Chapter 6: Thermochemistry

Transcription

1 Chapter 6: Thermochemistry 1. Light the Furnace: The Nature of Energy and Its Transformations a. Thermochemistry is the study of the relationships between chemistry and energy i. This means that we will be shifting our focus from matter to energy b. Vocabulary: i. Energy the capacity to do work ii. Work the result of a force acting through a distance iii. Heat the flow of energy caused by a temperature difference iv. Kinetic energy energy associated with the motion of an object v. Thermal energy energy associated with the temperature of an object vi. Potential energy energy associated with the position or composition of an object vii. Chemical energy energy associated with the relative positions of electrons and nuclei in atoms and molecules viii. Law of Conservation of Energy states that energy cannot be created nor destroyed (but it can be transferred or assume different forms) c. The SI unit of energy is the joule (J) i. d. Other units include: Calorie (C) and kilowatt-hour (kwh) [see table 6.1 for conversations] 2. The First Law of Thermodynamics: There Is No Free Lunch a. Thermodynamics is the general study of energy and its interconversions b. The first law of thermodynamics states the total energy of the universe is constant i. Basically the law of conservation of energy c. The internal energy (E) of a system is the sum of the kinetic and potential energies of all of the particles that compose the system i. This is a state function its value depends only on the state of the system, not on how the system arrived at that state 1

2 ii. So all we care about is the initial and final states In a chemical reaction the final state is the products and the initial state is the reactants d. The sign of ΔE is positive when energy flows into the system and out of the surroundings and the sign of ΔE is negative when energy flows out of the system and into the surroundings e. According to the first law: i. Table 6.3 Sign conventions for q, w, and ΔE* q (heat) + system gains thermal e system loses thermal e w (work) + work done on system work done by system ΔE (change in int. energy) + energy flow into system energy flows out of system f. Homework Problems: i. Page 234 #5, 6, 7, 8, 10, Quantifying Heat and Work a. Remember that thermal energy always flows from matter at higher temperatures to matter at lower temperatures this is why a hot beverage will cool down (because it is transferring heat to the surroundings) i. Also remember that temperature is a measure of thermal energy of a sample of matter ii. This means that heat is directly proportional to the change in temperature 1. This results in a formula: b. C is heat capacity or the quantity of heat required to change the temperature by 1 C and is measured in J/g C 2

3 c. Specific heat capacity is the amount of heat required to raise the temperature of 1 gram of the substance by 1 C and is measured in J/g C i. Table 6.4 for the specific heat capacity of common substances (pg. 213) ii. Molar heat capacity is the amount of heat required to change the temperature of 1 mole of a substance and is measured in J/mol C d. There is a formula that relates heat and specific heat: * e. To calculate work (in particular pressure-volume work which occurs when the force is the result of a force acting through a distance): * f. Homework Problems: i. Page 234 #14, 16, 18, 20, Measuring ΔE for Chemical Reactions: Constant-Volume Calorimetry a. We now have a complete picture of how a system exchanges energy with the surroundings: b. When we have a chemical reaction, the volume is constant, which changes our energy equation to: (where q V is heat at constant volume) i. In a bomb calorimeter (a device used to measure changes in internal energy for combustion reactions) the heat absorbed by the calorimeter is represented by: and because the heat gained by the calorimeter is equal to the heat released by the reaction* 5. Enthalpy: The Heat Evolved in a Chemical Reaction at Constant Pressure a. The enthalpy (H) of a system is defined as the sum of its internal energy and the pressure of its pressure and volume: i. At constant pressure we can simplify the equation: b. The value of ΔH for a chemical reaction is the amount of heat absorbed or evolved in the reaction under constant pressure 3

4 i. An endothermic reaction has a positive ΔH and absorbs heat from the surroundings (feels cold to the touch) ii. An exothermic reaction has a negative ΔH and gives off heat to the surroundings (feels warm to the touch)* c. The enthalpy change for a chemical reaction is also called the enthalpy of reaction or heat of reaction (ΔH rxn ): i. 1. This means when 1 mole of C 3 H 8 that reacts with 5 moles of O 2 to form 3 moles of CO 2 and 4 moles of H 2 O, 2044kJ of heat is emitted* d. Homework Problems: i. Page 234 #26, 28, 29, Constant-Pressure Calorimetry: Measuring ΔH rxn a. For many aqueous reaction, ΔH rxn can be measured using a coffee-cup calorimeter and if you know the specific heat capacity of the solution: i. ii. And since this happens under constant pressure: * b. Homework Problems: i. Page 235 #34, 35, Relationships Involving ΔH rxn a. There are three quantitative relationships between a chemical equation and ΔH rxn : i. If a chemical equation is multiplied by some factor, then ΔH rxn is also multiplied by the same factor 1. For example: ii. If a chemical equation is reversed, then ΔH rxn changes sign 1. For example: 4

5 iii. If a chemical equation can be expressed as the sum of a series of steps, then ΔH rxn for the overall equation is the sum of the heats of reactions for each step 1. For example: 2. This relationship is known as Hess Law* b. Homework Problems: i. #38, 40, Enthalpies of Reaction from Standard Heats of Formation a. There are different standards that we have been and will be talking about: i. Standard State 1. For a gas: the pure gas at a pressure of 1 atm 2. For a liquid/solid: the pure substance in its most stable form at 1 atm and the temp of interest (usually 25 C) 3. For substance in solution: concentration of 1M ii. Standard Enthalpy Change ( ) 1. The change in enthalpy for a process when all reactants and products are in their standard states. The degree sign indicates standard states. iii. Standard Enthalpy of Formation ( ) 1. For a pure compound: when 1 mole of the compound forms from its constituent elements in their standard states 2. For a pure element in its standard state: 3. See table 6.5 for some standard enthalpies of formation (or appendix IIB for a more complete list)* 5

6 b. We can also calculate the standard enthalpy change for a reaction using the following formula: * c. Homework Problems: i. Page 235 #44, 46, 48, 50 All chapter 6 homework problems: #5, 6, 7, 8, 10, 12, 14, 16, 18, 20, 22, 26, 28, 29, 30, 34, 35, 36, 38, 40, 42, 44, 46, 48, 50 (25) Review Problems: #55, 59, 62, 65 6