BackThermochemistry: Energy, Internal Energy, and Thermodynamic Systems
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Thermochemistry and Thermodynamics
Introduction to Energy and Thermochemistry
Thermochemistry is a branch of thermodynamics that focuses on the study of energy changes, particularly heat, that accompany chemical reactions. Understanding how energy is transferred and transformed is fundamental to predicting the behavior of chemical systems.
Energy is defined as the ability to do work or transfer heat.
Thermodynamics is the study of energy and its transformations.
Thermochemistry specifically examines the energy changes associated with chemical reactions, especially those involving heat.
Types of Systems in Thermodynamics
Classification of Systems
In thermodynamics, a system is the part of the universe we are studying, while the surroundings are everything else. Systems are classified based on their ability to exchange energy and matter with their surroundings.
Open System: Can exchange both matter and energy (usually as heat) with the surroundings.
Closed System: Can exchange energy but not matter with the surroundings.
Isolated System: Cannot exchange either matter or energy with the surroundings.
Example: A boiling pot of water without a lid is an open system; with a tightly sealed lid, it is a closed system; a thermos bottle is an example of an isolated system.
Internal Energy and the First Law of Thermodynamics
Definition of Internal Energy
The internal energy (E) of a system is the sum of all kinetic and potential energies of its components. The change in internal energy, , is a key concept in thermodynamics.
Change in Internal Energy:
If , the system has absorbed energy from the surroundings.
If , the system has released energy to the surroundings.
Example: When hydrogen and oxygen react to form water, energy is released to the surroundings, so .
Thermodynamic Quantities
Every thermodynamic quantity has three parts:
A number (magnitude)
A unit (e.g., Joules, J)
A sign (indicating direction of energy flow)
Sign Conventions:
A positive means the system gains energy from the surroundings.
A negative means the system loses energy to the surroundings.
First Law of Thermodynamics
The First Law of Thermodynamics states that energy cannot be created or destroyed, only transferred or transformed. For a system, the change in internal energy is given by:
Where q is heat and w is work.
Sign Conventions for q and w:
q > 0: Heat absorbed by the system (endothermic)
q < 0: Heat released by the system (exothermic)
w > 0: Work done on the system
w < 0: Work done by the system
Example: If a gas expands against a piston, it does work on the surroundings, so w is negative.
State Functions
Definition and Importance
A state function is a property whose value depends only on the current state of the system, not on the path taken to reach that state. Internal energy (E) is a state function, but heat (q) and work (w) are not.
State functions: Internal energy (E), enthalpy (H), pressure (P), volume (V), temperature (T)
Path functions: Heat (q), work (w)
Example: The change in internal energy when a battery is discharged is the same regardless of whether it is used to power a fan or shorted out directly.
