Chapter 5 Thermochemistry CHEMISTRY The Central Science 9th Edition Chapter 5 Thermochemistry David P. White Prentice Hall © 2003 Chapter 5

The Nature of Energy Kinetic Energy and Potential Energy Kinetic energy is the energy of motion: Potential energy is the energy an object possesses by virtue of its position. Potential energy can be converted into kinetic energy. Example: a bicyclist at the top of a hill. Prentice Hall © 2003 Chapter 5

The Nature of Energy Kinetic Energy and Potential Energy Electrostatic potential energy, Ed, is the attraction between two oppositely charged particles, Q1 and Q2, a distance d apart: The constant  = 8.99  109 J-m/C2. If the two particles are of opposite charge, then Ed is the electrostatic repulsion between them. Prentice Hall © 2003 Chapter 5

The Nature of Energy Units of Energy SI Unit for energy is the joule, J: We sometimes use the calorie instead of the joule: 1 cal = 4.184 J (exactly) A nutritional Calorie: 1 Cal = 1000 cal = 1 kcal Prentice Hall © 2003 Chapter 5

The Nature of Energy Systems and Surroundings System: part of the universe we are interested in. Surroundings: the rest of the universe. Prentice Hall © 2003 Chapter 5

The Nature of Energy Transferring Energy: Work and Heat Force is a push or pull on an object. Work is the product of force applied to an object over a distance: Energy is the work done to move an object against a force. Heat is the transfer of energy between two objects. Energy is the capacity to do work or transfer heat. Prentice Hall © 2003 Chapter 5

The First Law of Thermodynamics Internal Energy Internal Energy: total energy of a system. Cannot measure absolute internal energy. Change in internal energy, Prentice Hall © 2003 Chapter 5

The First Law of Thermodynamics Relating DE to Heat and Work Energy cannot be created or destroyed. Energy of (system + surroundings) is constant. Any energy transferred from a system must be transferred to the surroundings (and vice versa). From the first law of thermodynamics: when a system undergoes a physical or chemical change, the change in internal energy is given by the heat added to or absorbed by the system plus the work done on or by the system: Prentice Hall © 2003 Chapter 5

The First Law of Thermodynamics

The First Law of Thermodynamics Exothermic and Endothermic Processes Endothermic: absorbs heat from the surroundings. Exothermic: transfers heat to the surroundings. An endothermic reaction feels cold. An exothermic reaction feels hot. Prentice Hall © 2003 Chapter 5

The First Law of Thermodynamics State Functions State function: depends only on the initial and final states of system, not on how the internal energy is used. Prentice Hall © 2003 Chapter 5

State Functions

Zn(s) + 2H+(aq)  Zn2+(aq) + H2(g) Enthalpy Chemical reactions can absorb or release heat. However, they also have the ability to do work. For example, when a gas is produced, then the gas produced can be used to push a piston, thus doing work. Zn(s) + 2H+(aq)  Zn2+(aq) + H2(g) The work performed by the above reaction is called pressure-volume work. When the pressure is constant, Prentice Hall © 2003 Chapter 5

Enthalpy

Enthalpy Enthalpy, H: Heat transferred between the system and surroundings carried out under constant pressure. Enthalpy is a state function. If the process occurs at constant pressure, Prentice Hall © 2003 Chapter 5

Enthalpy Since we know that We can write When DH, is positive, the system gains heat from the surroundings. When DH, is negative, the surroundings gain heat from the system. Prentice Hall © 2003 Chapter 5

Enthalpy Prentice Hall © 2003 Chapter 5

Enthalpies of Reaction For a reaction: Enthalpy is an extensive property (magnitude DH is directly proportional to amount): CH4(g) + 2O2(g)  CO2(g) + 2H2O(g) DH = -802 kJ 2CH4(g) + 4O2(g)  2CO2(g) + 4H2O(g) DH = -1604 kJ Prentice Hall © 2003 Chapter 5

