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Chem 1046 Chapter 20

by: Kaitlyn Michaud

Chem 1046 Chapter 20 Chemistry 1046

Kaitlyn Michaud
Virginia Tech
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About this Document

These notes cover the first sections of chapter 20 and what will be on our next test. Thermodynamics, entropy, enthalpy, and Gibb's Free Energy are all expanded on.
General Chemistry
Dr. Neidigh
Class Notes
Chemistry, chem 1046, entropy, Enthalpy, thermodyanmics, gibb's free energy, Virginia Tech




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This 5 page Class Notes was uploaded by Kaitlyn Michaud on Friday April 1, 2016. The Class Notes belongs to Chemistry 1046 at Virginia Polytechnic Institute and State University taught by Dr. Neidigh in Spring 2016. Since its upload, it has received 39 views. For similar materials see General Chemistry in Chemistry at Virginia Polytechnic Institute and State University.

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Date Created: 04/01/16
Thermodynamics, Entropy, Free Energy, and the Direction of Chemical Reactions Chapter 20 beings with an introduction to the 2 ndLaw of Thermodynamics, but don not forget that the first law of thermodynamics is essentially energy cannot be created or destroyed, but rather is conserved. I. The Second Law of Thermodynamics: Predicting Spontaneous Change  A spontaneous change occurs under specified conditions, and does not require a continuous input of energy—think of a fire. o Under certain conditions it is very easy to start a fire, and then the fire will continue to burn spontaneously even though you are not putting more energy into it.  NOTE that when we say “spontaneous” we mean that if the process can occur on its own naturally, then it will. If the conditions are optimal for a brush fire then one will start. Spontaneous does NOT mean instantaneous or reveal anything about the time it will take for a change to occur.  A nonspontaneous change requires the surroundings to continuously supply energy—think of a bike. o For a bike to move you have to constantly input your own energy into peddling it.  Under given conditions, these changes will be opposing; if a change is spontaneous in one direction, it will NOT be spontaneous in the opposite direction. o Referring back to our fire example, you have to initially cause a spark for the fire to then spontaneously keep burning, however, you can’t go in the opposite direction of this process. The fire does not just appear. II. Relating Spontaneity to the First Law of Thermodynamics  Remember what we said about conservation of energy at the top of the page? Energy, (ex. In the form of heat) when released by some exothermic process will be absorbed by the surroundings. The heat doesn’t just disappear. o Ex. A chemical reaction (the system) releases heat into the surroundings. Anything not the system is considered the surroundings. o The system PLUS the surroundings = the whole universe.  Therefore, the total energy of the universe is constant (since energy is constantly being absorbed and released), so the change in energy of the universe is ZERO. Thermodynamics, Entropy, Free Energy, and the Direction of Chemical Reactions  SO, we have defined that energy is not created or destroyed, and the energy of the universe is constant and equal to 0, but how does this relate to spontaneity? o The first law accounts for the energy, but NOT the direction of a process, that being spontaneous or not spontaneous.  Ex. An ice cube melting in your hand is spontaneous and the energy is conserved, but the reverse change can never happen. The ice will not resolidify in your hand using the energy it just released via melting. So the first law cannot predict the direction of a spontaneous change. III. ΔH  Recall that enthalpy is denoted, “H,” so delta H is the change in enthalpy. When examining enthalpy of chemical reactions, a negative ΔH indicates an exothermic reaction, and a positive ΔH indicates an endothermic reaction. But does ΔH indicate spontaneity? o The answer is NO. Remember the definition for spontaneity says, “…under specified conditions…” o Some processes will be spontaneous under exothermic reactions, and others under endothermic reactions. ΔH does not equate temperature to spontaneity.  Examples: An ice cub melting in your hand is a spontaneous EXOthermic reaction. Conversely, most soluble salts dissolve ENDOthermically and spontaneously.  So, with the first law of thermodyanimcs, ΔH also does not predict the direction of spontaneity. IV. Direction of a Spontaneous Reaction and Freedom of Motion  We’ve discussed what does not predict the direction of a spontaneous change, but not what does predict—that being freedom of particle motion! o If you think back to our examples, both exothermic and endothermic, there is a common theme in all that can predict the direction of spontaneity, and that is: a process will favor spontaneous change if the freedom of motion of particles in a system increases with the change.  With our ice example, the molecules of water have more freedom to move in liquid form than solid. In solid, the molecules are held in a rigid crystalline structure. When the ice melts and releases heat, the molecules can move around in a fluid medium, they have more freedom of motion. Thermodynamics, Entropy, Free Energy, and the Direction of Chemical Reactions  So the spontaneous change from solid to liquid favored more freedom of motion of particles.  