HiSET Science: Laws of Thermodynamics

Laws of Thermodynamics

First law
The first law of thermodynamics states that energy cannot be created or destroyed, but only converted from one form to another. It is generally applied as Q = ΔU + W, where Q is the net heat energy added to a system, ΔU is the change in internal energy of the system, and W is the work done by the system. For any input of heat energy to a system, that energy must be either converted to internal energy through a temperature increase or expended in doing work. For a system that gives off heat, either the temperature of the system must decrease or work must be done on the system by its surroundings.

By convention, work done by the system is positive while work done on the system is negative.
For instance, suppose a gas is compressed by a piston while the gas temperature remains constant. If we consider the gas to be the system, the work is negative, since the work is being performed on the gas. Since the temperature remains constant, ΔU = 0. Thus Q must be a negative quantity, indicating that heat is lost by the gas. Conversely, if the gas does positive work on the piston while remaining at a constant temperature, the gas must be receiving heat input from the surroundings.

Second law
The second law of thermodynamics is primarily a statement of the natural tendency of all things toward disorder rather than order. It deals with a quantity called entropy, which is an inverse measure of the remaining useful energy in a system. If we take a system of a pot of hot water and an ice cube, the system entropy initially has a value of s1.
After the ice cube melts in the water and the system reaches an equilibrium temperature, the system has larger entropy value s2, which is the maximum entropy for the system. The system cannot return to its initial state without work input to refreeze the ice cube and reheat the water. If this is done and the system returns to a state with entropy s1, then the entropy of the surroundings must at the same time increase by more than s2 – s1, since the net entropy from any process is always greater than zero.

Reversible processes are those that may be accomplished in reverse without requiring additional work input. These processes do not exist in the real world, but can be useful for approximating some situations. All real processes are irreversible, meaning they require additional work input to accomplish in reverse. Another important concept is that of spontaneity, the ability of a process to occur without instigation. An ice cube located in an environment at a temperature above the freezing point will spontaneously melt. Although some processes can decrease system entropy at a cost to the entropy of the surroundings, all spontaneous processes involve an increase in system entropy.

Third and Zeroth Laws
The third law of thermodynamics regards the behavior of systems as they approach absolute zero temperature. Actually reaching a state of absolute zero is impossible. According to this law, all activity disappears as molecules slow to a standstill near absolute zero, and the system achieves a perfect crystal structure while the system entropy approaches its minimum value. For most systems, this would in fact be a value of zero entropy.
Note that this does not violate the second law since causing a system to approach absolute zero would require an immense increase in the entropy of the surroundings, resulting in a positive net entropy. This law is used to determine the value of a material's standard entropy, which is its entropy value at the standard temperature of 25 °C.
The zeroth law of thermodynamics deals with thermal equilibrium between two systems. It states that if two systems are both in thermal equilibrium with a third system, then they are in thermal equilibrium with each other. This may seem intuitive, but it is an important basis for the other thermodynamic laws.

Entropy
Entropy (S) is the amount of disorder or randomness of a system. According to the second law of thermodynamics, systems tend toward a state of greater entropy. The second law of thermodynamics can also be stated as ΔS > 0. Processes with positive changes in entropy tend to be spontaneous.

For example, melting is a process with a positive ΔS.
When a solid changes into a liquid state, the substance becomes more disordered; therefore, entropy increases. Entropy also will increase in a reaction in which the number of moles of gases increases due to the amount of disorder increasing. Entropy increases when a solute dissolves into a solvent due to the increase in the number of particles. Entropy increases when a system is heated due to the particles moving faster and the amount of disorder increasing.

Spontaneous / Reversible Processes
Some chemical processes are spontaneous. According to the second law of thermodynamics, systems or processes always tend to a state of greater entropy or lower potential energy. Some exothermic chemical systems are spontaneous because they can increase their stability by reaching a lower potential energy. If processes or reactions have products at a lower potential energy, these processes tend to be spontaneous. Spontaneous reactions have only one direction as given by the second law of thermodynamics.
Spontaneous processes go in the direction of greater entropy and lower potential energy. To be reversible, a reaction or process has to be able to go back and forth between two states. A spontaneous process is irreversible.

Concept of Change in Enthalpy
All chemical processes involve either the release or the absorption of heat. Enthalpy is this heat energy.

Enthalpy is a state function that is equivalent to the amount of heat a system exchanges with its surroundings.

For exothermic processes, which release heat, the change in enthalpy (ΔH) is negative because the final enthalpy is less than the initial enthalpy. For endothermic processes, which absorb heat, the change in enthalpy (ΔH) is positive because the final enthalpy is greater than the initial enthalpy.




Gibbs Energy
Gibbs energy (G), also known as Gibbs free energy, is the energy of the system that is available to do work. Gibbs energy determines the spontaneity of chemical and physical processes. Some processes are spontaneous because ΔH < 0 or because ΔS > 0.

If one of the conditions is favorable but the other condition is not favorable, Gibbs energy can be used to determine if a process is spontaneous. Gibbs energy is given by G = H – TS.

For processes that occur at constant temperature, ΔG= ΔH – TΔS. If ΔG is equal to zero, then the reaction is at equilibrium and neither the forward nor the reverse reaction is spontaneous. If ΔG is less than zero, then the forward reaction is spontaneous. If ΔG is greater than zero, then the reverse reaction is spontaneous.