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UNIT A: THERMODYNAMICS CHAPTER ELEVEN&TWELVE Energy Demands & Resources Personal Use of Chemical Energy Food (energy from photosynthesis) Fossil Fuels (combustion reactions) Access through cellular respiration Heat home, cook, transportation Alternative forms: try to be more eco-friendly Solar Wind Water DO YOU REMEMBER?? THE LAW OF CONSERVATION OF ENERGY During physical and chemical processes, energy may change form, but it may never be created nor destroyed. If a chemical system gains energy, the surroundings lose energy If a chemical system loses energy, the surroundings gain energy Examples: When octane (C3H8, the main component of gasoline) is burned in your car engine, chemical bond energy (potential energy) is converted into mechanical energy (pistons moving in the car engine; kinetic energy) and heat. When we turn on a light switch, electrical energy is converted into light energy and, you guessed it, heat energy. DO YOU REMEMBER?? EXOTHERMIC ENDOTHERMIC A change in a chemical A change in chemical energy energy where energy/heat EXITS the chemical system Results in a decrease in chemical potential energy where energy/heat ENTERS the chemical system Results in an increase in chemical potential energy DO YOU REMEMBER?? SYSTEMS Types of Systems System: the part of the universe we are interested in studying Surroundings: the rest of the universe Open: exchange matter & energy Closed: doesn’t exchange matter, but does energy Isolated: doesn’t exchange matter or energy open closed isolated An Introduction to Thermodynamics Kinetic Energy (Ek) is related to the motion of an entity Molecular motion can by translational (straight-line), rotational and vibrational Chemical Potential Energy (Ep) is energy stored in the bonds of a substance and relative intermolecular forces Thermal Energy is the total kinetic energy of all of the particles of a system. Increases with temperature. Symbol (Q), Units (J), Formula used (Q=mcΔT) Temperature is a measure of the average kinetic energy of the particles in a system Heat is a transfer of thermal energy. Heat is not possessed by a system. Heat is energy flowing between systems. An Introduction to Thermodynamics Which has more thermal energy, a hot cup of coffee or an iceberg? An iceberg! Put very roughly, thermal energy is related to the amount of something you have multiplied by the temperature. Let's assume your iceberg is at the freezing point of water - 0 degrees Celsius (~273 Kelvin). Now your cup of coffee might be 75 degrees Celsius (~350 Kelvin). 350 isn't a whole lot more than 270, but an iceberg is thousands of times larger than a cup of coffee. Even though the iceberg is at a lower temperature, it contains more thermal energy because the particles are moving and it's much larger than the cup of coffee. An Introduction to Thermodynamics Heat energy transferred will be related to the temperature change of the system. It takes different amounts of heat energy to raise the temperature of 1 g of different substances by 1 oC. This number is called the specific heat capacity (c), and is measured in units of J g C Thermal Energy Calculations There are three factors that affect thermal energy Q = mcΔt Mass (m) – unit is g or kg Type of substance (c) – unit is J/goC or kJ/kgoC Temperature change (Δt) Water’s Specific Heat Capacity Water has a c value of 4.19 J g C (can also use 4.19kJ/kgoC) This means that it takes 4.19 J of heat to raise the temperature of 1 g of water by 1oC Water has a very large c compared to most other common substances. Consider a bathtub and a teacup of water! They both contain water with the same heat capacity, but it would take much longer to heat up the water in the bathtub! Thermal Energy Calculations After coming in from outside, a student makes a cup of instant hot chocolate by heating water in a microwave. What is the gain in thermal energy if a cup (250 mL) of tap water is increased in temperature from 15oC to 95oC? (1mL of water = 1g) Thermal Energy Calculations A backpacker uses an uncovered pot to heat lake water on a single-burner stove. If the water temperature rises from 5.0oC to 97oC, and the water gains 385 kJ of thermal energy, what is the mass of water heated? Thermal Energy Calculations #2 A 750 g sample of ethanol absorbs 23.4 kJ of energy. Its initial temperature was 35oC. The specific heat capacity of ethanol is 2.44 J/goC . What is the final temperature of the ethanol sample? Thermal Energy Calculations Experiments have shown that a thermal energy change is affected by mass, specific heat capacity, and change in temperature. What happens to the thermal energy if a) The mass is doubled? b) The specific heat capacity is divided by two? c) The change in temperature is tripled? d) All three variables are doubled? Thermal Energy Calculations If the same quantity of energy were added to individual 100g samples of water, aluminum, and iron, which substance would undergo the greatest temperature change? Explain. (see page 3 of data booklet) How do we measure Q? With a simple laboratory calorimeter, which consists of an insulated container made of three nested polystyrene cups, a measured quantity of water, and a thermometer. The chemical is placed in or dissolved in the water of the calorimeter. Energy transfers between the chemical system and the surrounding water is monitored by measuring changes in the water temperature. “Calorimetry is the technological process of measuring energy changes of an isolated system called a calorimeter” Includes: Thermometer, stirring rod, stopper or inverted cup, two Styrofoam cups nested together containing reactants in solution Comparing Q’s Negative Q value Positive Q value An exothermic change An endothermic change Heat is lost by the system The temperature of the surroundings increases and the temperature of the system decreases Heat is gained by the system The temperature of the system increases and the temperature of the surroundings decreases Example: Cold Pack Question Tips: “What heat is required?” Example: Hot Pack Question Tips: “How much energy is released?” Other Calorimetry assumptions. . . • All the energy lost or gained by the chemical system is gained or lost (respectively) by the calorimeter; that is, the total system is isolated. • All the material of the system is conserved; that is, the total system is isolated. • The specific heat capacity of water over the temperature range is 4.19 J/(g•°C). (** IN YOUR DATA BOOK) • The specific heat capacity of dilute aqueous solutions is 4.19 J/(g•°C). • The thermal energy gained or lost by the rest of the calorimeter (other than water) is negligible; that is, the container, lid, thermometer, and stirrer do not gain or lose thermal energy. Calorimetry Questions •Remember to keep the system and the surroundings separate in your calculations. Because we assume the system is isolated, then we assume that: 1. the energy lost by the system = energy gained by the surroundings (water) OR 2. the energy gained by the system = the energy lost by the surroundings. Calorimetry Questions A 15.0 g piece of copper at 100oC is placed in a calorimeter with 100.0mL of water at 25.00oC. The final temperature of the water and copper is 25.35oC. What is the specific heat capacity of copper? Calorimetry Questions Suppose a piece of iron with a mass of 21.5 g at a temperature of 100.0oC is dropped into an insulated container of water. The mass of the water is 132.0 g and its temperature before adding the iron is 20.0oC. What will be the final temperature of the system? (ciron = 0.449 J/goC) Calorimetry Questions A student collected the following data from a calorimetry lab: Mass of empty cup 2.31 g Mass of cup + water 180.89 g Mass of cup + water + metal 780.89 g Initial temperature of water 17.0oC Initial temperatuer of metal 52.0oC Final temperature of system 27.0oC What is the specific heat capacity of the metal? Calorimetry Questions Suppose we mixed 40.0 g of hot water at 60oC with 100.0 g of cold water at 30.0oC. What is the final temperature?