HEAT CHANGES AND THERMOCHEMISTRY:
Energy is the capacity to do
work or to transfer heat.
Remember, that HEAT is
transferred, NOT cold !
2 general kinds of energy: KINETIC & POTENTIAL.
KINETIC
= energy of motion. Ex.: car moving at 60 mph.
Ekinetic = 1/2 mv2, where m = mass of car
and v = velocity
POTENTIAL
= energy due to position or composition.
Ex.: water held behind a dam. When we release the
water, the potential energy is converted into kinetic
energy as the water runs downhill.
Chemical energy comes from
the POTENTIAL ENERGY
stored in
the atoms and molecules due to their
arrangements.
When chemicals react and
release their potential energy
in the form
of heat, then these reactions are
EXOTHERMIC.
Ex.: C8H18(l) + O2(g) ----> CO2(g)
+ H2O(l) + J (heat)
(not balanced!!!)
C9H20(l) + O2(g) ----> CO2(g) + H2O(l) + J (heat)
In both of these reactions,
the total amount of energy in
the products
is less than the total amount of energy in
the
reactants (therefore, heat is RELEASED in the
reaction =
EXOTHERMIC).
**Show Potential Energy
diagram vs. Progress of Reaction
for
both EXOthermic and ENDOthermic reactions.
Also, when there are phase
changes, energy is either
released or
required - to allow the phase change to occur.
Ex.:
When you boil water to produce steam, you GIVE
energy to
the water; you heat it. Vice-versa, when
steam
condenses to form water, it GIVES off energy
as it
condenses. Or, you take heat away from water
to form ice.
Examples: 1) ammonium
nitrate dissolving in water:
spontaneous and ENDOthermic!
2) water(s) = ice ===> water(l) is
spontaneous and ENDOthermic!
3) acid + base ==> salt + water:
EXOthermic and spontaneous!
BOTTOM LINE:: DH is NOT
the predictor of spontaneity
of a reaction.
There are 3 laws of thermodynamics:
1- The total
amount of energy in the universe is constant.
OR: Energy is neither created nor
destroyed in ordinary
chemical reactions and physical changes, but only
converted from one form to another (Law of
Conservation of Energy). REMEMBER DHo
2- In spontaneous changes, the entropy of the universe
increases, and in an equilibrium
process, the entropy
of the universe remains unchanged.
Entropy is a
measure of the randomness or the disorder of a
system. Particles in
the solid state are MORE ordered
than particles in the liquid state which are
more
ordered than particles in the vapor state
(EX.: water!!).
Standard entropy = DSo is the entropy of substance
at 1 atm and 25oC. Elements do have DSo, in units
of J/K or
J/K-mol.
Positive value of DSo means MORE disorder, and
vice-versa!!
3- The entropy of a hypothetical pure, perfect, crystalline
substance
at absolute zero (0 Kelvin) is zero.
DEFINITIONS:
SYSTEM: all the substances involved in the chemical &
physical changes that we are studying. DSsys
SURROUNDINGS:
everything in the system's
environment. DSsurr
UNIVERSE: the SYSTEM plus its SURROUNDINGS.
DSuniv
DSuniv = DSsys + DSsurr
Changes
are described as: D
D
(ANYTHING) = ANYTHINGfinal
- ANYTHINGinitial
ENTHALPY CHANGES: The
quantity of heat transferred
into or out of a system when it
undergoes a chemical or
physical change at constant
pressure, qp, is called the
enthalpy
change, DH, of the process,
OR,
"heat of reaction".
DH = Hfinal -
Hinitial
- - - - - - - - - - - - - - - - -
- - - - - - - - - - - - - - - - - - - - - - - -
We know how to find DH for a reaction.
When
EXOthermic process occurs in a system, heat is
transferred to the surroundings, with a concurrent
increase in the disorder of the surroundings at the
molecular level, and the entropy of surroundings
increases.
Vice-versa,
ENDOthermic process, heat is absorbed BY
the system from the surr., and the entropy of the surr.
decreases.
Therefore,
DSsurr = -DHsys
where
+DS means more entropy and -DH means
release of heat.
