The Second Law of Thermodynamics
The Second Law of Thermodynamics
expresses the notion that there is an inherent direction in which any system
moves if it is not at equilibrium. For
instance, when a shiny nail is left outdoors, it rusts. This process occurs without outside
intervention and is said to be spontaneous.
For every spontaneous process, there is a reverse process. We can imagine that a rusty iron nail is
transformed into a shiny one, but it is inconceivable that this process is
spontaneous.
Our years of observing
nature have impressed upon us the simple rule: Processes that are spontaneous in one direction are not spontaneous in
the reverse direction.
The fact that a process is spontaneous does not mean that it will occur at an observable rate. A spontaneous reaction may be fast or slow. Thermodynamics can tells us the direction and extent of a reaction, but not the speed. Reaction rates are the subject of chemical kinetics.
The fact that a process is spontaneous does not mean that it will occur at an observable rate. A spontaneous reaction may be fast or slow. Thermodynamics can tells us the direction and extent of a reaction, but not the speed. Reaction rates are the subject of chemical kinetics.
Spontaneity and Entropy
The spontaneous motion of a brick released from your hand is toward
the ground. As the brick falls, it loses
potential energy. This potential energy
is first converted into kinetic energy, the energy of the motion of the
brick. When the brick hits the floor,
its kinetic energy is converted into heat.
The overall result of the brick’s fall is thus a conversion of the
potential energy of the brick into heat in its surroundings. This phenomenon is result of a system’s tendency to seek a resting place of
minimum energy.
The First Law of Thermodynamics (aka the Law of Conservation of Energy) states that energy is neither created nor destroyed in any process, only converted from one form into another.
The First Law of Thermodynamics (aka the Law of Conservation of Energy) states that energy is neither created nor destroyed in any process, only converted from one form into another.
The
tendency for a system to achieve the lowest possible energy is one of the
driving forces that determines the behaviour of molecular systems. The brick possesses potential energy because
of its position relative to the floor.
Likewise, a chemical substance possesses potential energy relative to
other substances because of the arrangements of nuclei and electrons. When these arrangements change, energy may be
released.
For example, the combustion of
propane, which is clearly a spontaneous process, is very exothermic:
C3H8(g) + 5O2(g) ↔
3CO2(g) + 4H2O(l) ΔH=-2202 kJ
The
rearrangements of nuclei and electrons in going from reactants to products lead
to a lower chemical potential energy and so heat is evolved. Reactions that are highly exothermic are
generally spontaneous. However, the
tendency toward minimum enthalpy is not the only factor that determines
spontaneity in molecular processes.
Spontaneity & Entropy Change
There
are several processes that are spontaneous even though they are not
exothermic. For example, consider the
melting of ice at room temperature. The
process
H2O(s) →
H2O(l)
is very spontaneous at 22°C,
yet it is an endothermic change. The
molecules of water that make up an ice crystal are held rigidly in place in the
ice crystal lattice. When the ice melts,
the water molecules are free to move about.
Thus, in liquid water the individual water molecules are more randomly
distributed than in the solid. A highly
ordered solid structure is replaced by the much more disordered liquid
structure.
As
the above example illustrates, spontaneity
is associated with an increase in randomness or disorder of a system. The randomness is expressed by a
thermodynamic quantity called entropy,
given the symbol S. The more disordered
a system, the greater its entropy.
Like
enthalpy, entropy is a state function and the change of entropy of a system, ΔS=Sfinal-Sinitial,
depends only on initial and final states.
A positive ΔS indicates an increase in randomness
or disorder and signals a spontaneous process.
A negative ΔS indicates a decrease in randomness and
indicates a non-spontaneous process.
Molecular Interpretation of Entropy & The Third Law of Thermodynamics
A molecule can engage in three types of
movement, translational, vibrational and rotational motion. These forms of motion are ways that the molecule
has of storing energy. As the
temperature of a system increases, the amounts of energy stored in these forms
of motion increase.
To see how this pertains to entropy,
imagine a pure substance that forms a perfect crystalline lattice at the lowest
possible temperature, absolute zero. The Third Law of Thermodynamics states that
the entropy of a pure crystalline substance at absolute zero is zero, S(0K)=0. As the substance is heated the molecules have
increased freedom of motion and this adds to the entropy content of the
substance. Therefore, as the temperature
of a system rises, so does the entropy of the system.
Calculation of Entropy Changes
The entropy values of substances in
their standard states (1 atm) are known as standard entropies and are denoted S°. The entropy change in a chemical reaction is
given by the sum of the entropies of the products minus the sum of the
entropies of the reactants. If this equation looks familiar, it should. We have seen a similar equation in the past. Use this new equation just like in the past, except we use S values instead of ΔH values.
ΔS°=∑S°(products)
- ∑S°(reactants)
Predicting Placement of the Heat Term
The
entropy of a system can be used to determine the placement of the heat term in
the chemical reaction equation. The heat
term is always placed in the equation, opposite the side of maximum
disorder. For instance, consider the
Haber process.
N2(g) + 3H2(g) ↔
2NH3(g)
Maximum disorder occurs on the reactant
side of the equation because four moles of gas are more disordered than two
moles of gas. Since the reactant side of
the equation is more disordered, these molecules contain more stored
energy. When the system shifts towards
products, the stored energy of the system drops and the excess energy is
released as heat. Thus the heat term is
placed on the product side of the reaction, indicating that this process is
exothermic (ΔH<0):
N2(g) + 3H2(g) ↔
2NH3(g) + heat
****************************************************
To determine which side has the greatest disorder, first compare the number of moles of gas on each side (since gases have the greatest disorder). The side with the higher amount of gas, will be more disordered and so the heat term will be placed opposite to this. If there are no gases, instead compare the amounts of aqueous solutions on each side. No solutions? Look for liquids and if there are no liquids, look at the amount of solids.
Tryits!:
Place the heat term on the proper side of each reaction equation.
H2O(l) ↔ H2O(s) + heat
heat + CaCO3(s) ↔
CaO(s) + CO2(g)
*highlight the line for the above rxns with your cursor to reveal the answers
Homework #31-35
