Important Vocab

• Thermochemistry: the study of the relationships between chemistry and energy

• Energy: the capacity to do work

• Work: the result of a force acting through a distance

• Thermodynamics: study of energy and its interconversions

• Internal energy: the sum of the kinetic and potential energies of all of the particles that compose the system

• Pressure–volume work: occurs when a force (caused by a change in volume) acts through a distance against an external pressure

• Heat: the transfer of thermal energy

• Thermal equilibrium: Surroundings & object are same temp, no additional net transfer of temp

• Heat capacity: the quantity of heat required to change its temperature by 1 °C

• Calorimetry: measure the thermal energy exchanged between the reaction (defined as the system) and the surroundings by observing the change in temperature of the surroundings

• Enthalpy: the sum of a system’s internal energy and the product of its pressure and volume

Types of Energy

Kinetic

• Associated with the motion of an object

• Ex: Moving ball

Thermal

• Associated with the temperature of an object

• Type of kinetic energy

• Arises from the motions of atoms or molecules within a substance

• Ex: Hot cup of coffee

Potential

• Associated with the position or composition of an object

• Ex: compressed spring, ball held up above the ground

Chemical

• Type of potential energy

• Often stored in chemical bonds

• Associated with the relative positions of electrons and nuclei in atoms and molecules

Thermodynamics

First Law of Thermodynamics

• Also known as the law of energy conservation

• Energy is neither created nor destroyed

• Internal energy: the sum of the kinetic and potential energies of all of the particles that compose the system

  • Internal energy is a state system (value depends only on the state of the system)

• Energy flow rules:

  • Reactants have a higher internal energy than the products, is negative and energy flows out of the system into the surroundings

  • If the reactants have a lower internal energy than the products, is positive and energy flows into the system from the surroundings

Heat

• Heat: the transfer of thermal energy

• Thermal equilibrium: Surroundings & object are same temp, no additional net transfer of temp

• Heat capacity: the quantity of heat required to change its temperature by 1 °C

  • Depends on:

    • The amount of matter being heated

    • Specific heat capacity/molar capacity (q)

• Pressure–volume work: occurs when a force (caused by a change in volume) acts through a distance against an external pressure

• w = F * D

• Calorimetry: measure the thermal energy exchanged between the reaction (defined as the system) and the surroundings by observing the change in temperature of the surroundings

  • Measurement tool: bomb calorimeter and coffee-cup calorimeter

    • Bomb calorimetry occurs at constant volume and measures ΔE for a reaction

    • Coffee-cup calorimetry occurs at constant pressure and measures ΔH for a reaction

Enthalpy

• Enthalpy: the sum of a system’s internal energy and the product of its pressure and volume

• H = E + PV

• Negative delta H = endothermic reaction

• Positive delta H = exothermic reaction

• The value of ΔH for a chemical reaction is the amount of heat absorbed or evolved in the reaction under conditions of constant pressure

• An endothermic reaction has a positive ΔH and absorbs heat from the surroundings. An endothermic reaction feels cold to the touch

• An exothermic reaction has a negative ΔH and gives off heat to the surroundings. An exothermic reaction feels warm to the touch

• Standard heat of formation

• Standard State

  • For a Gas: The standard state for a gas is the pure gas at a pressure of exactly 1 atm.

  • For a Liquid or Solid: The standard state for a liquid or solid is the pure substance in its most stable form at a pressure of 1 atm and at the temperature of interest (often taken to be 25 °C).

  • For a Substance in Solution: The standard state for a substance in solution is a concentration of exactly 1 M.

• Standard Enthalpy Change (ΔH°)

  • The change in enthalpy for a process when all reactants and products are in their standard states. The degree sign indicates standard states.

• Standard Enthalpy of Formation ()

  • For a Pure Compound: The change in enthalpy when 1 mol of the compound forms from its constituent elements in their standard states.

  • For a Pure Element in Its Standard State: delta H = 0

Hess’ Law

• If a chemical equation can be expressed as the sum of a series of steps, then for the overall equation is the sum of the heats of reaction for each step

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