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Understanding Gibbs Free Energy, Phase Changes, and Chemical Equilibrium in Thermodynamics

This article explores fundamental concepts in thermodynamics, including the formula for Gibbs Free Energy (ΔG = ΔH - TΔS) and its variables. You'll learn how to determine if a reaction is spontaneous, exothermic, or endothermic, as well as concepts like entropy. The article also covers the properties of solids, liquids, and gases, outlines endothermic and exothermic phase changes, and explains how to read phase change diagrams. Additionally, it discusses chemical equilibrium, Le Chatelier’s principle, and how to calculate equilibrium constants.

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Understanding Gibbs Free Energy, Phase Changes, and Chemical Equilibrium in Thermodynamics

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  1. Daily Sciencepg. 106 • What is the formula for Gibbs Free energy? • What does each variable represent? • How can you tell if a reaction will be spontaneous? • How can you tell if a reaction is exothermic or endothermic? • What is entropy? • Determine if the following reaction will be spontaneous/nonspontaneous, exo/endo, increase/decrease is entropy • ΔH= -26700 J, T= 387, ΔS= -234

  2. Daily Science pg. 106 • Name one property of a solid, liquid, and gas. • Separate the phase changes into endothermic and exothermic processes. • How do you find the boiling point of a substance on a phase change diagram? • Draw a sketch of a phase change diagram and label the solid, liquid, and gas phases. • How do you know a state change is taking place on a freezing and boiling point graph? • How many moles of a gas are present at a pressure of 798 torr, a volume of 278 mL, and a temperature of 45 oC? • How many liters of NH3 can be made from 5.6 g of H2? • N2+ 3H2 2NH3 • What is the specific heat of a 4.3 g piece of metal that absorbs 1987 J of heat energy and the temperature changed from 16 oC to 54 oC • Using the Gibbs free energy equation, describe the following process: • ΔH= -25 KJ, T= 124 K, and ΔS= 5.o KJ/K

  3. Pg. 105 Chemical Equilibrium

  4. Equilibrium • Most reactions do not go to completion. They are reversible • Use a double arrow to show that an equation can be reversed • Chemical equilibrium happens when the forward and backward reactions balance each other because they take place at equal rates • As the concentration of reactants decrease, concentration of products increase

  5. Example • The formation of NH3. Used in fertilizers • N2 + 3H2 2NH3

  6. Determining equilibrium (homogeneous) • Can determine equilibrium constant- Keq • Concentrations of the products raised to the power of their coefficients/ concentrations of the reactants raised to the power of their coefficients • aA + bBcC + dDKeq= ([C]c[D]d)/ ([A]a[B]b) • K > 1: more products that reactants @ equil • K < 1: more reactants than products @equil

  7. Example • Write the expression for finding the equilibrium constant for the following reaction: • N2 + 3H2 2NH3 • Find the equilibrium constant if the concentration of N2 is 0.533 mol/L, H2 is 1.600 mol/L, and NH3 is 0.933 mol/L

  8. Le Chatelier’s principle • Says if a stress is applied to a system at equilibrium, the system shifts in the direction that relieves the stress • Can predict how changes in temp, concentration, and volume (pressure) will effect equilibrium • V and P related- if you increase the volume, you decrease the pressure • Longer arrow shows which way equilibrium will shift ( or )

  9. Changes in concentration • An increase in the concentration of a reactant shifts the reaction toward the product • Ex. CO +3H2 CH4 + H2O • If products are removed, it will shift toward the products

  10. Changes in volume • If you decrease the volume (increase pressure) reaction will shift to the side with the smaller amount of moles. • Ex. CO +3H2 CH4 + H2O

  11. Changes in temperature • Alters both equilibrium position and constant • The reaction will shift in the opposite direction when the temperature is added • Exothermic shift to left • Endothermic shift to right

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