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# Chapter 17.2 : Driving Force of Reactions

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• 1. Chapter 17.2
• 2. 1) Explain the relationship between enthalpy change and the tendency of a reaction to occur. 2) Explain the relationship between entropy change and the tendency of a reaction to occur. 3) Discuss the concept of free energy, and explain how the value of this quantity is calculated and interpreted. 4) Describe the use of free energy change to determine the tendency of a reaction to occur.
• 3. • The great majority of chemical reactions in nature are exothermic. • The tendency throughout nature is for a reaction to proceed in a direction that leads to a lower energy state. • Some endothermic reactions do occur spontaneously. • Something other than enthalpy change can help determine whether a reaction will occur.
• 4. • Melting is one example of a naturally occurring endothermic process. • An ice cube melts spontaneously at room temperature as energy is transferred from the warm air to the ice. • The well-ordered arrangement of water molecules in the ice crystal is lost, and the less-ordered liquid phase of higher energy content is formed. • A system that can go from one state to another without an enthalpy change does so with an increase in entropy.
• 5. • The decomposition of ammonium nitrate: 2NH4NO3(s) 2N2(g) + 4H2O(l) + O2(g) • On the left side are 2 mol of solid ammonium nitrate. • The right-hand side of the equation shows 3 mol of gaseous molecules plus 4 mol of a liquid. • The arrangement of particles on the right-hand side of the equation is more random than the arrangement on the left side and hence is less ordered.
• 6. • There is a tendency in nature to proceed in a direction that increases the randomness of a system. • A random system is one that lacks a regular arrangement of its parts. • This tendency toward randomness is called entropy. • Entropy, S, can be defined in a simple qualitative way as a measure of the degree of randomness of the particles, such as molecules, in a system.
• 7. • To understand the concept of entropy, consider the comparison between particles in solids, liquids, and gases. • In a solid, the particles are in fixed positions, and we can easily determine the locations of the particles. • In a liquid, the particles are very close together, but they can move around. Locating an individual particle is more difficult. The system is more random, and the entropy is higher. • In a gas, the particles are moving rapidly and are far apart. Locating an individual particle is much more difficult, and the system is much more random. The entropy is even higher.
• 8. • Absolute entropy, or standard molar entropy, of substances are recorded in tables and reported in units of kJ/(mol•K). • Entropy change, which can also be measured, is defined as the difference between the entropy of the products and the reactants. • An increase in entropy is represented by a positive value for ∆S • A decrease in entropy is represented by a negative value for ∆S
• 9. • Processes in nature are driven in two directions: 1. toward least enthalpy 2. toward largest entropy • As a way to predict which factor will dominate for a given system, a function has been defined to relate the enthalpy and entropy factors at a given temperature and pressure. • This combined enthalpy-entropy function is called the free energy, G, of the system; it is also called Gibbs free energy.
• 10. • Only the change in free energy can be measured. It can be defined in terms of enthalpy and entropy. • At a constant pressure and temperature, the free-energy change, ∆G, of a system is defined as the difference between the change in enthalpy, ∆H, and the product of the Kelvin temperature and the entropy change, which is defined as T∆S: ∆G0 = ∆H0 – T∆S0
• 11. ∆G0 = ∆H0 – T∆S0 • The expression for free energy change is for substances in their standard states. • The product T∆S and the quantities ∆G and ∆H have the same units, usually kJ/mol. • If ∆G < 0, the reaction is spontaneous. • ∆H and ∆G can have positive or negative values. This leads to four possible combinations of terms.
• 12. For the reaction NH4Cl(s) → NH3(g) + HCl(g), at 298.15 K, ∆H0 = 176 kJ/mol and ∆S0 = 0.285 kJ/(mol•K). Calculate ∆G0, and tell whether this reaction is spontaneous in the forward direction at 298.15 K. Given: ∆H0 = 176 kJ/mol at 298.15 K ∆S0 = 0.285 kJ/(mol•K) at 298.15 K Unknown: ∆G0 at 298.15 K Solution: The value of ∆G0 can be calculated according to the following equation: ∆G0 = ∆H0 – T∆S0 ∆G0 = 176 kJ/mol – 298 K [0.285 kJ/(mol•K)] ∆G0 = 176 kJ/mol – 84.9 kJ/mol ∆G0 = 91 kJ/mol Does Not occur naturally