Enthalpies of Reaction When we reverse a reaction, we change the sign of DH: CO2(g) + 2H2O(g)  CH4(g) + 2O2(g) DH = +802 kJ Change in enthalpy depends on state: H2O(g)  H2O(l) DH = -88 kJ Prentice Hall © 2003 Chapter 5

Calorimetry Heat Capacity and Specific Heat Calorimetry = measurement of heat flow. Calorimeter = apparatus that measures heat flow. Heat capacity = the amount of energy required to raise the temperature of an object (by one degree). Molar heat capacity = heat capacity of 1 mol of a substance. Specific heat = specific heat capacity = heat capacity of 1 g of a substance. Prentice Hall © 2003 Chapter 5

Calorimetry Constant Pressure Calorimetry Atmospheric pressure is constant! Prentice Hall © 2003 Chapter 5

Calorimetry Constant Pressure Calorimetry Prentice Hall © 2003 Chapter 5

Calorimetry Bomb Calorimetry (Constant Volume Calorimetry) Reaction carried out under constant volume. Use a bomb calorimeter. Usually study combustion. Prentice Hall © 2003 Chapter 5

Hess’s Law Hess’s law: if a reaction is carried out in a number of steps, H for the overall reaction is the sum of H for each individual step. For example: CH4(g) + 2O2(g)  CO2(g) + 2H2O(g) H = -802 kJ 2H2O(g)  2H2O(l) H = -88 kJ CH4(g) + 2O2(g)  CO2(g) + 2H2O(l) H = -890 kJ Prentice Hall © 2003 Chapter 5

Hess’s Law Note that: H1 = H2 + H3 Prentice Hall © 2003 Chapter 5

Enthalpies of Formation If 1 mol of compound is formed from its constituent elements, then the enthalpy change for the reaction is called the enthalpy of formation, Hof . Standard conditions (standard state): 1 atm and 25 oC (298 K). Standard enthalpy, Ho, is the enthalpy measured when everything is in its standard state. Standard enthalpy of formation: 1 mol of compound is formed from substances in their standard states. Prentice Hall © 2003 Chapter 5

Enthalpies of Formation If there is more than one state for a substance under standard conditions, the more stable one is used. Standard enthalpy of formation of the most stable form of an element is zero. Prentice Hall © 2003 Chapter 5

Enthalpies of Formation Prentice Hall © 2003 Chapter 5

Enthalpies of Formation Using Enthalpies of Formation of Calculate Enthalpies of Reaction We use Hess’ Law to calculate enthalpies of a reaction from enthalpies of formation. Prentice Hall © 2003 Chapter 5

Enthalpies of Formation Using Enthalpies of Formation of Calculate Enthalpies of Reaction For a reaction Prentice Hall © 2003 Chapter 5

Foods and Fuels Foods Fuel value = energy released when 1 g of substance is burned. 1 nutritional Calorie, 1 Cal = 1000 cal = 1 kcal. Energy in our bodies comes from carbohydrates and fats (mostly). Intestines: carbohydrates converted into glucose: C6H12O6 + 6O2  6CO2 + 6H2O, DH = -2816 kJ Fats break down as follows: 2C57H110O6 + 163O2  114CO2 + 110H2O, DH = -75,520 kJ Prentice Hall © 2003 Chapter 5

Foods and Fuels Foods Fats: contain more energy; are not water soluble, so are good for energy storage. Prentice Hall © 2003 Chapter 5

Foods and Fuels Fuels In 2000 the United States consumed 1.03  1017 kJ of fuel. Most from petroleum and natural gas. Remainder from coal, nuclear, and hydroelectric. Fossil fuels are not renewable. Prentice Hall © 2003 Chapter 5

Foods and Fuels Prentice Hall © 2003 Chapter 5

Foods and Fuels Fuels Fuel value = energy released when 1 g of substance is burned. Hydrogen has great potential as a fuel with a fuel value of 142 kJ/g. Prentice Hall © 2003 Chapter 5

End of Chapter 5: Thermochemistry Prentice Hall © 2003 Chapter 5