Nowadays, physics has become a very common science. Literally, it is everywhere. The most elementary example: an apple tree grows in your yard, and fruits are ripened on it, the time comes and the apples begin to fall, but in what direction do they fall? Thanks to the law of gravity, our fruit falls to the ground, that is, it goes down, but not up. This was one of the most famous examples of physics, but let's pay attention to thermodynamics, or rather, to phase equilibria, which are no less important in our lives.
Thermodynamics
To get started, let's look at this term. Θερμοδυναμική- this is how the word in Greek looks like. The first part Θερμ means "heat", and the second δυναμική - "strength". Thermodynamics is a branch of physics that studies the properties of a macroscopic system, as well as various ways of converting and transferring energy. In this section, various states and processes are specially studied so that the concept of temperature can be introduced into the description (this is a physical quantity that characterizes the thermodynamic system and is measured using certain instruments). All processes in thermodynamic systems are described only by microscopic values (pressure and temperature, as well as the concentration of components).
Clapeyron-Clausius equation
Every physicist knows this equation, but let us analyze it in parts. It refers to the equilibrium processes of the transition of a certain matter from one phase to another. This is clearly seen in such examples: melting, evaporation, sublimation (one of the methods of preserving products, which takes place by completely removing moisture). The formula clearly shows the ongoing processes:
- n = PV / RT;
- where T is the temperature of the substance;
- P-pressure
- R-specific heat of the phase transition;
- V-change in specific volume.
History of the equation
The Clapeyron-Clausius equation is an excellent mathematical explanation of the second law of thermodynamics. Also referred to as "Clausius inequality." Naturally, the theorem was developed by the scientist himself, who wanted to explain the relationship between the heat flux in the system and entropy, as well as its environment. Clausius developed this equation in his attempts to explain entropy and determine its quantities. In the literal sense, the theorem enables us to determine whether a cyclic process is reversible or irreversible. This inequality offers us a quantitative formula for understanding the second law.
The scientist was one of the first who worked on the idea of entropy, and even gave this process a name. What is now known as the Clausius theorem was first published in 1862 in the sixth work of Rudolph “On the use of the transform equivalence theorem for work in the interior”. The scientist tried to show a proportional relationship between entropy and energy flow by heating (δ Q ) in the system. In construction, this thermal energy can be converted into work, and it can be transformed into heat through a cyclic process. Rudolph proved that "the algebraic sum of all the transformations that take place in a cyclic process can only be less than zero or, in extreme cases, equal to zero."
Closed Insulated System
An isolated system is one of the following methods:
- The physical system is far from others that do not interact with them.
- The thermodynamic system is closed by rigid fixed walls through which neither matter nor energy can pass.
Despite the fact that the subject is internally related to his own gravity, an isolated system is usually taken outside the external gravitational and other distant forces.
This can be contrasted with what (in a more general terminology used in thermodynamics) is called a closed system, covered by selective walls through which energy can be transferred in the form of heat or work, but not matter. And with an open system into which matter and energy enter or exit, although it may have various impenetrable walls in parts of its borders.
An isolated system is subject to the law of conservation. Most often in thermodynamics, matter and energy are considered as separate concepts.
Thermodynamic transitions
To understand quantum phase transitions, it is useful to compare them with classical transformations (also called thermal conversions). CPT describes the cusp in the thermodynamic properties of the system. It signals a particle reorganization. A typical example is the freezing transition of water, which describes the smooth transition between a liquid and a solid. Classical phase outgrowths are due to competition between the energy of the system and the entropy of its thermal fluctuations.
The classical system has no entropy at zero temperature and, therefore, phase transformation cannot occur. Their order is determined by the first discontinuous derivative of the thermodynamic potential. And, of course, it has the first order. The phase transformations from a ferromagnet to a paramagnet are continuous and have a second order. These constant changes from the ordered to the disordered phase are described by the order parameter, which are equal to zero. For the aforementioned ferromagnetic transformation, the order parameter will be the total magnetization of the system.
