RVRH

Phase (matter)

In the physical sciences, a phase is a set of states of a macroscopic physical system that have relatively uniform chemical composition and physical properties (i.e. density, crystal structure, index of refraction, and so forth).

Phases are sometimes confused with states of matter, but there are significant differences. States of matter refers to the differences between gases, liquids, solids, etc. If there are two regions in a chemical system that are in different states of matter, then they must be different phases. However, the reverse is not true — a system can have multiple phases which are in equilibrium with each other and also in the same state of matter. For example, diamond and graphite are both solids but they are different phases, even though their composition may be identical. A system with oil and water at room temperature will be two different phases of differing composition, but both will be the liquid state of matter. This difference is especially important when considering the Gibbs’ phase rule, which governs the number of allowed phases.

In general, two different states of a system are in different phases if there is an abrupt change in their physical properties while transforming from one state to the other. Conversely, two states are in the same phase if they can be transformed into one another without any abrupt changes. There are, however, exceptions to this statement — for example the liquid-gas critical point discussed below in the Phase Diagrams section.

An important point is that different types of phases are associated with different physical qualities. When discussing the solid, liquid, and gaseous phases, we talked about rigidity and compressibility, and the effects of varying the pressure and volume, because those are the relevant properties that distinguish a solid, a liquid, and a gas. On the other hand, when discussing paramagnetism and ferromagnetism, we look at the magnetization, because that is what distinguishes the ferromagnetic phase from the paramagnetic phase. Several more examples of phases will be given in the following section.

In more technical language, a phase is a region in the parameter space of thermodynamic variables in which the free energy is analytic; between such regions there are abrupt changes in the properties of the system, which correspond to discontinuities in the derivatives of the free energy function. As long as the free energy is analytic, all thermodynamic properties (such as entropy, heat capacity, magnetization, and compressibility) will be well-behaved, because they can be expressed in terms of the free energy and its derivatives. For example, the entropy is the first derivative of the free energy with temperature.

When a system goes from one phase to another, there will generally be a stage where the free energy is non-analytic. This is a phase transition. Due to this non-analyticity, the free energies on either side of the transition are two different functions, so one or more thermodynamic properties will behave very differently after the transition. The property most commonly examined in this context is the heat capacity. During a transition, the heat capacity may become infinite, jump abruptly to a different value, or exhibit a “kink” or discontinuity in its derivative. See also differential scanning calorimetry.

The distribution of kinetic energy among molecules is not uniform, and it changes randomly. This means that at, say, the surface of a liquid, there may be an individual molecule with enough kinetic energy to jump into the gas phase. Likewise, individual gas molecules may have low enough kinetic energy to join other molecules in the liquid phase. This phenomenon means that at any given temperature and pressure, multiple phases may co-exist.

For example, under standard conditions for temperature and pressure, a bowl of liquid water in dry air will evaporate until the partial pressure of gaseous water equals the vapor pressure of water. At this point, the rate of molecules leaving and entering the liquid phase becomes the same (due to the increased number of gaseous water molecules available to re-condense). The fact that liquid molecules with above-average kinetic energy have been removed from the bowl results in evaporative cooling. Similar processes may occur on other types of phase boundaries.

Gibbs’ phase rule relates the number of possible phases, variables such as temperature and pressure, and whether or not an equilibrium will be reached.

States of matter

States of matter refers to the differences between gases, liquids, solids, etc

In addition to the specific chemical properties that distinguish different chemical classifications chemicals can exist in several phases. For the most part, the chemical classifications are independent of these bulk phase classifications; however, some more exotic phases are incompatible with certain chemical properties. A phase is a set of states of a chemical system that have similar bulk structural properties, over a range of conditions, such as pressure or temperature. Physical properties, such as density and refractive index tend to fall within values characteristic of the phase. The phase of matter is defined by the phase transition, which is when energy put into or taken out of the system goes into rearranging the structure of the system, instead of changing the bulk conditions.

Sometimes the distinction between phases can be continuous instead of having a discrete boundary, in this case the matter is considered to be in a supercritical state. When three states meet based on the conditions, it is known as a triple point and since this is invariant, it is a convenient way to define a set of conditions.

The most familiar examples of phases are solids, liquids, and gases. Many substances exhibit multiple solid phases. For example, there are three phases of solid iron (alpha, gamma, and delta) that vary based on temperature and pressure. A principle difference between solid phases is the crystal structure, or arrangement, of the atoms. Less familiar phases include plasmas, Bose-Einstein condensates and fermionic condensates and the paramagnetic and ferromagnetic phases of magnetic materials. While most familiar phases deal with three-dimensional systems, it is also possible to define analogs in two-dimensional systems, which has received attention for its relevance to systems in biology.

Molecules

A molecule is the smallest indivisible portion of a pure compound or element that retains a set of unique chemical properties. Molecules differ from other chemical entities in that they can and often do exist as single electrically neutral units. Salts, for example, do not consist of molecular units but rather of many cations and anions in a crystal lattice. Molecules are typically a set of atoms bound together by covalent bonds, such that the structure is electrically neutral and all valence electrons are paired with other electrons either in bonds or in lone pairs.

Substance

A chemical substance is a general term that can be an element, compound or a mixture of compounds, elements or compounds and elements. Most of the matter we encounter in our daily life are one or another kind of mixtures, e.g. air, alloys, biomass etc.

Compounds

A compound is a substance with a fixed ratio of chemical elements which determines the composition, and a particular organization which determines chemical properties. For example, water is a compound containing hydrogen and oxygen in the ratio of two to one, with the oxygen between the hydrogens, and an angle of 104.5° between them. Compounds are formed and interconverted by chemical reactions.

