CHEM 162: Gilbert Chapter 14 page 1 Chapter 14: Thermodynamics: Spontaneous Processes, Entropy, and Free Energy Problems: 14.1-14.62, 14.69-14.73, 14.75-14.83, 14.86-14.87, 14.91-14.92 Why do some reactions occur but others don’t? We can answer these questions using thermodynamics, the study of energy transformations. 1st Law of Thermodynamics: Energy is neither created nor destroyed. – Essentially the law of conservation of energy – While energy is converted from one form to another, the energy of the universe is constant. system: that part of the universe being studied surroundings: the rest of the universe outside the system In an exothermic reaction like methane burning, CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(g) the bonds formed in the products are stronger than the bonds broken in the reactants. → The difference in energy between the bonds broken and the bonds formed is released to the surroundings. → Heat flows from the system to the surroundings. In an endothermic reaction like the following, N2(g) + O2(g) → 2 NO(g) more energy is required to break the stronger bonds in the reactants compared to the bonds formed in the reactants. → The difference in energy must be absorbed from the surroundings for the reaction to occur. → Heat flows from the surroundings to the system when the reaction occurs.CHEM 162: Gilbert Chapter 14 page 2 14.1 SPONTANEOUS PROCESSES AND ENTROPY spontaneous process: takes place naturally, without continuous outside intervention – e.g., a drop of food coloring will spread in a glass of water or once ignited H2 gas burns in O2 gas. nonspontaneous process: only takes place as a result of continuous intervention – e.g., electricity is required to convert water to H2 gas and O2 gas Keep in mind that spontaneous processes may be fast or slow. a. An explosion is fast and spontaneous. b. The process of rusting is slow and spontaneous. Consider the following spontaneous processes at room temperature: • A drop of food coloring will spread in a glass of water. • Methane (CH4) burns in O2 gas. • Ice melts in your hand. • Ammonium chloride dissolves in a test tube with water, making the test tube colder. Example: What do these processes or reactions have in common? Are they all exothermic? Are they all endothermic?CHEM 162: Gilbert Chapter 14 page 3 The result of all of these processes/reactions is an increase in entropy. Entropy (S): a measure of the molecular randomness or disorder of a system – A thermodynamic function that describes the number of arrangements (positions and/or energy levels) available to a system at a given instant. The more ways a particular state can be achieved, the more likely or the greater the probability of finding the system in that state. For example, imagine a brand new deck of cards with the jokers removed. – If the deck is divided and shuffled, the deck goes from being ordered to disordered; the more it’s shuffled, the more disordered it becomes. But why? – In 1877 Ludwig Boltzmann, an Austrian physicist, introduced the concept of entropy to explain why cards become more disordered the more they’re shuffled. – He determined there were 8.066×1067 ways a deck of cards can be organized compared to only one way for it to be perfectly ordered like a new pack of cards. → There are more states for a disordered deck of cards. → When shuffled a deck of cards is more likely to be disordered. → Nature proceeds towards the states that have the highest probabilities of existing. → Thus, the universe tends towards the most probable disordered states. Ex. Consider the ideal gas in the following image. Explain what will happen when the valve between the bulbs is opened. Now, consider why the gas spreads. – Consider the many arrangements resulting in about the same number of atoms in each of the two bulbs compared to all of the atoms gathering in one bulb or many more atoms in one bulb versus the other. → If we consider a particular arrangement of particles a microstate, then many more microstates exist where particles are evenly distributed in both bulbs and a much smaller number of microstates exist where the particles are in only one bulb or mostly concentrated in one bulb. → Thus, if a gas is placed in one bulb of an empty container, and the connecting valve is opened, the gas will spontaneously expand to fill the entire vessel. → However, the opposite process (the gaseous particles filling a vessel all moving to one bulb leaving the other bulb completely or mostly empty), while not completely impossible, is highly improbable.CHEM 162: Gilbert Chapter 14 page 4 The expansion of a gas also demonstrates the idea of positional probability, a type of probability that depends on the number of arrangements in space that yield a particular state. – The greater the space or volume available to particles, the higher the positional entropy. → For example, positional entropy increases from solid to liquid and from liquid to gas. – Solids have the smallest volume and relatively few positions available for particles. – Gases have much greater volume with many positions available for gas particles. Positional Entropy and Solutions – The formation of a solution—when a solid dissolves (e.g., NaCl dissolving in water) or two liquids mix—increases the positional entropy of a system. → The overall volume available to each component increases. → Thus, the number of positions available to the particles in the system increases.CHEM 162: Gilbert Chapter 14 page 5 Entropy and Energy States – Positional entropy results from the different energy levels available to a system. – To understand entropy we must recognize how atoms and molecules behave at the molecular level. For example, the kinetic energy of a molecule can be in any of the following forms: – translational: motion through space – rotational: motion about its center of mass – vibrational: stretching, compression, bending and twisting of chemical bonds The more complex the molecule, the more energy levels available to it. → The higher its entropy. – Compare the limited vibrational energy available for the NO molecule compared to the various forms of vibrational energy available for the NO2 molecule below. Consider an O2 molecule and the