Thermodynamics: Energy, Heat, and Matter

• Thermodynamics deals with temperature, heat, energy, radiation, work, and properties of matter.
• Energy can be electrical, mechanical, or chemical.
• William Thomson coined the term thermodynamics in 1749.
• Thermodynamics is derived from the Greek words "thermes" meaning heat, and "dynamikos" meaning powerful.
• Thermodynamics is the branch of science concerned with the relationship between forms of energy and heat.
• Thermodynamics is the branch of science dealing with heat, temperature, energy, work, radiation and properties of matter.
• Thermodynamics explains how thermal energy is converted to/from other energies and how matter is affected by this process.
• Thermal energy is the energy coming from heat and is generated by the movement of particles within an object.
• It is imperative to define a system and its surroundings because that concept becomes the basis for many types of descriptions and calculations in thermodynamics.
• A goal of thermochemistry is determining heat quantity exchanged between a system and its surroundings.
• The system is the part of the universe being studied.
• Surroundings are the rest of the universe that interacts with the system.
• Systems and surroundings can range in scale, such as rain forests or beakers. The type of system dictates the laws of thermodynamics in chemistry.
• Focusing attention on the chemical reaction defines the system in chemistry (e.g. 2H₂ + O₂ -> 2H₂O).
• A system consists of the molecules that react.
• Surroundings are everything else in the universe.
• An open system freely exchanges both energy and matter with its surroundings.
• An example of an open system is boiling soup in a saucepan; energy and matter transfer to the surroundings through steam..
• A saucepan exemplifies an open system, allowing matter transfer (spices) and energy transfer (heating, steam leaving).
• Putting a lid on a saucepan makes it a closed system because this only exchanges energy with its surroundings, not matter.
• Putting a lid on a beaker makes it a closed system.
• Even though a closed system can't transfer matter, it can still transfer energy.
• A thermos is an example of an isolated system that does not allow heat or matter transfer.
• An isolated system does not exchange energy or matter with its surroundings, as exemplified by soup poured into an insulated container.
• Creating a completely isolated system is difficult.

First Law of Thermodynamics

• The first law of thermodynamics is also known as the conservation of energy principle.
• Energy can neither be created nor destroyed, but can be changed from one form to another.

First Law of Thermodynamics Examples

• Fans convert electrical energy to mechanical energy.
• Plants convert the radiant energy of sunlight to chemical energy through photosynthesis.
• Eating plants converts chemical energy into kinetic energy when we swim, walk, breathe, and scroll.

Second Law of Thermodynamics

• Energy in the form of heat only flows from regions of higher temperature to that of lower temperature.
• The second law has a big impact; it costs money to run air conditioning.
• Human bodies obey the second law of thermodynamics.

Second Law of Thermodynamics Examples

• Sweating in a crowded room exemplifies the second law, the body uses sweat to cool itself, transferring heat to the surroundings.

Third Law of Thermodynamics

• The third law states that the entropy of a perfect crystal is zero when the crystal's temperature equals absolute zero (0 K).
• Entropy is "waste energy".
• There is no heat energy at absolute zero, so there is no waste energy.

Third Law Of Thermodynamics Examples

• Steam has high entropy because the molecules within move freely at higher temperatures.
• As temperature decreases below 100°C, steam converts to water, restricting molecule movement and decreasing water's entropy.
• When water cools below 0°C, it becomes solid ice, further restricting molecule movement and reducing system entropy.
• As ice temperature lowers, molecules are more restricted, reducing the substance's entropy.
• In reality, cooling any substance to absolute zero is impossible (therefor the entropy should be zero)

Laws of Thermodynamics

• The laws of thermodynamics define the physical quantities of energy, temperature, and entropy that characterize thermodynamic systems at thermal equilibrium.
• The laws represent how the quantities behave under various circumstances.
• There are four laws of thermodynamics: Zeroth, First, Second, and Third.

Zeroth Law of Thermodynamics

• The Zeroth Law is the basis for temperature measurement.
• Two bodies in thermal equilibrium with a third body are in thermal equilibrium with each other.
• Two cups, A and B, with boiling water illustrate the Zeroth Law.
• A thermometer placed in cup A reads 100°C.
• The thermometer is in equilibrium with cup A when it reads 100°C.
• Moving the thermometer to cup B will result in readings of 100°C again.

Terms in Thermodynamics

• An endothermic process absorbs heat.
• An exothermic process loses heat.
• A state function is the property of a system that is determined by specifying the system's condition like temperature and pressure.
• The value of a state function depends only on the present state and not the path the system took to attain the state.
• Enthalpy (H) accounts for heat flow in processes occurring at constant pressure.
• A positive ΔH implies the system gained heat.
• A negative ΔH means the system lost heat to the surroundings.
• The enthalpy change that accompanies a reaction is called the enthalpy of reaction.
• A negative ΔH value identifies the reaction as exothermic.
• A positive ΔH value means the reaction is endothermic.
• Enthalpy is an extensive property.
• The magnitude of ΔH is directly proportional to the reactant amount consumed in the process.
• The enthalpy for a reaction is equal in magnitude, but opposite in sign, to ΔH for the reverse reaction.
• The enthalpy change for a reaction depends on the state of the reactants and products.
• A thermodynamic process is a passage of a thermodynamic system from an initial to a final state of thermodynamic equilibrium.
• Classical Thermodynamics analyzes the behavior of matter with a macroscopic approach. Units, properties such as temperature and pressure are taken into consideration, to help calculate characteristics of the matter/individuals that are undergoing the process.

Thermodynamics

• Internal energy (U)
• Enthalpy (H)
• Entropy (S)
• Gibb's free energy (G)

Some Terms in Thermodynamics

• The change in internal energy ΔU is the difference between U final and U initial.
• A positive ΔU indicates the system as gained energy from its surrounding.
• A negative ΔU indicates that the system as lost energy to its surroundings
• ΔU = Q + W
• Q is the heat absorbed or evolved by the system.
• U increases when work is done on a system or heat is added to a system.

More Terms

• +Q = The system gains heat.
• -Q = The system loses heat.
• +W = Work is done on the system.
• -W = Work is done by the system.
• +ΔU = Net gain fo energy by the system.
• -ΔU = Net loss of energy by the system.
• Energy is the ability to do work. Work occurs when an object moves against a force.
• W = FD:
• W represents work.
• F represents force.
• D represents distance.
• When a force is applied, work is done (e.g. gravity).

Thermodynamics explores the relationships between energy forms and heat, detailing how thermal energy converts to other energies and affects matter. Thermal energy, arising from particle movement, necessitates defining systems and surroundings for accurate descriptions and calculations in thermodynamics.