CH 6: Thermochemistry

6.1 Nature of Energy Thermochemistry – study of energy changes during chemical reactions –Aspects of thermochemistry are studied in both physics and chemistry

6.1 Energy – the capacity to do work or produce heat Law of conservation of energy states that energy can be converted from one form to another, but it cannot be created or destroyed. –Energy of the universe is constant!

6.1 Two forms of energy 1. Potential energy – stored energy Energy of position or composition Examples 2. Kinetic energy – energy of motion Heat, light, electricity In this chapter our focus will (eventually) be on the heat aspects of thermochemistry.

6.1 Consider the diagram on page 237. Ball A has potential energy due to its position –Some of this energy is released as heat as A rolls down the hill –Ball A hits Ball B and work is done as B moves up the incline Work = force acting over a distance

6.1 The original potential energy of A is equal to the new potential energy of B plus the heat released as friction. Energy can be released as: –Heat (friction) –Work (ball rolling)

6.1 The total energy released when ball A rolls down the hill is fixed – how it’s released is not. –How much is released as heat vs. work depends upon the conditions –Total energy released does not depend on the pathway.

6.1 Total energy change is a state function. –State function is a property of a system that depends only on its current state not its past See last paragraph on page 237 –Energy of a system is a state function. Heat and work are not state functions – they depend on the path taken

6.1 Chemical Energy When studying chemical change we consider the system and the surroundings. –System includes the reactants and products of a given reaction –Surroundings are everything else in the universe! Including the water if the reaction is done in solution

Chemical Energy Exothermic reactions – energy flows out of the system –Products are of lower potential energy than the reactants Energy is released to the surroundings Energy lost by the system = energy gained by the surroundings – There’s a loss of energy by the system….  Energy of the system is negative

Chemical Energy Endothermic reactions – energy flows into the system –Products are of higher potential energy than the reactants The system absorbs energy from the surroundings The energy gained by the system = the energy lost by the surroundings There’s a gain of energy by the system…….  Energy of the system is positive

Chemical Energy Exothermic:  E < 0 –System _______ energy Endothermic:  E > 0 –System ________ energy

More on Energy of the System The energy of the system can change as a result of 2 factors: –Heat (q) –Work (w) Change of energy of the system = heat flowing in/out of system + work being done to/by the system  E = q + w

 E of the System Heat (q) –When heat flows into the system, q > 0 Endothermic –When heat flows out of the system, q < 0 Exothermic Work (w) –When work is done to the system, w > 0 –When work is done by the system, w < 0

More on Work Most common form of work done by a system is the expanding or compressing of a gas, called pressure – volume work –Work – force applied over a distance Moving an object a distance = work

Pressure Volume Work Consider a gas in a cylinder with a movable piston on top Work = force x distance Force = pressure x area Distance = change in height of gas in cylinder (  h) Work = P x A x  h  Volume Page 241

Pressure Volume Work Work = P x  V –The sign ( +/-) on work is assigned so that: When the gas expands, it is doing work on the surroundings (w < 0) When the gas is compressed, work is done on the system (w > 0)

Pressure Volume Work Final version of the equation that shows both magnitude and sign on work: W = - P x  V When the gas expands  V is positive and w < 0 (work is done by the system) When the gas is compressed  V is negative and w > 0 (work is done on the system)

6.2 Entahalpy Enthalpy (H) –Enthalpy is defined as: H = E + PV E is the energy of the system P is the pressure of the system V is the volume of the system

Enthalpy In chemistry we consider enthalpy at constant pressure –After much math this results in the formula:  q p  H is called the heat of the reaction  H is a measure of the flow of heat (q) into/out of the system at constant pressure

Enthalpy When heat leaves the system  H < 0 –Exothermic process When heat enters the system  H > 0 –Endothermic process

Finally – the Applications! CH O 2  CO H energy  H = kJ 1.Is the reaction exothermic or endothermic? 2.How much energy will be ___________ if 6.50 grams of CH 4 is burned at constant pressure? Next…..#44 on page 277

Calorimetry The heat changes associated with a chemical reaction are often measured in a calorimeter.

Calorimetry Exothermic: The heat released by the reaction is used to heat up a known quantity of water. More heat released the hotter the water gets Endothermic: The heat absorbed by the reaction comes from the water More heat absorbed the colder the water gets

Terms Terms all describe the energy needed to heat or cool some amount of a given substance –Heat Capacity (C) –Specific heat capacity –Molar heat capacity

Terms Heat Capacity (C) –Amount of energy needed to raise the temperature of a substance by 1 0 C –Units: J/ 0 C –Substance is an entire/specific object…e.g.

Terms Specific Heat Capacity –Amount of energy needed to raise the temperature of 1 gram of a substance by 1 0 C –Units: J/g 0 C –See page 245

Terms Molar Heat Capacity –Amount of energy needed to raise the temperature of 1 mole of a substance by 1 0 C –Units: J/mol 0 C Molar heat capacity = specific heat x molar mass J/mol 0 C = J/g 0 C x g/mol

Calculating Heat Capacities Heat Capacity = heat absorbed  T Specific Heat = energy (mass) (  T) Page 277: 52, 54

Determining Specific Heat of a Metal Lab demonstration of experiment 27:E, page 347 of the lab manual –Assume the heat capacity of the ccc is 0 J/g 0 C