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)