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.
E_kinetic =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 - must balance it!!!)

        C₉H₂₀(l) + O₂(g) ---->  CO₂(g)  + H₂O(l) +  J (heat)
           (not balanced - must balance it!!!)

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.

  BOTTOM LINE::  DH is NOT the predictor of spontaneity
                            of a reaction.

18.2 SPONTANEOUS PROCESSES AND ENTROPY

18.3 THE SECOND LAW OF THERMODYNAMICS

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 = DS^o is the entropy of substance
         at 1 atm and 25^oC.  Elements do have DS^o, in units
         of J/K or J/K-mol.

       Positive value of DS^o means MORE disorder, and
         vice-versa!!

The Third Law of Thermodynamics and Absolute Entropy

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:  DS_sys

  SURROUNDINGS:  everything in the system's environment:
                  DS_surr

Changes are described as:  D
            D  (ANYTHING) = ANYTHING_final - ANYTHING_initial

ENTHALPY CHANGES:  The quantity of heat transferred into
or out of a system when it undergoes a chemical or physical
change at constant pressure, q_p, is called the enthalpy
change, DH, of the process, OR,  "heat of reaction".

                        DH = H_final - H_initial
- - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - -
We know how to find  DH for a reaction. 

When EXOthermic process occurs in a system, heat is
transferred to the surroundings vy the system, with a
concurrent increase in the disorder of the surroundings
at the molecular level, and the entropy of the
surroundings increases.  

Vice-versa, in an ENDOthermic process, heat is absorbed
BY the system from the surroundings, and the entropy
of the surroundings decreases.

Therefore,  DS_surr  =  -DH_sys 

where +DS means more entropy and -DH means release
of heat.

Since it is more difficult to release heat from a system
to the surroundings at higher T, then we want to rewrite
the equation above to be:

            DS^o = -199 J/K            Therefore, this is DS_sys

Calculate DS_surr:
    DS_surr  =  -DH_sys/T  =  -(-92,600 J)/298K = 311 J/K

Therefore:      DS_univ  = DS_sys  +  DS_surr

                            DS_univ  = -199 J/K + 311 J/K = +112 J/K

How do we find  DS^o for a reaction?

 DS^o_reaction = Sum[nDS^o(prods)] - Sum[mDS^o(reacts)]

If a reaction produces more gas molecules than there were
gas reactants, DS^o is positive = more disorder; and
vice-versa = more gas reactants than products, less
disorder and DS^o is negative.

STANDARD STATES & STANDARD ENTHALPY CHANGES:

Spontaneous Processes:
Occur without outside intervention, e.g., drop two eggs..
        they break!
    BUT, 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.....the bigger the 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 = S_final - S_initial
If DS > 0, the randomness increases (less order).
If  DS < 0, the order increases.

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