Thermodynamics: Energy, Heat, Temperature, and Phase Changes Chapter 16

16.1 Energy A. Energy – “the capacity to do work or cause the flow of heat” (work = force x distance)

16.1 Energy 1. Kinetic Energy 2. Potential Energy “energy due to motion” KE = 1/2 mv2 Ex. The rock actually falling on Wiley Coyote. 2. Potential Energy Energy due to position or arrangement Ex. The rock actually above his head…levitating there

16.1 Energy 3. Chemical Potential Energy Energy due to chemical bonding Attractions and repulsions due to ionic and covalent bonding B. Law of conservation of energy Energy is not created or destroyed, only transformed Most of the time true except for nuclear reactions

16.1 Energy D. Heat “q”- energy transfer between a system and its surroundings caused by difference in temp Flows High T  Low T stops when system and surrounding the same T K.E. transfer from system to surroundings What was temperature again???

16.1 Energy Measuring Heat Temperature is used to monitor the flow of heat in and out of a system

16.1 Energy 1. Units of Heat calorie (cal) Joule (J) “Quantity of heat that will raise 1.0 g of water 1.0 oC” Based on water (common substance) so it’s easy to calculate Joule (J) Unit of energy used to measure all forms of energy, not just heat SI Unit of heat (our favorite one in Chem) Ex. 60 Watt Light bulb used 60 J/s of energy

16.1 Energy Calorie (Cal): Conversions 1 calorie = 4.184 joules typically known as the food calorie 1 kcal = 1000 calories Conversions 1 calorie = 4.184 joules 1 Calorie = 1000 calories

16.1 Energy Practice: How many joules of heat are there in 325 calories? 400 Calories?

16.1 Energy Practice: How many joules are in 1.11 Calories?

16.1 Energy Practice The average baked potato contains 164 Calories. How much energy is this in Joules?

16.1 Energy Heat in Chemical Reactions and Processess A. Chemical energy and the universe 1. System – the part of the universe under study. 2. Surroundings – the rest of the universe. 3. Universe – the system and the surroundings.

16.1 Energy Exothermic vs. Endothermic Reactions Enthalpy-the stored energy of a system Enthalpy (heat) of a reaction – the change in enthalpy between the products and reactants of a reaction (ΔH)

16.1 Energy Energy diagrams

16.1 Energy Thermochemical Reactions 1. Exothermic reaction   a. 4 Fe (s) + 3 O2 (g)  2 Fe2O3 (s) + 1625 kJ b. 4 Fe (s) + 3 O2 (g)  2 Fe2O3 (s) + Heat c. 4 Fe (s) + 3 O2 (g)  2 Fe2O3 (s) DH = - 1625 kJ

16.1 Energy Endothermic reaction a. NH4NO3 (s) + 27 kJ  NH4+1 (aq) + NO3- (aq) NH4NO3 (s) + Heat  NH4+1 (aq) + NO3- (aq) c. NH4NO3(s)  NH4+1(aq) + NO3-(aq) DH = 27 kJ

16.2 Heat in Chemical Reactions and Processes A. Measuring Heat using a calorimeter Calorimeter- Device used to measure heat, based on the law of Conservation of Energy Energy gained by one substance had to be lost by another

16.2 Heat in Chemical Reactions and Processes Coffee Cup Calorimeter Thermometer -Heat lost by hot solids is gained by water in cup -From mass + temp change of water, you can calculate a quantity of heat H2O Why a Styrofoam Cup? 1) Good insulator 2) Won’t absorb heat for itself

16.1 Energy C. Equation for measuring heat transfer in a Calorimeter q = m x c x ΔT q = heat absorbed or released in joules or calories m = mass of the sample in grams c = specific heat of the substance in joules or calories/g °C ΔT = Tf – Ti change in temperature in Celsius

16.1 Energy F. Specific Heat 1. Definition 2. Units - Amount of energy required to raise 1 gram of substance by 1 degree Celsius 2. Units -

16.1 Energy Table I: Specific Heats of Common Substances at 298 K (25 ˚C) Substance Specific heat J/(g˚C) Water (liquid) 4.184 Water (ice) 2.03 Water (steam) 2.01 Ethanol 2.44 Aluminum 0.987 Granite 0.803 Iron 0.449 Lead 0.129 Silver 0.235 Gold

Examples: 1. How much heat, in joules, is needed to raise the temperature of 5000 grams of water from 20oC to 30oC in calories?

Examples: 2. How many joules of heat are given off when 5.0 grams of water cool from 75oC to 25oC?

Examples: 3. What mass of water can be heated from 15.0oC to 40.0oC with the addition of 3,870 joules?

Examples: 4. What is the specific heat of a metal if a 5.7 gram block of it absorbs 629 joules to raise the temperature from 22 oC to 145 oC?

Examples: 5. What will the final temperature be if 6.50 kilocalories of heat is absorbed by 500 grams of water with an initial temperature of 50.0 oC?   a. First solve for ΔT: b. Then solve for Tf:

16.1 Energy 3. Applications of specific heat Ex. Pot of water on the stove How fast does the pot heat up? The water? Why is water so special?

16.3 Thermochemical Equations C. Heating Curve / Cooling Curve