4: Thermodynamics

Lecture 4 Announcements

• Today's Topics: Brown 5.1-5.7, nature of chemical energy, first law of thermodynamics, enthalpy, enthalpy of reaction, calorimetry, Hess's law, enthalpies of formation
• Problem Set 3: Due the day before Exercise #4 and upload to Moodle
• Problem Set 4: Posted on Moodle, due before Exercise #5
• Study Center: Wednesdays 6:00 PM - 8:00 PM in ETA F5
• Office Hours: Prof. Norris and Brisby, Thursdays 5:00 PM - 6:00 PM in LEE P210

Lecture 5

• Next Week's Topics: Brown 19.1-19.6, spontaneous processes, entropy, second law of thermodynamics, molecular interpretation, third law of thermodynamics, entropy changes in chemical reactions, Gibbs free energy, free energy and temperature

Review

• Lecture 3 Review: Aqueous solutions, electrolytes, nonelectrolytes, precipitation reactions, exchange reactions, molecular reactions, complete ionic reactions, net ionic reactions, acids, bases, neutralization reactions, oxidation-reduction reactions, oxidation numbers, oxidation states, activity series, and concentration: molarity

Today: Understanding Energy

• Topic: Hydrogen combustion, work, heat, calculation
• Question: How much work, how much heat, how to calculate?

Thermodynamics I

• Today's Topic: Study of energy and its transformation; relates heat, work, and chemical reactions; laws of thermodynamics, enthalpy, and entropy
• Purpose: Explains why reactions occur or don't occur
• Today's Topic: Thermochemistry, chemical reactions produce/absorb heat

Energy

• Two Types: Kinetic energy (motion) and potential energy (stored energy)
• In Chemistry: Potential energy stored in chemical bonds
• 1st Law of Thermodynamics: Energy can be converted between different types but cannot be created or destroyed
• System: Reactants and products
• Surroundings: Everything else

System Types

• Open: Matter and energy exchanged with surroundings
• Closed: Energy exchanged with surroundings, but not matter
• Isolated: Neither matter nor energy exchanged with surroundings

Total Energy for System

• Internal Energy (E): Total kinetic plus potential energy in system
• Instead: Focus on changes in E, ΔE = E_final - E_initial
• Usually: Occurs during a chemical reaction

Energy Changes

• Sign is Important: ΔE > 0 if system gains energy; ΔE < 0 if system loses energy
• For Chemical Reactions: E_final is for products and E_initial is for reactants

1st Law Revisited

• As energy cannot be created or destroyed: It must go somewhere
• ΔE = q + w where q = heat added to system, w = work done on system
• Sign Convention:
  • q > 0: heat added to system
  • q < 0: heat removed from system
  • w > 0: work done on system
  • w < 0: work done by system

Heat and Work: Sign Convention

• Internal Energy: Energy deposited into system or energy withdrawn from system
• ΔE > 0 if energy deposited to system, < 0 if energy withdrawn

Endothermic and Exothermic Processes

• Endothermic: q > 0, heat absorbed by system (ex. ice melting)
• Exothermic: q < 0, heat released by system (ex. H₂ combustion)
• State Functions: Internal energy, E, is an example; value depends on current state of system, not path taken to get there, also depends on quantity of matter (extensive property)

E is State Function

• ΔE: Only depends on initial and final state, not the path
• Contrast: q and w are not state functions, path matters

Enthalpy, H

• H = E + PV: Internal energy + (pressure·volume)
• E, P, and V are state functions: Thus, H is a state function
• Pressure-volume work: Work done against constant pressure

Change in Enthalpy

• ΔH = ΔE + Δ(PV): At constant pressure, ΔH = q_p = heat at constant pressure
• Chemistry is typically at constant P and qp: Can determine ΔH for a reaction
• ΔH > 0: System gained heat
• ΔH < 0: System lost heat

Enthalpies of Reaction

• ΔH = H_final - H_initial = H_products - H_reactants
• Enthalpy is extensive property: ΔH_forward = -ΔH_reverse
• ΔH_rxn depends on state (s, l, g): of reactants and products

Calorimetry

• Used to measure ΔH: Heat flows change T
• Heat Capacity: Heat required for ΔT = 1K for particular substance
• Molar heat capacity (C_m): J/(mol·K)
• Specific heat capacity (C_s): J/(g·K)

Ex: Constant Pressure

• System (reactants) and surroundings (H₂O) combined in insulated calorimeter
• If rxn is exothermic: q_rxn < 0, q_H₂O > 0
• If rxn is endothermic: q_rxn > 0, q_H₂O < 0

Ex: Constant Volume

• Can also seal vessel ("bomb")
• Then ΔE_rxn = q_rxn = −q_H₂O

Hess's Law

• If a reaction can be written in steps: ΔH_rxn = sum of enthalpy changes for the steps

Other Enthalpies

• ΔH_fus: Heat of fusion (melting)
• ΔH_vap: Heat of vaporization (boiling)

Thermodynamics, So Far...

• We have only considered ΔE and ΔH: But this is not the whole story
• We need additional information about spontaneity and entropy: To determine if a chemical process will occur
• Next Time: 2nd Law of Thermodynamics

What We Learned

• 1st Law of Thermodynamics: System and surroundings, internal energy (E), energy diagrams, ΔE = q + w, endothermic/exothermic processes, state functions (enthalpy - H), pressure-volume work, enthalpies of reaction, calorimetry, heat capacity, Hess's law, heats of reaction, heats of formation

This quiz covers topics from Brown 5.1-5.7, including chemical energy, thermodynamics laws, and calorimetry. Additionally, it reviews spontaneous processes, entropy, Gibbs free energy, and chemical reaction principles outlined in Lecture 4.