Unit 6: Thermochemistry

1
Unit 6 Thermochemistry

• Introduction
• Heat and Work
• Specific Heat
• Enthalpy (DH)
• Enthalpy of Reaction
• Phase Diagram

2
Introduction

• Most daily activities involve processes that
  either use or produce energy
• Activities that produce energy
• Metabolism of food
• Burning fossil fuels
• Activities that use energy
• Photosynthesis
• Pushing a bike up a hill
• Baking bread

3
Introduction

• Thermodynamics
• The study of energy and its transformations
• Thermochemistry
• A branch of thermodynamics
• The study of the energy (heat) absorbed or
  released during chemical reactions

4
Introduction

• Objects can have two types of energy
• Kinetic energy
• Energy of motion
• Thermal energy
• The type of kinetic energy a substance possesses
  because of its temperature
• Potential energy
• Energy of position
• stored energy resulting from the attractions
  and repulsions an object experiences relative to
  other objects

5
Introduction

• Units of Energy
• SI unit joule (J)
• 1 J the kinetic energy of a 2 kg mass moving at
  a speed of 1 m/s
• A very small quantity
• Kilojoule (kJ)
• 1 kJ 1000 J

6
Introduction

• Units of Energy (cont)
• Calorie (cal)
• Originally defined as the amount of energy needed
  to raise the temperature of 1g of water from
  14.5oC to 15.5oC.
• 1 cal 4.184 J (exactly)
• Kilocalorie (kcal)
• 1 kcal 1000 cal

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Introduction

• Example Convert 3.02 kJ to J.
• Given 3.02 kJ
• Find J

8
Introduction

• Example Convert 725 cal to kJ.
• Given 725 cal
• Find kJ

9
Introduction

• When using thermodynamics to study energy
  changes, we generally focus on a limited,
  well-defined part of the universe.
• System
• The portion of the universe singled out for study
• Surroundings
• Everything else

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Introduction
The system is usually the chemicals in the
flask/reactor.
The system
The flask and everything else belong to the
surroundings.
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Introduction

• Open system
• A system that can exchange both matter and energy
  with the surroundings
• Closed system
• A system that can exchange energy with the
  surroundings but not matter

A cylinder with a piston is one example of a
closed system.
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Introduction

• In a closed system energy can be gained from or
  lost to the surroundings as
• Work
• Heat
• Work
• Energy used to cause an object to move against a
  force
• Lifting an object
• Hitting a baseball

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Introduction

• Heat
• The energy used to cause the temperature of an
  object to increase
• The energy transferred from a hotter object to a
  cooler one
• Energy
• The capacity to do work or to transfer heat

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Introduction

• The potential energy of a system can be converted
  into kinetic energy and vice versa.
• Energy can be transferred
• back and forth between the system and the
  surroundings
• as work and/or heat.

Potential energy Kinetic energy
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The First Law of Thermodynamics

• Although energy can be converted from one form to
  another and can be transferred between the system
  and the surroundings
• Energy cannot be created or destroyed.
• (First Law of Thermodynamics)
• Any energy lost by the system must be gained by
  the surroundings and vice versa.

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The First Law of Thermodynamics

• The First Law of Thermodynamics can be used to
  analyze changes in the Internal Energy (E) of a
  system.
• The sum of all kinetic and potential energy of
  all components of a system
• For molecules in a chemical system, the internal
  energy would include
• the motion and interactions of the molecules
• the motion and interactions of the nuclei and
  electrons found in the molecules

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The First Law of Thermodynamics

• Internal Energy
• Extensive property
• depends on mass of system
• Influenced by temperature and pressure
• Has a fixed value for a given set of conditions
• State function

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The First Law of Thermodynamics

• The internal energy of a system is a state
  function.
• A property of the system that is determined by
  specifying its condition or its state in terms of
  T, P, location, etc
• Depends only on its present condition
• Does not depend on how the system got to that
  state/condition

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The First Law of Thermodynamics

• The internal energy of a system can change when
• heat is gained from or lost to the surroundings
• work is done on or by the system.
• The change in the internal energy
• D E Efinal - Einitial
• DE change in internal energy
• Efinal final energy of system
• Einitial initial energy of system

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The First Law of Thermodynamics

• If Efinal gt Einitial,
• DE gt0 (positive)
• the system has gained energy from the
  surroundings.
• endergonic

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The First Law of Thermodynamics

• The decomposition of water is endergonic (DE gt
  0)
• 2 H2O (l) 2 H2 (g) O2 (g)

H2 (g), O2 (g)
Energy must be gained from the surroundings.
final
E
H2O (l)
initial
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The First Law of Thermodynamics

• If Efinal lt Einitial,
• DE lt 0 (negative)
• the system has lost energy to the surroundings.
• exergonic

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The First Law of Thermodynamics

• The synthesis of water is exergonic (DE lt 0)
• 2 H2 (g) O2 (g) 2 H2O (l)

Energy is lost to the surroundings in this
reaction.
initial
E
final
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The First Law of Thermodynamics

• The internal energy of a system can change when
  energy is exchanged between the system and the
  surroundings
• Heat
• Work
• The change in internal energy that occurs can be
  found
• D E q w
• Where q heat
• w work

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The First Law of Thermodynamics

• By convention
• q positive
• Heat added to the system
• w positive
• Work done on the system by the surroundings
• q negative
• Heat lost by the system
• w negative
• Work done by the system on the surroundings

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The First Law of Thermodynamics

• Example Calculate the change in internal energy
  of the system for a process in which the system
  absorbs 140. J of heat from the surroundings and
  does 85 J of work on the surroundings.
• Given system absorbs 140. J heat
• system does 85 J work
• Find D E

140. J - 85J
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The First Law of Thermodynamics

• D E q w
• D E 140 J (-85 J)
• D E 55 J