Post on 09-Feb-2018
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:
Potential Energy
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
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 by:
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
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:
WorkWhen 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). This is called P-V work
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 in chemistry) and the only work done is this P-V 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 the 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 chemical 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
Specific Heat can be determined experimentally:
Specific Heat
Specific heat is heat capacity expressed on a per gram basis
s =q
m × ΔT
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
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
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
An enthalpy of formation, ΔHf, 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 formationThe standard enthalpy of formation, ΔHo
f, 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)