of this aqueous solution and the calorimeter is 21. 4° C. Then he adds 0. 252 g of Mg(s). The final mass of this solution is 100. g. The final temperature of this solution is 31. 1° C. The specific heat of this solution is 3. 62 Jg-1 K-1. Assume the heat capacity of the calorimeter is 78. 7 JK-1. Mg(s) + 2 HCl(aq) → H 2(g) + Mg. Cl 2(aq) • How many moles of Mg(s) reacted? The atomic weight of Mg is 24. 305 g/mole. 0. 252 g of Mg /24. 305 g of Mg mole-1 = 0. 0104 moles of Mg

MEASURING THE HEAT OF A REACTION • A student puts 100. m. L of 2. 0 M HCl(aq) (an excess) in a calorimeter. The temperature of this aqueous solution and the calorimeter is 21. 4° C. Then he adds 0. 252 g of Mg(s). The final mass of this solution is 100. g. The final temperature of this solution is 31. 1° C. The specific heat of this solution is 3. 62 Jg-1 K-1. Assume the heat capacity of the calorimeter is 78. 7 JK-1. Mg(s) + 2 HCl(aq) → H 2(g) + Mg. Cl 2(aq) • What is the heat of reaction in k. J/mole?

SAFETY • Give at least 1 safety concern for the following procedures that will be used in today’s experiment. • Heating with a flame. • Injury from a burn or causing a fire. Be careful. Do not wear loose clothing or long hair. Glass can shatter when heating; wear your goggles at all times. • Using 2. 0 M HCl, and Mg(s). • These are irritants. Wear your goggles at all times. Immediately clean all spills. If you do get either of these in your eye, immediately flush with water. • Generating H 2(g). • Hydrogen gas is flammable. Do NOT generate H 2(g) until all the Bunsen burners in the laboratory are extinguished. Wear your goggles at all times. • Your laboratory manual has an extensive list of safety procedures. Read and understand this section. • Ask your instructor if you ever have any questions about safety.

SOURCES • Barnes, D. S. , J. A. Chandler. 1982. Chemistry 111 -112 Workbook and Laboratory Manual. Amherst, MA: University of Massachusetts. • Mc. Murry, J. , R. C. Fay. 2004. Chemistry, 4 th ed. Upper Saddle River, NJ: Prentice Hall. • Petrucci, R. H. 1985. General Chemistry Principles and Modern Applications, 4 th ed. New York, NY: Macmillan Publishing Company.