Unit 16: Thermodynamics
Why some reactions run themselves and others never will, told through counting arrangements and one equation weighing energy against disorder.
8 topics · 29 questions
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16.2 Spontaneity and the Second Law of Thermodynamics
- How it worksHow does the second law of thermodynamics require the total entropy of the universe to increase for any spontaneous process?
- Common mistakeWhy isn't a spontaneous process automatically a fast one?
- Going furtherWhat happens to a nonspontaneous process, like separating a mixture back into pure components, if you're willing to add outside energy?
Your turn1 of 2
Iron rusting is a spontaneous process, yet it can take years to happen noticeably. Why isn't a spontaneous process automatically a fast one?16.1 Entropy and Counting Microstates
- How it worksHow does counting the number of microstates available to a system connect directly to its entropy through Boltzmann's equation?
- How it worksWalk me through comparing the microstates available to gas molecules confined to half a container versus spread through the whole container.
- Worked exampleFind which of two systems has the higher entropy: one with 2 possible microstates or one with 1,000,000.
- Common mistakeWhy isn't entropy just a fancy word for messiness or disorder?
Your turn1 of 2
Using Boltzmann's equation S = k ln(W), where k = 1.38 times 10−23 J/K, find the entropy of a system with 696418 possible microstates.16.2 Predicting the Sign of Entropy Change
- How it worksHow do more possible microstates in a gas than a liquid, and a liquid than a solid, predict the sign of entropy change in a phase transition?
- Worked exampleFind the sign of the entropy change for the reaction 2 H2 gas plus O2 gas forming 2 H2O liquid.
- Common mistakeWhy doesn't dissolving always increase a system's entropy, even though mixing usually does?
- Going furtherWhat happens to a reaction's entropy change when the number of moles of gas increases from reactants to products?
Your turn1 of 2
What is the sign of the entropy change for the reaction 2 H2 gas plus O2 gas forming 2 H2O liquid?16.2 The Third Law of Thermodynamics
- How it worksHow does the third law give chemists an absolute zero point to measure every substance's standard entropy against?
- Common mistakeWhy isn't a substance's standard entropy value ever negative, unlike its enthalpy of formation?
- Going furtherWhat happens to the residual entropy of a real crystal that isn't perfectly ordered even at absolute zero?
Your turn1 of 2
Enthalpy of formation values are routinely negative for many compounds, but a substance's standard entropy value is never negative. Why is that?16.3 Gibbs Free Energy
- How it worksHow does Gibbs free energy combine a system's own enthalpy and entropy changes into one number that predicts spontaneity without tracking the surroundings separately?
- Worked exampleFind the sign of delta G for a reaction with delta H = negative 150 kJ and delta S = negative 200 J/K at 298 K.
- Common mistakeWhy isn't a reaction with a negative delta H automatically spontaneous?
Your turn1 of 2
Find delta G for a reaction with delta H = -126 kJ and delta S = -73 J/K at 390 K (delta G = delta H minus T delta S; convert delta S to kJ/K by dividing by 1000).16.3 How Temperature Flips a Reaction's Spontaneity
- How it worksHow does the equation delta G equals delta H minus T delta S let temperature flip a reaction's spontaneity depending on the signs of delta H and delta S?
- Worked exampleFind the temperature at which a reaction with delta H = 90 kJ and delta S = 200 J/K becomes spontaneous.
- Worked exampleFind whether a reaction with delta H negative and delta S positive is spontaneous at every temperature, checking all four sign combinations.
- Common mistakeWhy doesn't a positive delta S guarantee that a reaction is spontaneous at every temperature?
Your turn1 of 2
Find the temperature at which a reaction with delta H = 145 kJ and delta S = 212 J/K becomes spontaneous (delta G equals zero when T equals delta H divided by delta S, converting delta S to kJ/K).16.4 Standard Free Energy and the Equilibrium Constant
- How it worksHow does the equation delta G naught equals negative RT ln K link one standard-state number to an entire equilibrium constant?
- Worked exampleFind the equilibrium constant K for a reaction with a standard free energy change of negative 25.0 kJ/mol at 298 K.
- Common mistakeWhy doesn't delta G naught tell you which direction a reaction runs right now, only what happens under standard-state conditions?
- Ties togetherHow does the relationship between delta G naught and K connect back to Le Chatelier's principle predicting a reaction's shift?
Your turn1 of 2
Find the equilibrium constant K for a reaction with a standard free energy change of -16 kJ/mol at 306 K (delta G naught equals negative RT ln K, R = 8.314 J per mol per K).16.4 Free Energy Under Nonstandard Conditions
- How it worksHow does the equation delta G equals delta G naught plus RT ln Q let real, nonstandard concentrations override the standard-state prediction?
- Worked exampleFind delta G for a reaction with delta G naught = 5.0 kJ/mol at 298 K when Q = 0.010.
- Common mistakeWhy isn't delta G naught the number that decides whether a reaction proceeds under real, everyday conditions?
- Going furtherWhat happens to delta G as a reaction proceeds toward equilibrium and Q approaches K?
Your turn1 of 2
Find delta G for a reaction with delta G naught = 14 kJ/mol at 306 K when Q = 0.466 (delta G equals delta G naught plus RT ln Q, R = 8.314 J per mol per K).