Elements

A chemical element, or element for short, is a type of atom that is defined by its atomic number; that is, by the number of protons in its nucleus. The term is also used to refer to a pure chemical substance composed of atoms with the same number of protons.[1]

Common examples of elements are hydrogen, nitrogen, and carbon. In total, 117 elements have been observed as of 2007, of which 94 occur naturally on Earth. Elements with atomic numbers greater than 82 (i.e,. bismuth and those above), are inherently unstable and undergo radioactive decay. In addition, elements 43 and 61 (technetium and promethium) have no stable isotopes, and also decay. However, the unstable elements up to atomic number 94 with no stable nuclei are found in nature as a result of the natural decay processes of uranium and thorium.[2]

All chemical matter consists of these elements. New elements are discovered from time to time through artificial nuclear reactions.

An element is a class of atoms which have the same number of protons in the nucleus. This number is known as the atomic number of the element. For example, all atoms with 6 protons in their nuclei are atoms of the chemical element carbon, and all atoms with 92 protons in their nuclei are atoms of the element uranium.

The most convenient presentation of the chemical elements is in the periodic table of the chemical elements, which groups elements by atomic number. Due to its ingenious arrangement, groups, or columns, and periods, or rows, of elements in the table either share several chemical properties, or follow a certain trend in characteristics such as atomic radius, electronegativity, etc. Lists of the elements by name, by symbol, and by atomic number are also available. In addition, several isotopes of an element may exist.

Atoms

In chemistry and physics, an atom (Greek ἄτομος or átomos meaning “indivisible”) is still characterized as a chemical element.[2] (átomos is usually translated as “indivisible” or “uncuttable.” Until the advent of quantum mechanics, dividing a material object was invariably equated with cutting it.) Whereas the word atom originally denoted a particle that cannot be cut into smaller particles, the atoms of modern parlance are composed of subatomic particles:

electrons, which have a negative charge, a size which is so small as to be currently unmeasurable, and which are the least heavy (i.e., massive) of the three;
protons, which have a positive charge, and are about 1836 times more massive than electrons; and
neutrons, which have no charge, and are the same size as protons.
Protons and neutrons make up a dense, massive atomic nucleus, and are collectively called nucleons. The electrons form the much larger electron cloud surrounding the nucleus.

Atoms can differ in the number of each of the subatomic particles they contain. Atoms of the same element have the same number of protons (called the atomic number). Within a single element, the number of neutrons may vary, determining the isotope of that element. The number of electrons associated with an atom is most easily changed, due to the lower energy of binding of electrons. The number of protons (and neutrons) in the atomic nucleus may also change, via nuclear fusion, nuclear fission, bombardment by high energy subatomic particles or photons, or certain (but not all) types of radioactive decay. In such processes which change the number of protons in a nucleus, the atom becomes an atom of a different chemical element.

Atoms are electrically neutral if they have an equal number of protons and electrons. Atoms which have either a deficit or a surplus of electrons are called ions. Electrons that are furthest from the nucleus may be transferred to other nearby atoms or shared between atoms. By this mechanism atoms are able to bond into molecules and other types of chemical compounds like ionic and covalent network crystals.

Atoms are the fundamental building blocks of chemistry, and are conserved in chemical reactions.

http://en.wikipedia.org/wiki/Atom

Chemistry

Chemistry (from Egyptian kēme (chem), meaning “earth”[1]) is the science treating matter at the atomic to macromolecular scale, the reactions, transformations and aggregations of matter, as well as accompanying energy and entropy changes during these processes. In short, chemistry studies molecules, crystals, and metal/nonmetals and is concerned with the composition and statistical properties of such structures, as well as their transformations and interactions to become materials encountered in everyday life. According to quantum mechanics, all physical and chemical properties of materials are generally determined by their structure at the molecular or atomic scale, which is itself defined by interatomic electromagnetic forces, and laws of quantum mechanics. Robert Boyle (1661), Antoine Lavoisier (1787), and John Dalton (1808) can be considered the three fathers of modern chemistry,[2] while some consider the earlier chemist Geber (d. 815) to be the “father of chemistry”.[3]

The word chemistry comes from the earlier study of alchemy, which is basically the quest to make gold from earthen starting materials.[19] As to the origin of the word “alchemy” the question is a debatable one; it certainly can be traced back to the Greeks, and some, following E. Wallis Budge, have also asserted Egyptian origins. Alchemy, generally, derives from the old French alkemie from the Arabic al-kimia - “the art of transformation”. The Arabs borrowed the word “kimia” from the Greeks when they conquered Alexandria in the year 642 AD. A tentative outline is as follows:

• Egyptian alchemy [5,000 BC – 400 BC], formulate early “element” theories such as the Ogdoad.
• Greek alchemy [332 BC – 642 AD], the Greek king Alexander the Great conquers Egypt and founds
• Alexandria, having the world’s largest library, where scholars and “wise” men gather to study.
• Arabian alchemy [642 AD – 1200], the Arabs take over Alexandria; Jabir is the main chemist
• European alchemy [1300 – present], Pseudo-Geber builds on Arabic chemistry
• Chemistry [1661], Boyle writes his classic chemistry text The Sceptical Chymist
• Chemistry [1787], Lavoisier writes his classic Elements of Chemistry
• Chemistry [1803], Dalton publishes his Atomic Theory
• Thus, an alchemist was called a ‘chemist’ in popular speech, and later the suffix “-ry” was added to this to describe the art of the chemist as “chemistry”.