CAcT Home
Chemical Energy
Skills to develop
- Discuss the energy aspect of chemical reactions
- Identify chemical reactions that provide energy
- Calculate the amount of energy accompanying a chemical reaction
- Identify exothermic and endothermic changes
- Estimate work performed due to volume change under constant pressure,
including the application of ideal gas law
- Explain enthalpy and standard enthalpy of a reaction
- Calculate the energy change using standard enthalpy of reaction and
standard enthalpy of formation.
Chemical Energy
Propane, C3H8, natural gas, CH4,
and phosphorous, P4 react with oxygen O2,
and these reactions release energy in the form of heat and light.
No doubt, according to the principle of conservation of energy,
energy is required to reverse the reactions. Thus, energy stored
in chemicals (compounds) and energy released or absorbed in chemical
reactions are called chemical energy, which also covers topics
such as bond energy, ionization potential, electron affinity,
electronegativity, lattice energy, etc.
For example, at standard conditions, the combustion of 1.0 mole hydrogen
with oxygen releases 285.8 kJ of energy. We represent the reaction.
H2(g) + 1/2 O2 -> H2O (l),
dH = -285.8 kJ/mol
where dH represent the heat (or
enthalpy) of reaction, and a negative value
means that the heat is released. Usually, dH is represented by
DH in textbooks, but using dH
notation is much less work on Internet documents.
For the reverse reaction, 285.8 kJ/mol is required, and the sign for
dH value changes.
H2O (l) -> H2(g) + 1/2 O2,
dH = +285.8 kJ/mol
| A Chemical Energy Level Diagram
|
|---|
------------H2(g) + 1/2O2
|
| |
286 kJ | | -286 kJ
| |
| ¯
------------ H2O
|
We can also use an energy level diagram to show the relative
content of energy. The energy content of H2(g) + 0.5 O2
is 285.8 kJ higher than a mole of water, H2O.
Oil, gas, and food are often called energy by the news media, but more
precisely they are sources of (chemical) energy -- energy stored in
chemicals with a potential to be released in a chemical reaction.
The released energy performs work or causes physical and chemical changes.
It is obvious that the amount of energy released in a chemical reaction
is related to the amount of reactants. For example, when the amount is
doubled, so is the amount of energy released.
2 H2(g) + O2 -> 2 H2O (l),
dH = -571.6 kJ/mol
Example 1 shows the calculation when the amount of reactants is only
a fraction of a mole.
Example 1
How much energy is release when a balloon containing 0.15 mole of hydrogen
is ignited in the air?
Solution
The amount released is 0.15 mol * 285.8 kJ/mol = 42.9 kJ
Discussion
The sudden release of energy causes an explosion.
Endothermic and Exothermic Reactions
A reaction that releases energy is called exothermic reaction.
Energy is released in the form of heat, light and (pressure-volume) work.
For example, when methane or propane is oxidized by O2,
the heat released causes the gas to expand (explosion in some cases);
releasing heat & light and doing work at the same time. In this case, the
energy source came from chemical reactions instead of at the expense of
internal energy
described in the previous module.
Endothermic reactions absorb energy, and in all cases, the energy
is supplied from another source, in the form of electrical energy,
heat or light.
Pressure-volume Work in Chemical Reactions
Many chemical reactions involve gases, and when a gas is formed, it displaces
other gases by pushing them out against a pressure. Work, defined in
Newtonian physics as a force times the distance along the force direction,
is performed in such an action. The work is called the pressure-volume
(P-V) work, which is a form of energy and it must be analyzed and its
quantity included in chemical energy calculations.
The SI units for pressure are N m-2 and that of the volume is
m3. Pressure times volume gives the unit of N m, which is
the definition of joule,
1 Pa * 1 m3 = 1 N m-2 m
= 1 N m
= 1 J
Since 1 atm = 101300 Pa, and 1 L = 0.001 m3. Thus,
1 atm L = 101.3 J.
The P-V work under constant pressure (P) is simply the pressure
times the change in volume dV.
w = - P dV
This method applies to reactions that produce gases, which are released
into the atmosphere. When work is done by the system, the work is negative,
as the formula indicates. In reactions where gases are consumed to produce
liquid or solid, work is done on the system by its environment.
