Thermodynamics
·
A study of energy
changes in chemical and physical processes.
·
Energy is the
capacity to do work or to transfer heat
Energy can be stored in
molecules
·
Reactions involve a
rearrangement of bonds, and consequently may result in a net intake or output of
energy.
·
Reactions occur
when reactant molecules collide.
This collision must involve "enough" energy and must involve the correct
orientation.
·
The amount of
energy required for a successful collision is referred to as the activation
energy. All reactions require
an input of some level of activation energy.
·
Some reactions are
exothermic -- the overall process of going from reactants to products
results in a net release of energy as heat
·
Endothermic
reactions -- result in a net intake of energy
·
The energy taken in
or given off during a reaction is referred to as the change in enthalpy
(DH). For an
exothermic reaction DH < 0; for an endothermic reaction DH > 0.
·
All of this
information can be summarized in a potential energy
diagram:
The potential energy diagram
shown above applies to the rxn CH4 + 2 O2 à CO2 + 2H2O
·
Do the products
have more energy or less energy than the reactants?
·
Would this relate
to an endothermic or exothermic process?
·
Would ∆H be greater
than or less than zero?
·
What is the source
of the activation energy when we burn methane?
Physical processes also involve
energy changes (ex. Melting, condensing, dissolving,
etc.)
First Law of Thermodynamics:
the total amount of energy in the universe is constant
·
Corresponds to Law
of Conservation of Energy: energy is neither created nor destroyed in ordinary
chemical and physical changes
Definitions:
·
System: the
substances involved in a chemical or physical change
·
Surroundings:
everything outside the system
·
Universe: the
system and its surroundings.
·
Thermodynamic state
of a system: the conditions that specify all the properties of the system (P, V,
T, # of moles, physical state, etc.)
·
Changes in state
functions = final state – initial state; changes are path
independent.
Ex. A
temperature change from -5° C to +5° C is the same as a temperature change from
-5° C to +15° C to +5° C. Both
DT = +10° C.
·
Enthalpy change
(DH) – the quantity of heat transferred into or out of a
system as it undergoes a chemical or physical change at constant
pressure
While it is impossible to measure the absolute enthalpy
of a system, enthalpy change can be measured. (Ex.
Calorimetry)
Using a calorimeter to measure enthalpy
change
·
A calorimeter is an
insulated container which is part of the system in which a reaction
occurs.
·
The heat capacity
of the calorimeter must be determined.
Ex.: After adding 100.0 g of
water at 58.5°C to 100.0 g of water at 22.8°C in a calorimeter, the final temperature of the water
is 39.7°C.
·
Heat gained by cool
water:
·
Heat lost by warm
water:
·
Heat gained by
calorimeter:
·
Heat capacity of
calorimeter:
Once the heat capacity of the
calorimeter is determined, enthalpy changes from rxns can be
measured.
Ex. 20.00 mL of 0.625 M NaOH is
reacted with 30.00 mL of 0.500 M CH3COOH in a calorimeter. Both are initially at 21.40°C; after the reaction the temperature is
24.35°C and the density of the solution is 1.02
g/mL.
·
Heat released by
rxn = heat absorbed by soln + heat absorbed by calorimeter
Molar heat of neutralization =
amount of heat released/ 1 mole of water formed.
·
Find limiting
reactant and # of moles of H2O formed:
·
Calculate molar
heat of neutralization:
The thermochemical equation for
this reaction would be written as follows:
NaOH
(aq) + CH3COOH (aq) à
NaCH3COO (aq) + H2O (l) DH
=
·
In a thermochemical
equation, DH always refers to the number of moles of substances
shown and the states of matter shown.
·
If this reaction
was run using sufficient quantities of acid and base to produce two moles of
water, the enthalpy change observed would be
.
Use the following data to
calculate DH for the thermochemical equation.
·
The following
reaction is used in the attitude-controlled engines of the space
shuttles:
CH6N2 (l) + 5/4
N2O4 (l) à
CO2 (g) + 9/4 N2 (g) +
3H2O (l)
The two reactants ignited instantly on contact,
producing a flame temperature of 3000 K.
