Temperature, Heat, and Heating Curves I. Temperature A. What is temperature? 1. Heat and Temperature are not the same. 2. Definition: Temperature is a measure of average kinetic energy of the particles of a sample of matter. 3. Kinetic energy(KE) = energy of motion 4. KE = ½ m v2, where m = mass of particle and v = velocity of particle 5. So, temperature reflects the average velocity of the particles in a sample of matter a. If temp is high, particles have high velocity. b. If temp is low, particles have low velocity. B. Measuring temp 1. Two common scales are used in chemistry: Celsius and Kelvin 2. Celsius a. Abbreviation: ⁰C b. Based on freezing and boiling points of water a. 0 ⁰C is freezing point of water b. 100 ⁰C boiling point of water 3. Kelvin a. Abbreviation: K b. Based on absolute 0 a. Absolute 0 = point at which particles cease movement, i.e. point at which particles’ velocity equals 0 4. Kelvin vs. Celsius a. Conversion: K = 273.15 + ⁰C b. Example a. What is the temp in K at room temperature(25⁰C)? b. K = 273.15 + 25 = 298 K c. NOTE: a single degree in Celsius is the same size as a single degree in Kelvin. The scales just begin at different places. a. Absolute 0 in K = 0 K b. Absolute 0 in⁰C = -273.15⁰C II. Heat A. What is heat? 1. Definition: a. Heat is the transfer of energy between two objects due to a temperature difference between the objects. b. Another way to think of it: Heat is the energy used to cause the temperature of an object to increase. 2. Abbreviation: q 3. Heat always flows from the hot object to the cooler object. B. Measuring heat 1. Most common SI unit of heat is the joule (J). 2. 1 J = Kg x m2/s2 III. Specific heat A. What is specific heat? 1. Definition: Specific heat is the amount of energy required to raise the temp of 1 g of a substance by 1 degree K. 2. Significance: a substance with a high specific heat requires more heat to raise its temperature than a substance with a low specific heat. 3. Specific heat is unique to the substance. Water has a different specific heat than copper. 4. Specific heat is unique to the state of the substance. Specific heat of liquid water is different than that for gaseous water. B. Abbreviation 1. cp or s, depending on the book 2. Since s is used in AP Chem, I will use s. C. Units: 1. J/(g K) or J/(g °C) 2. since a degree K = a degree ⁰C, then J /(g K) = J/(g ⁰C) IV. Calculations with heat, specific heat, and temperature. A. The big equation: q = ms∆T 1. Heat transferred equals the mass of the substance times its specific heat times its change in temperature. 2. Note: ∆ means change 3. ∆T = Tfinal – T initial = final temperature of the object – the initial temperature of the object 4. The units for the equation are: q in J, m in g, s in J/(g °C), and ∆T in K or ⁰C. B. Example 1: How much heat is required to raise the temp of 10. g of water from 50. ⁰C to 75 ⁰C? The s for liquid water is 4.18 J/(g ⁰C). 1. q = ms∆T 2. q = 10.0g x 4.18 J/(g ⁰C) x (75.0-50.0 ⁰C) = 1040 J C. Example 2: What is the specific heat of a 214 g sample, if 24315 J are required to raise its temperature 52.7 degrees °C? 1. Rearrange the equation: s = q/m∆T 2. s = 24315J/( 214 g x 52.7°C) 3. s = 2.16 J/ g °C D. NOTE: This equation cannot be used to determine q, if a phase change occurs in the problem(i.e. liquid water heating to gasesous water). V. Phase Changes A. Types of Phase Changes 1. Solid to liquid: melting, fusion 2. Liquid to gas: vaporization 3. Solid to gas: sublimation 4. Liquid to solid: freezing 5. Gas to liquid: condensation 6. Gas to solid: deposition B. Heating Curves – ask Mr. Smith for example – What happens when a substance is heated. 1. Solid: heat is added and temp of the solid increases 2. Fusion: At phase change, the added heat goes into the phase change and temp does not change(heat goes into breaking intermo forces) 3. Liquid: heat increases KE of the liquid; 4. Vaporization: At bp, heat is used to change from the liquid state to the gas state – temp is not increased. 5. Gas: increased heat will increase KE of particles and temp will rise. 6. Reverse is true in other direction C. Heats of Phase Changes 1. Explanation a. To change phases, intermo forces must be formed or overcome. b. Heat is consumed or liberated during these process c. Heat is transferred until all molecules change phase d. No heat is available during these changes to increase temp! 2. Heats a. Heat of fusion = - heat of freezing b. Heat of vaporization = - heat of condensation c. Heat of sublimation = - heat of deposition D. Calculation Heat Change for a Heating Curve 1. How much heat is required to change ice at -15 °C into steam at 115°C. 2. This is a 5 step problem. 3. Step 1: Raising Temp of the Ice: heat is added to raise the temp to 0°C, the melting pt. a. Use q=msΔT b. Heat transferred = mass of ice x specific heat of ice x temp change of ice. c. s ice = 2.092 J/g °C 4. Step 2: Phase change from ice to water - heat is added to change phase from ice to water. a. Use q= mΔH fusion b. Heat transferred= mass of water x heat of fusion of water c. Heat of fusion of water = 334J/g 5. Step 3: Raising Temp of Water: heat is added to raise temp to 100°C, the boiling pt. a. Use q=msΔT b. Heat transferred= mass of water x specific heat of water x temp change of the water c. s water = 4.184 J/g °C 6. Step 4: Phase change from water to steam a. Use q= mΔH vaporization b. Heat transferred= mass of water x heat of vaporization of water c. Heat of vaporization of water = 2257 J/g 7. Step 5: Raising the Temp of Steam: heat is added to raise the temp of the steam from 100 °C to 115°C. a. Use q= msΔT b. Heat transferred = mass of steam x specific of steam x change of temp of steam c. s steam = 1.841 J/g°C 8. Total amount of heat transferred is the sum of Steps 1-5.
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