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States of matter are the distinct forms that different phases of matter take on. Historically, the distinction is made based on qualitative differences in bulk properties. Solid is the state in which matter maintains a fixed volume and shape; liquid is the state in which matter maintains a fixed volume but adapts to the shape of its container; and gas is the state in which matter expands to occupy whatever volume is available.

This diagram shows the nomenclature for the different phase transitions.

More recently, distinctions between states have been based on differences in molecular interrelationships. Solid is the state in which intermolecular attractions keep the molecules in fixed spatial relationships. Liquid is the state in which intermolecular attractions keep molecules in proximity, but do not keep the molecules in fixed relationships. Gas is that state in which the molecules are comparatively separated and intermolecular attractions have relatively little effect on their respective motions. Plasma is a highly ionized gas that occurs at high temperatures. The intermolecular forces created by ionic attractions and repulsions give these compositions distinct properties, for which reason plasma is described as a fourth state of matter.[1][2]

Forms of matter that are not composed of molecules and are organized by different forces can also be considered different states of matter. Fermionic condensate and the quark–gluon plasma are examples.

States of matter may also be defined in terms of phase transitions. A phase transition indicates a change in structure and can be recognized by an abrupt change in properties. By this definition, a distinct state of matter is any set of states distinguished from any other set of states by a phase transition. Water can be said to have several distinct solid states.[3] The appearance of superconductivity is associated with a phase transition, so there are superconductive states. Likewise, liquid crystal states and ferromagnetic states are demarcated by phase transitions and have distinctive properties.

Because solids have thermal energy or heat capacity, their atoms vibrate about fixed mean positions within the ordered (or disordered) lattice. Shown here are the one-dimensional normal modes of vibration in a crystalline solid. The amplitude of the motion has been exaggerated, and is actually much smaller than the lattice parameter. The entire spectrum of lattice vibrations in a crystalline or glassy network plays a key role in the kinetic theory of solids.


Classical states



The particles (ions, atoms or molecules) are packed closely together. The forces between particles are strong enough so that the particles cannot move freely but can only vibrate. As a result, a solid has a stable, definite shape, and a definite volume. Solids can only change their shape by force, as when broken or cut.

A crystalline solid: atomic resolution image of strontium titanate. Brighter atoms are Sr and darker ones are Ti.

In crystalline solids, the particles (atoms, molecules, or ions) are arranged in an ordered three-dimensional structure. There are many different crystal structures, and the same substance can have more than one structure (or solid phase). For example, iron has a body-centred cubic structure at temperatures below 912 °C, and a face-centred cubic structure between 912 and 1394 °C. Ice has fifteen known crystal structures, or fifteen solid phases which exist at various temperatures and pressures.[4]

Solids can be transformed into liquids by melting, and liquids can be transformed into solids by freezing. Solids can also change directly into gases through the process of sublimation.


Structure of a classical monatomic liquid. Atoms have many nearest neighbors in contact, yet no long-range order is present.

The volume is definite if the temperature and pressure are constant. When a solid is heated above its melting point, it becomes liquid. Intermolecular (or interatomic or interionic) forces are still important, but the molecules have enough energy to move relative to each other and the structure is mobile. This means that the shape of a liquid is not definite but is determined by its container. The volume is usually greater than that of the corresponding solid, the most well known exception being water, H2O. The highest temperature at which a given liquid can exist is its critical temperature.[5]


In a gas, the molecules have enough kinetic energy so that the effect of intermolecular forces is small (or zero for an ideal gas), and the typical distance between neighboring molecules is much greater than the molecular size. A gas has no definite shape or volume, but occupies the entire container in which it is confined. A liquid may be converted to a gas by heating at constant pressure to the boiling point, or else by reducing the pressure at constant temperature.

At temperatures below its critical temperature, a gas is also called a vapor, and can be liquefied by compression alone without cooling. A vapor can exist in equilibrium with a liquid (or solid), in which case the gas pressure equals the vapor pressure of the liquid (or solid).

