Chapter 5.6: Back to Phase Changes |
Now let us return to the situation with solids, liquids, and gases. How do we think about entropy in these systems? Doesn’t a substance become more ordered as we move it from gas to liquid to solid? Clearly the entropy of a solid is lower than that of a liquid (and the entropy of a liquid is lower than that of a gas). We can calculate how entropies change for materials as they go from gas to liquid to solid, and as we have predicted they decrease. How can it be that a change in which the entropy of the system decreases (for example freezing of ice) can occur? Are we forced to conclude that things we know to happen are impossible according to the Second Law of Thermodynamics? |
5.1 Systems |
Free Energy to the Rescue This idea, that we must consider changes in entropy for both the system and its surroundings when we are predicting which way a change will occur (that is, in which direction a process is thermodynamically favorable) is an important one. However, it is almost always easier to look at the system than it is to look at the surroundings (after all we define the system as that part of the universe we are studying.) It would therefor be much more convenient to use criteria for change that refer only to the system. Fortunately there is a reasonably simple way to do this. Let us think about water freezing again: we can measure the enthalpy change for this process. The thermal energy change for the system, ΔHfreezing, is about -6 kJ/mole. That is 6 kJ of thermal energy are released into the surroundings for every mole of water that freezes. We can relate this thermal energy release to the entropy change of the surroundings. Entropy is measured in units of J/K – that is energy/temperature. Since we know how much energy is added to the surroundings we can calculate the entropy change that this energy produces. Mathematically we can express this as ΔSsurroundings = ΔHsurroundings/T. And since we know that ΔHsystem = -ΔHsurroundings, (that is: the energy lost by the system = minus (-) the energy gained by the surroundings) we can express the entropy change of the surroundings in terms of measurable variables for the system. That is ΔSsurroundings = -ΔHsystem/T. The quantity -TΔStotal has units of energy, and is commonly known as the free energy change ΔG (or the Gibbs Free energy). The equation is normally written as: ΔG = ΔH – TΔS. The free energy change of a reaction is probably the most important thermodynamic term that you will need to learn about. In most biological and biochemical systems, it is ΔG that is commonly used to determine whether reactions are thermodynamically favorable. However, it is important to remember that ΔG is a proxy for the entropy change of the universe; if it is negative, universal entropy is increasing (and the reaction occurs), if it is positive, universal entropy would decrease if the reaction occurred (and so it does not). It is possible however for reactions with a positive ΔG to occur, but only if they are coupled with a reaction with an even greater negative ΔG (see chapters 8 and 9). What is “Free” about Gibbs Free Energy? The reason we use ΔG or ΔGº to describe many systems (and especially biological ones), is that both the magnitude and sign tell us a lot about how that system will behave. We use ΔG (the Gibbs free energy change) rather than ΔH (the enthalpy change) because ΔG tells us how much energy is actually available to bring about further change (or to do work). It factors out for us the energy produced from the change from what that is lost to the surroundings as increased entropy (and so is not available for us to use to do work.) As an example, when wood is burnt, it is theoretically impossible for all of the heat released to be used to do work, some of the energy goes to increase the entropy of the system. For any change in the system, some of the energy is always lost in this way to the surroundings; this is why it is impossible to build a machine that is 100% efficient in converting energy from one kind to another (although many have tried - google perpetual motion machines.) The term “free energy” then doesn’t mean that it is literally free - but rather means what is theoretically available for further transformations. When ΔG is negative, we know that the reaction will be thermodynamically favored.