The work is positive.
In cases the pressure varies, an integral approach is required to evaluate
the pressure volume work.
w = - Ó d (P V)
= - Ó P dV - (the integral of) V dP
The negative sign is retained, because the work done by the system
is negative. However, the integral of P V work depends on the
path, and we will not get into the detail discussion at this stage.
Example 2
In the reaction to produce oxygen,
KClO3(s) = KCl(s) + 3/2 O2(g),
calculate the pressure-volume work done by 8.2 g of KClO3.
Solution
The molar mass of KClO3 is 123.5 g/mol and 8.2 g is 0.067 mol.
Thus, the amount of oxygen produced is 0.10 (= 0.067*2/3) mol.
Apply the ideal gas law to the pressure volume work (P V),
w have
P V = n R T
w = - D P V
= - D n R T
= - 0.10 mol*8.312 (J / (mol K))*298 K
= - 248 J
Discussion
The work done is due to the formation of gas O2 which
expands against the atmosphere of 1.0 atm or 101.3 kPa. The volume
changes of the solids are insignificant compared to that of the gas.
In case both pressure and volume change, and the work is the difference
of the pressure-volume product,
DP V.
Enthalpy
The enthalpy, usually represented by H is the energy released
in a chemical reaction under constant pressure, H = qP.
The enthalpy is a convenient property to evaluate for reactions taking
place at constant pressure. Enthalpy differ from internal energy, E,
defined in
Energy as the energy
input to a system at constant volume.
The energy released in a chemical reaction raises the internal energy,
E, and does work under constant pressure at the expense of
energy stored in compounds. Thus,
H = qP = E + P dV
Of course, the enthalpy change (dH) of a chemical reaction
depends on the amount of reactants, the temperature, and pressure.
Under normal conditions, the
ideal gas law
can be applied to give reasonable results.
Like the internal energy, enthalpy is also a thermodynamic state function,
depending only on the initial and final states of the system,
but not on the rate of reaction.
Standard Enthalpy of Reaction
In order to make the data useful for scientific and engineering applications,
there is a general agreement to report and tabulate enthalpy changes
for a mole of reaction at a standard temperature and pressure.
Such quantities are called the standard enthalpy of reaction.
In handbooks and textbooks, the standard enthalpy change is represented
by 
Ho.
For simplicity, we use dHo to represent the changes
of standard enthalpy in our discussion to avoid (very) slow loading of the
delta onto your computer.
Example 3
The standard enthalpy for the combustion of methane is 890.4 kJ per mole,
CH4(g) + 2 O2(g) -> CO2(g) + 2 H2O(g),
dHo = -890.4 kJ/mol
calculate the standard enthalpy change when 1.0 cubic meter of
natural gas is burned converting to gaseous products.
Solution
When 1.0 mol or 22.4 L of CH4, at 273K and 1 atm, is oxidized
completely, the standard enthalpy change is 890.4 kJ. One cubic meter
is 1000 L (/22.4 = 44.6 mol). Thus, the standard enthalpy of change is,
dH = 44.6 mol * 890.4 kJ/mol = 39712 kJ or 39 million joules.
A problem can be made up using any of the following standard enthalpy of
reactions. These are given here to illustrate the type of reactions
and the representation of enthalpy of reactions.
