Lecture 8
Stability and reactivity
We tend to say that substances are ‘stable’ or •
‘unstable’, ‘reactive’ or ‘unreactive’ but these
terms are relative and may depend on many
factors.
Thermodynamic and kinetic factors can also •
be important.
Stability and reactivity can be controlled by
thermodynamic factors (depending only on
the initial and final states and not on the
reaction pathway) or kinetic ones (very
dependent on the reaction pathway).
Both factors depend on the conditions, and on
the possibility of different routes to
decomposition or reaction.
Enthalpy and Hess’ Law
The enthalpy change (ΔH) in a reaction is equal
to the heat input under conditions of constant
temperature and pressure.
Enthalpy is commonly used as a measure of the
energies involved in chemical reactions.
Endothermic reactions (positive ΔH) are ones
requiring a heat input, and exothermic
reactions (negative ΔH) give a heat output.
Hess’ Law states that ΔH does not depend on the
reaction pathway( taken between initial and
final states), and is a consequence of the First
Law of Thermodynamics.
ΔH can be expressed as the sum of the values for
individual steps:
ΔH = ΔH1+ ΔH2 +……..
Enthalpy change does depend on conditions of
temperature, pressure and concentration of
the initial and final states, and it is important
to specify these. Standard states are defined
as pure substances at standard pressure
(1bar), and temperature(298K).
Schematic thermodynamic cycle illustrating the use of
Hess’ Law
The standard enthalpy of formation of any
compound
refers to formation from its elements, all in
standard states. By definition,it is zero for any
element in its stable(standard ) state.
enthalpy change ΔHΘ in any reaction to be
calculated from
Entropy and free energy
- Entropy (S) is a measure of molecular
‘disorder’, or more precisely ‘the number of
microscopic arrangements of energy possible
in a macroscopic sample’.
- Entropy increases with rise in temperature and
depends strongly on the state
- Entropy changes (ΔS) are invariably positive for
reactions that generate gas molecules.
The Second Law of Thermodynamics asserts
that the total entropy always increases in a
spontaneous process, and reaches a maximum
value at equilibrium.
Both internal and external changes are taken
account of by defining the Gibbs free energy
change (ΔG):
for a reaction taking place at constant
temperature (T, in kelvin)
Gibbs free energy change (ΔG)
Gibbs free energy change (ΔG)
From the Second Law it can be shown that ΔG is
always negative for a feasible reaction at
constant temperature and pressure and is
zero at equilibrium.
ΔS and ΔG for reactions do not depend on the
reaction pathway. They depend even more
strongly than ΔH on concentration and
pressure.
Tabulated standard entropies may be used to estimate
changes in a reaction from where SΘ values are not zero for
elements
Equilibrium constants
For a general reaction such as
aA +bB
cC + dD
the equilibrium constant is
K =[A]a[B]b/[C]c[D]d
where the terms [A], [B] as concentrations or
partial pressures. (This assumes ideal
thermodynamic behavior and is a much better
approximation for gases than in solution.)
A very large value (≫1) of K indicates a strong
thermodynamic tendency to react, so that
very little of the reactants (A and B) will
remain at equilibrium. Conversely, a very small
value (≪1) indicates very little tendency
to react: in this case the reverse reaction (C and D
going to A and B) will be very favorable.
For any reaction K may be related to the standard
Gibbs free energy change (ΔGΘ) according to
(ΔGΘ) = - R T lnK
where R is the gas constant (=8.314 J K−1 mol−1)
and T the absolute temperature (in K). Thus
equilibrium constants can be estimated from
tabulated values of and trends may often be
interpreted in terms of changes in ΔHΘ and ΔSΘ.
Important notes
- Equilibrium constants change with
temperature in a way that depends on ΔHΘ
for the reaction
- In accordance with Le Chatelier’s principle, K
increases with rise in temperature for an
endothermic reaction, and decreases for an
exothermic one.
Reaction rates
- The rate of reaction generally depends on the
concentration of reactants.
Rate =k[A]n[B]m
- where k is the rate constant and n and m are the orders
of reaction with respect to reactants A and B.
- Orders of reaction depend on the mechanism and are
not necessarily equal to the stoichiometric
coefficients a and b.
The rate constant depends on the mechanism and
especially on the energy barrier or activation energy
associated with the reaction pathway.
- High activation energies (Ea) give low rate
constants because only a small fraction of
molecules have sufficient energy to react.
- This proportion may be increased by raising
the temperature, and rate constants
approximately follow the Arrhenius equation.
Arrhenius equation
K =Ae-Ea/RT
- Large activation energies arise in reactions
where covalent bonds must be broken before
new ones are formed, or where atoms must
move through solids.
- Reactions involving free radicals, or ions in
solution, often have small (sometimes zero)
activation energies.
Catalyst
Reactions may be accelerated by the presence of a catalyst, which acts by providing an
alternative pathway with lower activation
energy.
- A true catalyst by definition can be recovered
unchanged after the reaction, and so does not
alter the thermodynamics or the position of
equilibrium