Download - Unit 6: Thermochemistry
Unit 6: Thermochemistry
Introduction Heat and Work Specific Heat Enthalpy ( Enthalpy of Reaction Phase Diagram
Introduction
Most daily activities involve processes that either use or produce energy:Activities that produce energy
Metabolism of foodBurning fossil fuels
Activities that use energy:PhotosynthesisPushing a bike up a hillBaking bread
Introduction
ThermodynamicsThe study of energy and its
transformations
Thermochemistry:A branch of thermodynamicsThe study of the energy (heat)
absorbed or released during chemical reactions
Introduction
Objects can have two types of energy:Kinetic energy
Energy of motionThermal energy
The type of kinetic energy a substance possesses because of its temperature
Potential energyEnergy of position“stored” energy resulting from the attractions and repulsions an object experiences relative to other objects
Introduction
Units of EnergySI 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
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
Introduction
Example: Convert 3.02 kJ to J.
Given: 3.02 kJFind: J
J = 3.02 kJ x J = 3.02 kJ x 1000 J1000 J = 3020 J = 3020 J1 kJ1 kJ
1 kJ = 1000 J
Introduction
Example: Convert 725 cal to kJ.
Given: 725 calFind: kJ
J = 725 cal x J = 725 cal x 4.184 J4.184 J x x 1 k J1 k J = 3.03 kJ = 3.03 kJ1 cal1 cal 1000 J1000 J
1 cal = 4.184 J1 kJ = 1000 J
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
Introduction
The system
The system is usually the chemicals in the flask/reactor. The flask and everything else belong to the surroundings.
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.
Introduction In a closed system
energy can be gained from or lost to the surroundings as:
WorkHeat
Work:Energy used to cause
an object to move against a forceLifting an object Hitting a baseball
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
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
work
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.
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
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
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
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 energyE = Efinal - Einitial
E = change in internal energyEfinal = final energy of systemEinitial = initial energy of system
The First Law of Thermodynamics
If Efinal > Einitial,
E >0 (positive) the system has gained energy
from the surroundings.
endergonic
The First Law of Thermodynamics
The decomposition of water is endergonic (E > 0):
2 H2O (l) 2 H2 (g) + O2 (g)
H2 (g), O2 (g)
H2O (l)
E
Energy must be gained from the
surroundings.
final
initial
The First Law of Thermodynamics
If Efinal < Einitial,
E < 0 (negative) the system has lost energy to
the surroundings.
exergonic
The First Law of Thermodynamics
The synthesis of water is exergonic (E < 0)
2 H2 (g) + O2 (g) 2 H2O (l)
H2 (g), O2 (g)
H2O (l)
E
Energy is lost to the
surroundings in this
reaction.
initial
final
The First Law of Thermodynamics
The internal energy of a system can change when energy is exchanged between the system and the surroundingsHeatWork
The change in internal energy that occurs can be found:
E = q + w
Where q = heatw = work
The First Law of Thermodynamics
By convention:q = positive
Heat added to the systemw = positive
Work done on the system by the surroundings
q = negativeHeat lost by the system
w = negativeWork done by the system on the surroundings
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: E
+ 140. J- 85J
The First Law of Thermodynamics
E = q + w
E = +140 J + (-85 J)
E = +55 J