AP Chemistry Unit 6:
Thermodynamics

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6.1 Endothermic and Exothermic Processes

• A change in temperature indicates an energy change
• Heat

• Heat is the transfer of energy between 2 objects due to a temperature difference.
• High-temperature transfers heat to low-temperature (second law of thermodynamics)
• Temperature

• The measure of the average kinetic energy (KE)
• Warmer = Higher KE
• Cooler = Lower KE

• Transfer of heat

• Heating or cooling of a substance
• Phase changes
• Dissolving solutes
• Chemical reactions

• System

• The specific part of the chemical reaction (where the reaction takes place)

• Surroundings

• Everything outside the system

• Endothermic (+∆H°)

• Heat flows into the system into the surroundings
• Breaking of a bond or attractive force

• Heating an object
• Melting, Vaporization, sublimation
• Chemical reactions and dissolving of a solute that feels cold

• Exothermic (–∆H°)

• Heat flows out of the system to the surroundings
• Forming of a bond or attractive force

• Cooling of an object
• Freezing, Condensation, Deposition
• Chemical reactions and dissolving of a solute that feels hot

        Source: thoughtco.com

6.2 Energy Diagrams

• 
  Source: theory.labster.com

• Activation Energy

• The minimum amount of energy required to overcome the initial barriers for a chemical reaction to occur.

• ∆H = HP–HR

• Energy released (–∆H) and Energy absorbed (+∆H)

6.3 Heat Transfer and Energy Equilibrium

• Heat Transfer

• Heat flows from warmer objects (higher KE) to cooler objects (lower KE), transferring Kinetic energy until both objects will have the same averge kinetic energy (temperature), reaching the state of thermal equilibrium.
  

6.4 Heat Capacity and Calorimetry

• Specific Heat Capacity (c)

• Amount of heat (J) to change 1 gram of a substance by 1° (Celsius or Kelvin)
• J/°C g or J/K g

• Molar Heat Capacity

• Amount of heat (J) to change 1 mole of a substance by 1°
• J/°C mol or J/K mol

• Heat Capacity (CP)

• Amount of heat (J) to change a substance by 1°
• J/°C or J/K

• Trends

• Lower Heat Capacity → Less Energy Required
• Higher Heat Capacity → More Energy Required

6.5 Energy of Phase Changes

• Phase Change

• Physical process where a substance transitions between phases of matter.

• Heating Curve (endothermic)

• 
  Source: cognitoedu.org
• X-axis

• Reaction progress over time

• Y-axis

• Change in temperature

• Plateaus (change of state)

• Represent the place where energy is required to break the intermolecular forces to allow for the transition between phases.
• Temperature does not change
• Melting (∆Hfusion) vs Vaporizing (∆Hvaporization)

• Where heat energy is used to break apart the IMFs
• ∆Hvaporization is almost always greater than the ∆Hfusion because it takes more energy for larger intermolecular separation

• Cooling Curves (exothermic)

• 
  Source: cognitoedu.org
• Essentially identical to heating curves but with ∆Hcondensation and ∆Hfreezing, respectively, which have opposite signs of the enthalpies of fusion/vaporization

6.6 Introduction to Enthalpy of Reaction

• Enthalpy (extensive)

• Heat
• ∆Hrxn

• The heat of reaction/reaction’s change in enthalpy
• Endothermic (∆H+)

• Generally thermodynamically unfavorable

• Exothermic (∆H–)

• Generally thermodynamically favorable

6.7 Bond Enthalpies

• Chemical Bond

• A force that holds groups of atoms together

• Bond Energy/Bond Enthalpy (BE)

• Potential energy stored in a chemical bond
• Breaking bonds requires energy (endothermic)
• Forming bonds requires energy (exothermic)

• ∆Hrxn = ∑(Bonds Broken BE) – ∑(Bonds Formed BE)
                  energy required                energy released

6.8 Enthalpy of Formation

• Standard State

• Standard Enthalpy (∆H°)

• Pressure is exactly 1 atmosphere for gases
• Concentration is exactly 1 molar for solutions
• Temperature is exactly 25°C

• Standard Heat of Formation

• Amount of heat needed to form 1 mole of a compound from its elements in the standard state
• ∆H°rxn = ∑np∆Hf°(Products) – ∑nr∆Hf°(Reactants)

• Heats of formation to figure out the standard heat of reaction
• np and nr being the stoichiometric coefficients of the reactants and products, respectively

• Bond Enthalpies vs. Heat of Formation

• Bond enthalpies are usually positive.
  ∆H = reactants – products
• Enthalpy of formation can be both positive and negative.
  ∆H = products – reactants

6.9 Hess’s Law

• Hess’s Law

• States that enthalpy is a state function where the change in enthalpy is the same whether the reaction takes place in one step or a series of steps
• If multiple reactions are combined, ∆Hrxn would be the individual reaction’s enthalpy changes combined
• If a reaction is reversed, ∆H is also reversed (opposite sign)
• If a reaction’s coefficients are multiplied/divided by a factor, ∆H is multiplied/divided by the same factor