The total entropy of a system either increases or remains constant in any process; it never decreases. For example, heat transfer cannot occur spontaneously from cold to hot, because entropy would decrease. Entropy is very different from energy. Entropy is not conserved but increases in all real processes.

A campfire is an example of entropy. The solid wood burns and becomes ash, smoke and gases, all of which spread energy outwards more easily than the solid fuel. Ice melting, salt or sugar dissolving, making popcorn and boiling water for tea are processes with increasing entropy in your kitchen.

The entropy only goes to zero if the system is definitely in a single quantum state, since log(1)=0. That turns out to basically just be a definition of T, with the understanding that dQ is the heat flow into the system as it stays in thermal equilibrium. At low T S grows a lot as heat flows in, at high T less so.

Entropy changes (ΔS) are estimated through relation ΔG=ΔH−TΔS for finite variations at constant T.

Entropy is a measure of randomness or disorder of the system. The greater the randomness, the higher the entropy. It is state function and extensive property. Its unit is JK−1mol−1.

And when the change of internal energy equals 0, q=-w. and since Delta S=q/T, you can plug in the equation we just derived in for q. q=nRT*ln(V2/V1). So, Delta S=(nRT*ln(V2/V1))/T.

joule/K

The symbol for entropy is S and the standard entropy of a substance is given by the symbol So, indicating that the standard entropy is determined under standard conditions. The units for entropy are J/K⋅mol.

As long as a system has the same number of atoms and the same number of quanta of energy to share between them, it is plausible that the system possesses a minimum number of possible microstates—and a minimum entropy. …

Unlike energy functions, entropy is not conserved in natural process or in isolated systems.

Entropy is not a conserved quantity. It can be created. In every irreversible cyclic process it gets created. However, entropy cannot be destroyed.

The trivial, everyday phenomenon of something cooling down is the prototypical example of entropy decreasing. It’s as simple as that. Now of course, as the tea cooled, the room warmed. But because at all times the tea was warmer than the room, the room gained more entropy than the tea lost.

Entropy increases when a substance is broken up into multiple parts. The process of dissolving increases entropy because the solute particles become separated from one another when a solution is formed. Entropy increases as temperature increases.

Refrigerator compressors and chemical hand warmers are both real-life examples of enthalpy. Both the vaporization of refrigerants in the compressor and the reaction to the iron oxidation in a hand warmer generate a change in heat content under constant pressure.

Enthalpy, the sum of the internal energy and the product of the pressure and volume of a thermodynamic system. In symbols, the enthalpy, H, equals the sum of the internal energy, E, and the product of the pressure, P, and volume, V, of the system: H = E + PV.

What Is the Importance of Enthalpy? Measuring the change in enthalpy allows us to determine whether a reaction was endothermic (absorbed heat, positive change in enthalpy) or exothermic (released heat, a negative change in enthalpy.) It is used to calculate the heat of reaction of a chemical process.

At constant pressure, the heat of reaction is equal to the enthalpy change of the system. Most chemical reactions occur at constant pressure, so enthalpy is more often used to measure heats of reaction than internal energy.

Enthalpy is a state function. If there is no non-expansion work on the system and the pressure is still constant, then the change in enthalpy will equal the heat consumed or released by the system (q). ΔH=q. This relationship can help to determine whether a reaction is endothermic or exothermic.

Heat is a transfer of energy due to a temperature difference. Enthalpy is the change in amount of heat in a system at constant pressure. You can only use heat and enthalpy interchangeably if there is no work being done to the system.

Q is the energy transfer due to thermal reactions such as heating water, cooking, etc. anywhere where there is a heat transfer. You can say that Q (Heat) is energy in transit. Enthalpy (Delta H), on the other hand, is the state of the system, the total heat content.

The quantitative relationship between heat transfer and temperature change contains all three factors: Q=mcΔT Q = mc Δ T , where Q is the symbol for heat transfer, m is the mass of the substance, and ΔT is the change in temperature. The symbol c stands for specific heat and depends on the material and phase.

Enthalpy changes Enthalpy

Every chemical reaction involves a change in free energy, called delta G (∆G). The change in free energy can be calculated for any system that undergoes a change, such as a chemical reaction. To calculate ∆G, subtract the amount of energy lost to entropy (denoted as ∆S) from the total energy change of the system.

Reactions with a negative ∆G release energy, which means that they can proceed without an energy input (are spontaneous). In contrast, reactions with a positive ∆G need an input of energy in order to take place (are non-spontaneous).