With our soluble salts example, when the salts dissolve in the solvent, the individual ions have more freedom to move when not held together in a molecular structure.  Thus again, the spontaneous change (though endothermic) was favored because freedom of motion of particles was higher in the end state.  Direction of spontaneous reaction can be determined by examining freedom of motion of particles in the start v. end phase. V. Entropy and Microstates  Quantization energy. When thinking of individual atoms, quantization energy includes the rotational, vibrational, and translational states (kinetic energy states)  Microstates. You can expand this thinking of quantization energy of atoms to a larger system, such as molecules. Each quantized state of the system is called a microstate, and the total energy of a system can be dispersed throughout one microstate.  Dispersal Energy. Each microstate has the same total energy as another, so each microstate is equally possible of occurring in a system. o Now that you have some definitions that sound very confusing, here is what you need to know about how they affect spontaneous changes:  A system with fewer microstates has a lower entropy  A system with more microstates has a higher entropy  These can be related through the equation: S= (k)ln(W) -23 Where S is entropy, k is the Boltzmann Constant 1.38 x 10 , and W is the # of microstates.  A greater number of microstates essentially means there are more ways to disperse the energy throughout the system and vice versa.  How can you tell if a microstate is increasing or decreasing? o Like we said, a microstate relates to the rotational, translational, or vibrational state of the system. Let’s look at some examples:  Why does gas expand to fill the volume of its container? (Refer to figure 20.2 in text). When you Thermodynamics, Entropy, Free Energy, and the Direction of Chemical Reactions increase the volume that a gas can occupy, you increase the translational energy levels because the particles now have more space to move freely. This in turn increases the number of microstates each particle can occupy in the larger volume, so more microstates=higher entropy.  How do microstates relate to temperature? o Think back to how temperature affects the phase of matter and how the molecules move. A piece of ice has very little freedom of motion regarding molecules. However, when you melt the ice you gain more freedom of movement in liquid form, and when you vaporize liquid you have the highest freedom of motion in gaseous form. As you continue to heat the gas particles, the molecules will move more rapidly—more microstates can be occupied= higher entropy. VI. Overview of Entropy and the Second Law  A change in entropy determines the direction of a spontaneous reaction, but you cannot look at just the system, but also the surroundings. nd  Another way to state the 2 law is: all natural processes occur spontaneously in the direction that increases the entropy of the universe (system + surroundings). o This highlights why you cannot just look at the entropy of the system. The system could decrease in entropy, but the surroundings could increase in entropy, so you have to look at both. REGARDLESS, real processes will occur in a fashion that increases the overall entropy of the universe. o Remember, energy is always conserved so the energy of the universe=0, but the entropy of the universe is always increasing VII. Entropy and the Third Law of Thermodynamics  The third law says that entropy CAN = 0. A perfect crystal can have zero entropy at absolute zero (-273 C or 0 K). o Emphases on a PERFECT crystal, meaning the particles are aligned flawlessly. At absolute 0, there so such minimal movement, almost virtually none, that there is only 1 microstate. If you refer to the equation on the previous page, S= k ln w. o The natural log of 1=0, so if there is only 1 microstate then entropy equals zero.  Standard Molar Entropies Thermodynamics, Entropy, Free Energy, and the Direction of Chemical Reactions o Entropies are compared at standard states: 1 atm (gases), 1 M (solutions), and the pure substance for solids and liquids. o Standard molar entropy (S°) is defined in units of J/mol x K o Standard molar entropy is directly affected by: temp, physical state, dissolution, atomic size, and molecular complexity. VIII. Gibb’s Free Energy  ΔS and ΔH combine to form the Gibb’s free energy equation. The sign of Gibb’s Free energy indicates the spontaneity of the reaction. Spontaneous Process Negative value Non spontaneous process Positive value Equilibrium Zero  G° = H° – TS° at STANDARD conditions. If the temperature is 25 degrees C and pressure is at 1 atm you can use an alternate version of this equation: G° = = nG° – f (products) nG° f (reactants) o In this equation, G°fis the free change that occurs when 1 mole of compound is formed from its element. (G° ff an element in its standard state=0) o N= the stoichiometric coefficient of any substance.  When calculating under non-standard conditions: G = G° + RT ln Q o G in this equation = free energy then temp is not 25 degrees C, pressure is not 1 atm, and molarity is not equal to 1. However, G° is the free energy under these conditions so once you add RTlnQ to that you get the Gibb’s free energy under different conditions. o R=8.314 X 10 (constant) o Q= reaction quotient, [products]/[reactants] o T= temp in Kelvin  At equilibrium: o Q=K and G=0


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