chapter 5 thermochemistry - university of victoria - …web.uvic.ca/~chem102/lee/chapter5.pdf ·...
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Chapter 5 Thermochemistry
• The Nature of Energy
• The First Law of Thermodynamics
• Enthalpy
• Enthalpies of Reaction
• Calorimetry
• Hess‟s Law
• Enthalpies of Formation
The ability to do work or transfer heat
5.1 The Nature of Energy
An object can possess energy in two forms, Kinetic Energy and Potential Energy.
Potential Energy
As the cyclist descends the hill potential energy is converted into other forms of
energy, primarily kinetic energy, the energy of motion
Forces arising from electrical charges are important when dealing with atoms
and molecules.
Chemical energy of substances is due to the potential energy stored in the
arrangement of atoms
Kinetic Energy
The energy of motion
Units of Energy
An older but still widely used non SI unit is the calorie
System and Surroundings
The SI unit of energy is the Joule (J)
Systems may be closed, open or isolated
In a chemical reaction, the reactants and products are the system. The
container and everything beyond it are considered the surroundings.
Closed System
Isolated System
Open system
We experience energy changes in the form of work or heat
Transferring Energy
Heat
Work
Heat is the energy transferred from a hotter object to a colder one
Sample Exercise 5.1
Energy can be neither created nor destroyed
5.2 The First Law of Thermodynamics
Internal Energy
The internal energy of a system is the sum of all the kinetic and potential
energies of all its components
The change in internal energy is given by
When a system undergoes any chemical or physical change, the magnitude
and sign of the accompanying change in internal energy (ΔE), is given:
Relating ΔE to heat and work
When heat is added to a system or work is
done on a system, its internal energy
increases
Note the sign conventions for ΔE, q and w
Sample exercise 5.2
Endothermic, Endo means “into”
Endothermic and Exothermic Processes
Exothermic, Exo means “out of”
Usually we have no way of knowing the exact value of the internal energy, E,
of a system, simply too complex, it does have a fixed value depending on
conditions.
State Functions
Internal Energy, E, is an example of a state function
A state function is the property of a system that is determined by specifying its
state (pressure, temperature)
The value depends only on the present state, not on how it arrived there
Internal Energy is a state function but q and w are not:
Work
When a process occurs in an open container, commonly the only work
done is a change in volume of a gas pushing on the surroundings (or
being pushed on by the surroundings).
When a reaction is carried out in a
constant volume container
so if PΔV is the only type of work done ΔE = q + w =
5.3 Enthalpy
When the system changes at constant pressure the change in Enthalpy is
given by:
If a process takes place at constant pressure (as
the majority of processes we study do) and the
only work done is this pressure-volume work,
we can account for heat flow during the
process by measuring the enthalpy of the
system.
Since ΔE = q + w and w = -PΔV (-ve system does work piston moves up)
then………
Since ΔE = q + w and w = -PΔV (-ve system does work piston moves up)
then………we can substitute these into the enthalpy expression
So at constant pressure
Remember sign convention
ΔH > 0 Endothermic system gains heat from the surroundings
ΔH < 0 Exothermic system gives out heat to the surroundings
The enthalpy of a chemical reaction, sometimes called the heat of reaction
ΔHrxn is given by the equation:
5.4 Enthalpies of Reaction
The magnitude of ΔH is directly proportional to the amount of reactant
consumed in the process.
The enthalpy change for a
reaction is equal in magnitude
and opposite in sign, to ΔH for the
reverse reaction
The enthalpy change for a reaction depends on the state of the reactants and
products
The enthalpy change for a reaction gives an indication as to whether a
reaction is likely to be „spontaneous‟ or thermodynamically favourable.
ΔH can be determined experimentally by measuring the heat flow
accompanying a reaction at constant pressure
5.5 Calorimetry
A calorimeter is a device used measure heat
accompanying a reaction at constant pressure
Heat Capacity and Specific Heat
The temperature change resulting from an object when it absorbs a certain
amount of heat is determined by its Heat Capacity, C
Specific Heat can be determined experimentally:
Specific Heat
Specific heat is heat capacity expressed on a per gram basis
Cs =q
m T
Sample exercise 5.5
This simple „coffee cup‟ calorimeter is not sealed, the
reaction occurs at constant atmospheric pressure
Constant-Pressure Calorimetry
The heat gained by the solution qsoln must be equal in magnitude and opposite
in sign to qrxn
Sample Exercise 5.6
Designed to study combustion reactions of (usually)
organic compounds.
Constant-Volume (Bomb) Calorimetry
The heat capacity of the Calorimeter Ccal is determined separately by combusting
a known mass of a compound that releases a known quantity of heat
Sample exercise 5.7 ΔE =ΔH
It is possible to calculate ΔH for a reaction using tabulated ΔH values from
other reactions, rather than make calorimetric measurements every time
5.6 Hess’s Law
Hess’s Law States: If a reaction carried out
in a series of individual steps, ΔH for the
overall reaction will equal the sum of the
individual enthalpy changes
Sample exercises 5.8 and 5.9
An enthalpy of formation, Hf, is defined as the enthalpy change for the
reaction in which a compound is made from its constituent elements in their
elemental forms
5.7 Enthalpies of formation
Standard enthalpy of formation
The standard enthalpy of formation, Hof, is defined as the enthalpy of
formation of one mole of compound when all the reactants and products are
in their standard states
We can uses Hess‟ Law to calculate ΔHorxn for any reaction in which we
know the ΔHof values for all reactants and products
Using Enthalpies of formation to calculate Enthalpies of Reaction
C3H8(g) + 5O2(g) → 3CO2(g) + 4H2O(l)
We can write this equation as the
sum of 3 formation reactions
Use the table of ΔHof
We can use Hess‟s Law to obtain the result that the standard enthalpy change
of a reaction is the sum of the standard enthalpies of formation of the products
MINUS the standard enthalpies of formation of the reactants
C3H8(g) + 5O2(g) → 3CO2(g) + 4H2O(l)
Sample exercises 5.11 and 5.12