Chem 1465 1st Edition Lecture 13Outline of Last Lecture 1. Enthalpy (H)A. Important relationships B. Thermodynamic propertyC. ΔH and q2. Thermochemical equationsA. EquationB. Three rules for thermochemical equations3. Hess’s law4. Enthalpies of formationOutline of Current Lecture 1. SpontaneityA. Spontaneous processB. Predicting if a process is spontaneous 2. EntropyA. DefinitionB. Boltzmann related microstates to entropy 3. Second law of thermodynamics A. DefinitionB. Quantifying entropyC. At constant pressure 4. Third law of thermodynamicsA. DefinitionB. Standard molar entropy These notes represent a detailed interpretation of the professor’s lecture. GradeBuddy is best used as a supplement to your own notes, not as a substitute.5. Gibbs free energy A. Equations B. Gibbs free energy and spontaneity Current Lecture1. SpontaneityA. Spontaneous process: a process that occurs by itself without any continuous intervention. Spontaneous does not mean quickly. Examples would be water below zero degrees Celsius forming ice, or in reverse, ice about zero degrees Celsius forming liquid water. B. How can we predict if a process is spontaneous? By using thermodynamics. Based onenthalpy- if ΔH is less than zero then it is exothermic and if ΔH is greater than zero, it is endothermic. Maybe exothermic processes are spontaneous. However, some endothermic reactions are spontaneous. 2. EntropyA. Definition: the driving force for a spontaneous process is an increase in entropy. It is viewed as a measurement of molecule randomness or disorder. Nature spontaneously proceeds toward a state that has highest probability or likelihood. B. Boltzmann related microstates to entropy: s = kb l Ω where s is entropy, kb is Boltzmann’s constant (1.38 x 10-23), and Ω is the number of energetically favorable ways to arrange components in a system. 3. The second law of thermodynamicsA. Definition: for any spontaneous process the entropy of the universe increases. B. Quantifying entropy: when a system exchanges heat with the surroundings it changes the entropy of the surroundings. C. At constant pressure we can use heat of the system to quantify ΔS of surroundings- A process that emits heat into the surrounding (heat of system is negative) increases the entropy of the surroundings (ΔS of surroundings is positive)- A process that absorbs heat from the surrounding (heat of system is positive) decreases the entropy of the surroundings (ΔS of surroundings is negative)- Magnitude of the heat of the system is proportional to magnitude of ΔS of the surroundings. - ΔSsurr = -qsys / T - ΔS = -ΔH / T4. The third law of thermodynamics A. Definition: entropy of a perfect crystal at 0 kelvin is zeroB. Standard molar entropy: Sᵒ is SME. This is the entropy for one mole of reaction. Because entropy is a state function it is path independent.- ΔSᵒrxn = Σ np (Sᵒproducts) – Σ nr (Sᵒreactants)5. Gibbs free energyA. Equations- G = H – TS- At constant temperature and pressure- changes in this equation allow for prediction of whether a process is spontaneous: ΔG = ΔH – (TΔS)- Initial equation: ΔSuniv = ΔSsys + ΔSsurr - We can change it to: ΔG = -TΔSuniv B. Gibbs free energy and spontaneity- ΔG is negative then the reaction is spontaneous in the forward direction- ΔG is positive then the reaction is nonspontaneous in the forward direction (but spontaneous in the reverse direction)- ΔG is zero then the reaction is at its thermodynamic