2 H(g) -> H2(g)
dHo = -436 kJ/mol
2 O(g) -> O2(g)
dHo = -498 kJ/mol
H2O(l) -> H2O(g)
dHo = 44 kJ/mol at 298 K
H2O(l) -> H2O(g)
dH = 41 kJ/mol at 373 K, non-standard condition
Mg(s) + S(s) -> MgS(s)
dHo = -598 kJ/mol
2 H(g) + O(g) -> H2O(g)
dHo = -847 kJ/mol
Cu(s) + 1/2O2(g) -> CuO(s)
dHo = -157 kJ/mol
1/2N2(g) + O2(g) -> NO2(g)
dHo = 34 kJ/mol
Mg(s) + 1/2O2(g) -> MgO(s)
dHo = -602 kJ/mol
2 P(s) + 3 Cl2(g) -> 2 PCl3(s)
dHo = -640 kJ/mol
2 P(s) + 5 Cl2(g) -> 2 PCl5(s)
dHo = -880 kJ/mol
C(graphite) + 2 O(g) -> CO2(g)
dHo = -643 kJ/mol
C(graphite) + O2(g) -> CO2(g)
dHo = -394 kJ/mol
C(graphite) + 2 H2(g) -> CH4(g)
dHo = -75 kJ/mol
2 Al(s) + Fe2O3(s) -> Al2O3(s) + 2Fe(s)
dHo = -850 kJ/mol
As we shall see, the application of
Hess Law will make these data very
useful. For example, applying Hess law using a few of these reactions
enable us to calculate the heat of combustion of methane to form
liquid water (as opposed to gaseous water) and carbon dioxide,
CH4 + 2 O2 -> 2 H2O(l) + CO2(g)
dH = -980 kJ/mol.
Enthalpy is an important topic in thermodynamics. Various methods have
been devised for the accurate measurement of heat of reaction under constant
pressure or under constant volume. This link gives a more advanced treatment
on enthalpy.
Standard Enthalpy of Formation, dHf
When the standard enthalpy is for a reaction that forms a compound from its
basic elements also at the standard state, the standard enthalpy of
reaction is called the standard enthalpy of formation, represented by
dHof. Unless specified, the temperature
is 298 K.
| Table of dHof
|
|---|
|
Compound | dHof
|
|---|
| MgS | -598 kJ/mol
|
| CuO | -157
|
| PCl3 | -320
|
| PCl5 | -440
|
| H2O | -286
|
| NO2 | + 34
|
| MgO | -602
|
| CO2 | -394
|
| CH4 | -75
|
In the above list, some of the equations lead to the formation of a compound
from its elements at their standard state. These equations and their
enthalpy of formation are given below:
Mg(s) + S(s) -> MgS(s)
dHof = -598 kJ/mol
P(s) + 3/2 Cl2(g) -> PCl3(g)
dHof = -320 kJ/mol
P(s) + 5/2 Cl2(g) -> PCl5(g)
dHof = - 440 kJ/mol
H2(g) + 1/2 O2(g) -> H2O(g)
dHof = -286 kJ/mol
1/2N2(g) + O2(g) -> NO2(g)
dHof = + 34 kJ/mol
Cu(s) + 1/2 O2(g) -> CuO(s)
dHof = -157 kJ/mol
Mg(s) + 1/2 O2(g) -> MgO(s)
dHof = -602 kJ/mol
C(graphite) + O2(g) -> CO2(g)
dHof = -394 kJ/mol
C(graphite) + 4 H2(g) -> CH4(g)
dHof = -75 kJ/mol
In all the above equations of reaction, the right hand side has only
one product and that its coefficient is 1. A general rule is to consider
standard enthalpy of formation of all elements at the standard condition
to be zero. Then, there is no need to write the complete equation
for the tabulation of the standard enthalpy of formation.
The above list can be simplified to give the table shown here.
A simple application of the standard enthalpy of formation is illustrated
by Example 4.
Example 4
For NH3, dHf = -46.1 kJ/mol. Estimate energy
released when 10 g of N2 react with excess of H2 to form
ammonia.
Solution
Ten grams of nitrogen is less than 1 mol, and we carry out the calculation
in the following manner:
1 mol N2 - 46.1 kJ
10 g N2 ---------- ---------- = - 32.9 kJ
28.1 g N2 0.5 mol N2
Thus, 32.9 kJ is released when 10 g of N2 is consumed.
Standard enthalpies of formation and standard entropies are important
thermodynamic data,
and this link gives an extensive table of values for some key compounds.
The principle of conservation of energy leads to the formulation of the
Hess law.
It's application makes the enthalpy of reaction and standard enthalpy
of formation very useful.
Confidence Building Questions
cchieh@sciborg.uwaterloo.ca