[Many people use the term spontaneous - but this is misleading since it infers that the reaction will happen right away. In fact ΔG tells us nothing about when the process will happen - only that it is favored. As we will see later the rate at which a process occurs is governed by other factors. ] The best case scenario is when ΔH is negative (an exothermic change - that is the system is losing energy to the surroundings and becoming more stable), and ΔS is positive - (that is the system is increasing in entropy). Because T is always greater than 0 (in degrees K), TΔS will be positive and this value will be subtracted from ΔH to given an even larger negative ΔG value. A good example of such a process is the reaction (combustion) of sugar (C6H12O6) with molecular oxygen (O2): This is an exothermic process and, as you can see from the reaction equation, it results in the production of more molecules than we started with (often a sign that entropy has increased.) Free energy, and temperature So we have two very clear cut cases that allow us to predict whether a process will occur - where the enthalpy and entropy predict the same outcome. But there are also two other situations where the enthalpy and entropy “point” in different directions. When this happens we need to use the fact that the free energy change is temperature dependent to predict the outcome. Recall the expression ΔG = ΔH – TΔS depends upon temperature. For a system where the entropy change is positive (+ΔS), an increase in temperature will lead to an increasingly negative contribution to ΔG. That is, as the temperature rises - a process that involves an increase in entropy will become more favorable. Conversely, if the system change involves a decrease in entropy, (that is ΔS is negative) then ΔG will become more positive (and less favorable) as the temperature increases. |
ΔH | ΔS | ΔG |
Negative (exothermic) | Positive (entropy increases) | Negative at all temperatures (always thermodynamically favored) |
Positive (endothermic) | Negative (entropy decreases) | Positive at all temperature (never thermodynamically favored)` |
Negative (exothermic) | Negative (entropy decreases) | Depends on the temperature, as the temperature increases ΔG will become more positive and the reaction will become less favored (go backwards) |
Positive (endothermic) | Positive (entropy increases) | Depends on the temperature, as the temperature increases ΔG will become more negative and the reaction will become favored (go forwards) |
Up to this point, we have been considering physical changes to a system, that is populations of molecules going from solid to liquid or liquid to gaseous states (and vice versa). Not really what one commonly thinks of as chemistry, but the fact is that these transformations involve the making and breaking of interactions between molecules. We can therefore consider phase transitions as analogous to chemical reactions, and because they are somewhat simpler, develop logic that applies to both processes. So let us begin by considering a phase change/reaction system H2O (liquid) ↔ H2O (gas). We use a double arrow ↔ to indicate that, depending upon the conditions, the reaction could go either to the right (boiling) or to the left (condensing). So, let us assume for the moment that we do not already know that water boils (changes from liquid to gas) at 100ºC, what factors would determine whether the reaction H2O (liquid) ↔ H2O (gas) favors the liquid or the gaseous state at a particular temperature? As we have seen the criterion for whether a process will “go” at a particular temperature is ΔG. We also know that the free energy change for a reaction going in one direction is the negative of the ΔG for the reaction going in the opposite direction. So that, the ΔG for the reaction: H2O (liquid) ↔ H2O (gas) is –ΔG for the reaction H2O (gas) ↔ H2O (liquid). When water boils all the intermolecular attractions between the
water molecules must be overcome, allowing the water molecules
to fly off into the gaseous phase. Therefore the process
of water boiling is endothermic (ΔHvaporization = +40.65
kJ/mol); it requires an energy input from the surroundings
(when you put a pot of water on the stove you have to turn
on the burner for it to boil). When the water boils, the
entropy change is quite large (ΔSvaporization = 109
J/mol K), as the molecules go from being relatively constrained
in the liquid to gas molecules that can fly around. At temperatures
lower than the boiling point the enthalpy term predominates
and ΔG is positive, but as you increase the temperature
in your pan of water, eventually it reaches a point where
the contributions to ΔG of ΔH and TΔS are
equal. That is ΔG goes from being positive to negative
and the process becomes favorable. At the temperature where
this “cross over” occurs ΔG = 0 and ΔH
= TΔS. At this temperature (373 K, 100 ºC) water
boils (at 1 atmosphere). At temperatures above the boiling
point ΔG is always negative and water will exist predominantly
in the gas phase. However if we let the temperature drop
below the boiling point, the enthalpy term becomes predominant
again and ΔG for boiling is positive - that is: water
will not boil at temperatures below 100 ºC (at one atmosphere). |
Water freezes at temperatures below 0 ºC. At temperatures where phase changes take place, (boiling point, melting point) ΔG = 0. Furthermore, if the temperature were kept constant - there would be no observable change in the system. We say that the system is at equilibrium; for any system at equilibrium, ΔG = 0. |
5.1
Systems |
Question to answer:
Questions to ponder and questions for later:
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28-Jun-2012 |