The energy liberated per 0.100 g of CH6N2 at
constant atmospheric pressure is 750 J.
·
This amount is
referred to as the enthalpy change/mol of rxn.
·
How many kJ are
liberated when 44.0 g of N2 is produced?
Standard States
·
Thermodynamic
standard state: the most stable state for a pure substance at 1 atm and
25°C.
·
For a gas in a
mixture, its partial pressure must be 1 atm in its standard
state
·
For an aqueous
solution, the standard state refers to a 1 M concentration
Standard enthalpy changes (DH°rxn)
·
Reactants and
products are specified as being in their standard states; all reactants are
converted to products.
·
The reaction is
often said to occur at constant T and P.
Changes in T and P may occur during the reaction, but then T and P of
products are returned to the original conditions.
Standard molar enthalpy of formation (DH°f)
·
Enthalpy change
when 1 mole of the substance is formed from its elements at standard
states.
·
DH°f value of
any element in its standard state is defined as being equal to zero. (See
Appendix 4)
·
DH°f for NaCl
(s) = -411.0 kJ/mol. Write the
thermochemical equation which corresponds to this
information.
Na (s) + ½ Cl2
(g) à NaCl (s) DH°f= -411.0
kJ
·
Is this an
exothermic or endothermic process?
Hess’s Law
·
The enthalpy change
for a reaction is the same whether it occurs by one step or by any series of
steps.
Use the following measured
enthalpies of reaction to determine DHrxn for
Ca2+ (aq) + 2
CaCO3 (s) à CaO (s) + CO2 (g)
DH = +178.1
kJ/mol rxn
CaO (s) + +
H2O (l) à Ca(OH)2 (s)
DH = -65.3
kJ/mol rxn
Ca(OH)2 (s) à
Ca2+ (aq) + 2
See problems 59 & 61, p
269
Hess’s Law also can be used
with DH°f values to
calculate the enthalpy change for a reaction.
DH°rxn =
å n DH°f (products)
- å n DH°f (reactants)
n = coefficients from the reaction
Ex. Calculate the standard
enthalpy change for the following reaction.
SF6 (g) + 3
H2O (l) à 6 HF (g) + SO3
(g)
Bond energies (See Table 8.4,
p. 351)
Energy is required to break
any chemical bond.
·
Bond energy: the
amount of energy needed to break one mole of bonds in a gaseous covalent
substance, forming gaseous products at constant T & P
·
Bond energies vary
based on the atoms involved and their surrounding atoms.
·
Larger values for
bond energy signify stronger (more stable) bonds.
·
Triple bonds are
stronger than double bonds; double bonds are stronger than single
bonds.
·
Bond energies can
be used to estimate enthalpy changes for gaseous state
reactions.
DH°rxn =
å B.E. (reactants) - å B.E. (products)
·
Energy must be put
in to break bonds of reactants; energy is given off as bonds of products form
·
Use bond energies
to estimate the enthalpy of reaction for the following gas phase
reaction.
CO + H2O à CO2 + H2
See problem 54a, p.
384
Internal energy refers to all
of the energy contained within a substance (kinetic energy, attractive and
repulsive forces, etc.)
·
Changes in internal
energy
DE = E final – E initial = E
products – E reactants = q + w
where q = heat and w = work
·
DE = heat absorbed by system + work done on
system
·
If heat is absorbed
by the system, q > 0; if heat is released by the system, q <
0
·
If work is done on
the system, w > 0; if work is done by the system, w <
0
·
Work usually
involves changes in P and V.
·
When pressure is
applied to a system from its surroundings, work is done on the system to
compress it.
·
When pressure is
applied by a system on its surroundings, work is done by the system as it
expands.
·
For reactions that
yield a net increase of gas molecules, work is done by the system: w < 0. (2
NH4NO3 (s) à 2 N2 (g) + 4 H2O
(g) + O2 (g))
·
For reactions that
yield a net decrease of gas molecules, work is done on the system: w> 0. (2 SO2 (g) +
O2 (g) à 2 SO3 (g))
·
If a system is held
in a constant volume sealed container (bomb calorimeter), no work is done and
DE = q.