Chemical Thermodynamics

Transcription

1 Chemical Thermodynamics Overview Everything in the world is a balance of energy, in various forms from biological processes to the rusting of a nail. Two of the most important questions chemists ask are: How fast does the reaction proceed? and How far does it proceed?. These questions were addressed in unit 8 and unit 9,10 & 11 respectively. In truth, the rate of chemical reaction and equilibria is closely tired to energy. In this section we will look at chemical thermodynamics which discusses the relationship between various forms of energy. Spontaneous Processes The first law of thermodynamics states that energy is conserved. In other words energy is neither created nor destroyed in any processes whether that process is a brick falling, a candle burning, or an ice cube melting. Energy can be transferred between a system and the surrounding and can be converted from one form to another, but the total energy of the universe remains constant. The first law of thermodynamics says nothing about how energy moves, just that it must be conserved. A spontaneous process is one that proceeds on its own without any outside assistance. For example at room temperature ice will spontaneously melt. However, the reverse process (freezing) is nonspontaneous. For any chemical reaction that is spontaneous at given conditions, the reverse will be nonspontaneous. Also, because a process is spontaneous doesn t mean that it will take place at an observable rate: it could be explosive or it could be very slow. Finally, a nonspontaneous process doesn t mean it won t happen it just means it won t happen under the current conditions. Predict whether each process is spontaneous or nonspontaneous: (a) Water at 40 C gets hotter when a piece of metal heated to 150 C is added (b) Water at room temperature decomposes into H 2 (g) + O 2 (g) (c) Benzene vapour, C 6 H 6 (g) at a pressure at 1 atm condenses to liquid benzene at the normal boiling point of benzene, 80.1 C

2 Reversible and Irreversible Processes A reversible process is a specific way in which a system changes its state. In a reversible process, the change occurs in such a way that the system and surroundings can be restored to their original states by exactly reversing the change. In other words, we can restore the system to its original condition with no net change to either the system or its surroundings. An irreversible process is one that cannot simply be reversed to restore the system and its surroundings to their original states. Chemical changes are irreversible processes. For example, if you place a hot object next to a cold one the laws of thermodynamics say that the heat will flow from the hot object to the cold object. You cannot make the reverse happen. This is an irreversible process. Let s imagine a scenario where the above change could be reversible: If two objects are placed side by side with infinitesimally small differences in temperature, so close in temperature that most people would say they are the same temperature. When placed side by side the two objects exchange heat at an infinitely slow rate. This process would be reversible since the change was extremely small and the time it occurred was very large. In short, any change we can observe is happening too fast to be considered reversible. To generalize further, any spontaneous process (any process that occurs without work input) is irreversible. Entropy Knowing that any spontaneous process is irreversible, can we make a prediction about whether a process is spontaneous or not if the process is unfamiliar? To do this we must understand a thermodynamic process called entropy. Entropy is a measure of the randomness of a system. That is, it tells us how much disorder there is in a system. Take a jar of marbles. In the jar the marbles are ordered, the balls are sitting together closely packed. If the jar breaks the marbles will spread out, separating from each other and increase their randomness. This would signal an increase in entropy. Quantitatively, entropy is a state function (like enthalpy) and is a characteristic of the system and the chemicals it contains. In experiment, the temperature of the system changes only by small amounts (with the exception of extremely exothermic processes). We call this an isothermal process. Assuming a process is reversible and the temperature is constant we can show that entropy is the heat change divided by the temperature the process is taking place, as shown to the right. Elemental mercury is a silver liquid at room temperature. Its normal freezing poijnt is 38.9 C and its molar enthalpy of fusion is H fusion = 2.29 kj/mol. What is the entropy of the system when 50.0g of Hg(l) freezes at the normal freezing point?

3 The Second Law of Thermodynamics As described earlier, the first law of thermodynamics is that energy must always be conserved in any process. Does that mean that entropy must also be conserved? Let s look at this situation by calculating the entropy change of a system including the entropy change of its surroundings when our system is 1 mol of ice melting in the palm of your hand. This process is irreversible. Given the enthalpy of fusion for water is 6.01 kj/mol and that your hand is at bond temperature, 37. Determine the change of entropy of the universe fore the melting of ice described above. So the second law of thermodynamics states that: The molecular Interpretation of Entropy and the Third Law of Thermodynamics Until now we have looked at the large scale setting, but how does Entropy changes we have noted look on the molecular level. The randomness of molecules is actually easy to understand when looking at the different states of matter. A solid has more order than a liquid which has more order than a gas. Ludwig Boltzmann decided to look at individual molecules and describe changes in entropy based on the smallest pieces of matter, Let s take a molecule of water. A single molecule of water can have different orientations: Each configuration is called a microstate, W

4 Boltzmann described these microstates to better describe the entropy of a system. To him, it made sense to say that the more microstates in a system the greater entropy of that system. He related the connection between entropy, S, and microstates by writing: where k = Boltzmann constant 1.38 x J/K So entropy of a system will increase for any change that takes place that increases the number of microstates. If the number of microstates decreases, so will the entropy of the system. So we can make qualitative predictions about a chemical reaction based on our predictions about how the change is system will increase the number of microstates: (1) The entropy of a system increases with increasing temperature. (2) The number of microstates increases with an increase in volume (3) The number of microstates increase with an increase in the number of moles (molecules). The Third Law of Thermodynamics states that the entropy (disorder) of a system approaches 0 J/K as the temperature of the system approaches absolute zero ( C or 0K). At this point all molecular motion stops so there is only 1 microstate. And there fore the ln (W) = ln (1) = 0. Predict the S for each process, assuming each occurs at constant temperature (a) H 2 O(l) H 2 O(g) (b) Ag +1 (aq) + Cl -1 (aq) AgCl(s) (c) 4 Fe(s) + 3O 2 (g) 2Fe 2 O 3 (s) (d) N 2 (g) + O 2 (g) 2NO(g)

5 Entropy Changes in Chemical Reactions Chemical reactions all experience changes is disorder. As is the case with enthalpy, there are standard molar entropies (measured in joules per Kelvin). These values are tabulated in Appendix C. you should note a few things about entropy, including some specific differences from enthalpy. 1. Unlike enthalpies of formation, standard molar entropies of elements at the reference temperature of 298K are NOT zero. 2. The standard molar entropies of gases are greater than those of liquids and solids, consistent with our interpretation of experimental observations. 3. Standard molar entropies generally increase with increasing molar mass. 4. Standard molar entropies generally increase with an increasing number of atoms in the formula of a substance. Using Hess s law we can determine the molar entropy change for a chemical process using these values in a similar manner as we did for enthalpy. Calculate the change in the standard entropy of the system, S, for the synthesis of ammonia from N 2 (g) and H 2 (g) at 298K N 2 (g) + 3H 2 (g) 2NH 3 (g)