Gibbs potential
Gibbs free energy is the maximum the number of works without expansion that can be removed from a thermodynamic closed system (which can exchange heat and work with the environment). Such a maximum result can be obtained only in a completely reversible process. When the system is transformed in the opposite way from the first state to the second, the decrease in Gibbs free energy is equal to that performed by the system in its environment, minus the work of pressure forces.
Equilibrium states
Thermodynamic and mechanical equilibrium is an axiomatic concept of thermodynamics. This is the internal state of one or more systems that are connected by more or less permeable or impermeable walls. In this state, there are no pure macroscopic flows from matter or energy, either within a system or between systems.
In its own conception of the state of internal equilibria, macroscopic change does not occur. Systems are simultaneously in mutual thermal, mechanical, chemical (constants), radiation equilibria. They can be in one form. In this process, all species are stored immediately and an infinite amount of time until the physical operation is disrupted. In macroscopic equilibrium, perfectly accurate balanced exchanges occur. The above evidence is a physical explanation of this concept.
The basics
Each laws, theorems, formulas have their own foundations. Let's analyze the 3 basics of the law of phase equilibrium.
- The phase is a form of matter that is homogeneous in chemical composition, physical state and mechanical equilibrium. Typical phases are solid, liquid and gaseous. Two immiscible liquids (or liquid mixtures with different compositions) separated by a separate boundary are considered two different phases and immiscible solid particles.
- The number of components ( C ) is the number of chemically independent components of the system. The minimum number of independent species required to determine the composition of all phases of the system.
- The number of degrees of freedom ( F ) in this context is the number of intensive variables that are independent of each other.
Classification by phase equilibria
- Solid net transfer reactions (often called solid phase reactions) occur between solids of different compositions. They may include elements found in liquids (H, C), but these elements are stored in solid phases, so that liquid phases (H 2 O, CO 2 ) are not involved as reagents or products. Solid net transfer reactions can be continuous or discontinuous, as well as terminal.
- Polymorphic are a special type of solid-state reaction, which includes phases of identical composition. Classical examples are reactions between aluminum silicates, kyanite-sillimanite-andalusite, the conversion of graphite to diamond at high pressure, and the equilibrium of calcium carbonate.
Laws of equilibrium
Gibbs factory rule was proposed by Josiah Willard Gibbs in his famous article entitled “The Equilibrium of Heterogeneous Substances,” which was published from 1875 to 1878. It is applied to non-reactive multicomponent heterogeneous systems in thermodynamic equilibrium and is a given equality:
- F = C-P + 2;
- where F is the number of degrees of freedom;
- C is the number of components;
- P is the number of phases in thermodynamic equilibrium with each other.
The number of degrees of freedom is the number of unoccupied intensive variables. The largest number of thermodynamic parameters, such as temperature or pressure, that can vary simultaneously and arbitrarily, without affecting each other. An example of a one-component system is a system with one pure chemical substance, and two-component systems, such as mixtures of water and ethanol, have two independent components. Typical phase transitions (phase equilibrium) are solids, liquids, gases.
Phase rule at constant pressure
For applications in the field of materials science regarding phase changes between different solid structures, constant pressure often arises (for example, in one atmosphere) and is ignored as a degree of freedom, therefore the rule becomes: F = C - P + 1.
This formula is sometimes called “condensed-phase rule”, but, as we know, it does not apply to these systems, which are subject to high pressures (for example, in geology), since the consequences of these pressures can cause catastrophic consequences.
It may seem that phase equilibrium is just an empty phrase, and there are few physical processes involved in this moment, but, as we saw, many laws we know do not work without it, so you need to get a little familiar with these unique, colorful, albeit slightly a little boring rules. This knowledge has helped many people. They learned to apply them to themselves, for example, electricians, knowing the rules of working with phases, can protect themselves from unnecessary danger.