Since
it is more difficult to release heat from a system
to the surr. at higher T, then we want to rewrite the
equation above to be:
DSsurr = -DHsys/T
----------------------------------------------------------------
Look at N2(g)
+ 3H2(g) ----> 2NH3(g) DHo = -92.6
kJ
Calc DSo = 2DSo(NH3) - [DSo(N2) + 3DSo(H2)]
= (2
mol)(193J/K-mol) -
[(1mol)(192J/K-mol)+(3mol)(131J/J-mol)]
= -199 J/K
Therefore, this is DSsys
DSsurr = -DHsys/T
= -(-92,600 J)/298K = 311 J/K
Therefore: DSuniv = DSsys + DSsurr
DSuniv = -199 J/K + 311 J/K = 112 J/K
We
predict that the reaction is spontaneous at 25oC.
---------------------------------------------------------------
How do we find DSo for a reaction?
DSoreaction
= Sum[nDSo(prods)] - Sum[mDSo(reacts)]
If a
reaction produces more gas molecules than were as
reactants, DSo is positive =
more disorder; and vice-versa
If no
net change in # molecules of gas, then DSo may
be + or -, but will be small numerically.
-----------------------------------------------------------------
STANDARD STATES & STANDARD ENTHALPY CHANGES:
a) For
a pure liquid or solid, the standard state IS the pure
liquid or solid.
b) For a
gas, the standard state IS the gas
at 1 atm.; in gas
mixtures, its partial pressure must be 1 atm.
c) For
solutions, the standard state is 1
M concentration.
------------------------------------------------------------------
Spontaneous Processes:
Occur without outside
intervention, e.g., drop two eggs...
they break!
The reverse
process is non-spontaneous.
The direction of a spontaneous
process can depend on
temperature.
Ice
turning to water is spontaneous at T > 0°C.
Water
turning to ice is spontaneous at T < 0°C.
Allowing 1 mol of ice to warm up is an irreversible process.
To get the reverse process to occur, the water temperature
must be lowered to 0°C (heat must be removed).
In any spontaneous process, the
path between reactants
and products
is irreversible, unless outside intervention
occurs.
E.g., water can be
frozen...then thawed by adding heat.
Thermodynamics gives us the direction of a process, but
not the speed
at which it will occur (this latter is
KINETICS).
How can endothermic reactions be spontaneous?
Entropy
and the Second Law of Thermodynamics:
Consider an initial state: two
flasks connected by a closed
stopcock. One flask is evacuated and the other contains
1 atm. of a
gas. Then open the stopcock....The final
state: two
flasks connected by an open stopcock.
Each
flask contains the gas at 0.5 atm. The expansion
of the gas
is isothermal (i.e., constant temperature).
Therefore,
the gas does NO work and heat is not
transferred.
Why does a gas expand?
ENTROPY!
S: is a measure of the disorder of a system.
The bigger
the disorder.....bigger entropy.
More order.....lower entropy (ice is more ordered
than water, which is more ordered than steam).
Spontaneous
reactions proceed to lower energy or
higher
entropy.
In the
connected flasks, is it more likely that all the gas
molecules
stay in one flask or that they spread out
randomly
to fill both flasks? The molecules are more
likely to
fill both flasks, and the system has moved to a
state of
higher entropy.
What about thawing ice?
In ice, the
molecules are very well ordered because of
the H-bonds
throughout the system (and the molecules
of water are
in fixed positions); therefore, ice has a low
entropy. As ice melts, the intermolecular forces are
broken
(which requires energy), but the order is
interrupted,
so entropy increases. Water is more random
than ice, so
ice spontaneously melts at room
temperature. There is always a balance between energy
and entropy
considerations.
Generally, an increase in entropy in one process is
associated
with a decrease in entropy in another; the
increase in
entropy usually dominates.
Entropy is a state
function: A state
function is one in which
the value depends only on the state of the system, and NOT
how they got there. Ex.:
P, V, T, DH, DS.
For a system, DS = Sfinal - Sinitial
If DS > 0, the randomness increases (less order);
if DS < 0, the order
increases.