A supercritical fluid (SCF) is a gas whose temperature and pressure are above the critical temperature and critical pressure respectively. It has the physical properties of a gas, but its high density confers solvent properties in some cases which lead to useful applications. For example, supercritical carbon dioxide is used to extract caffeine in the manufacture of decaffeinated coffee.[6]

Non-classical states

Crystalline vs. glassy

Schematic representation of a random-network glassy form (top) and ordered crystalline lattice (bottom) of identical chemical composition.

In crystalline solids, the atoms or molecules that compose the solid are packed closely together. In mineralogy and crystallography, a crystal structure is a unique arrangement of atoms in a crystal. A specific symmetry or crystal structure is composed of a Bravais lattice which is typically represented by a single unit cell. The unit cell is periodically repeated in three dimensions on a lattice.

Non-crystalline or amorphous or glassy solids are often referred to as supercooled liquids, but possess the mechanical properties of both a solid and a liquid, depending on the time scale under consideration. In their molecular structure, their molecules do not exhibit the long-range order exhibited by crystalline substances. In addition, while a glassy solid does exhibit some viscous flow and plastic deformation, this only occurs on geologic timescales. Thus, it behaves mechanically as a solid for all practical intents and purposes—and most experimental timescales. Common examples are silicate glasses, synthetic rubber and polystyrene and other polymers. Many amorphous solids soften into liquids when heated above their glass transition temperatures, at which the molecules become mobile.

Generally speaking, the atomic or molecular structure of glass exists in a metastable state with respect to its crystalline form. Glass would convert into a more stable crystalline form, but the rate of this conversion is slow. This essentially reflects the basic physics of glass, the glass transition, and its formation from a non-equilibrium supercooled liquid state.[7][8][9] Much work has been done to elucidate the primary microstructural features of glass forming substances (e.g. silicates) on both small (microscopic) and large (macroscopic) scales. One emerging school of thought is that a glass is simply the "limiting case" of a polycrystalline solid at small crystal size. Within this framework, domains, exhibiting various degrees of short-range order, become the building blocks of both metals and alloys, as well as glasses and ceramics. The microstructural defects of both within and between these domains provide the natural sites for atomic diffusion, and the occurrence of viscous flow and plastic deformation in solids.[10]

  • Note: Because solids have thermal energy, their atoms vibrate about fixed mean positions within the ordered (or disordered) lattice. The spectrum of lattice vibrations in a crystalline or glassy network provides the foundation for the kinetic theory of solids. This motion occurs at the atomic level, and thus cannot be observed or detected without highly specialized equipment—such as that used in spectroscopy.

Liquid crystal states

Liquid crystal states have properties intermediate between mobile liquids and ordered solids. For example, the nematic phase consists of long rod-like molecules such as para-azoxyanisole, which is nematic in the temperature range 118–136 °C.[11] In this state the molecules flow as in a liquid, but they all point in the same direction (within each domain) and cannot rotate freely.

Other types of liquid crystals are described in the main article on these states. Several types have technological importance, for example, in liquid crystal displays.


Transition metal atoms often have magnetic moments due to the net spin of electrons which remain unpaired and do not form chemical bonds. In some solids the magnetic moments on different atoms are ordered and can form a ferromagnet, an antiferromagnet or a ferrimagnet.

In a ferromagnet—for instance, solid iron—the magnetic moment on each atom is aligned in the same direction (within a magnetic domain). If the domains are also aligned, the solid is a permanent magnet, which is magnetic even in the absence of an external magnetic field. The magnetization disappears when the magnet is heated to the Curie point, which for iron is 768 °C.

An antiferromagnet has two networks of equal and opposite magnetic moments which cancel each other out, so that the net magnetization is zero. For example, in nickel(II) oxide (NiO), half the nickel atoms have moments aligned in one direction and half in the opposite direction.

In a ferrimagnet, the two networks of magnetic moments are opposite but unequal, so that cancellation is incomplete and there is a non-zero net magnetization. An example is magnetite (Fe3O4), which contains Fe2+ and Fe3+ ions with different magnetic moments.

Low-temperature states


Superconductors are materials which have zero electrical resistance, and therefore perfect conductivity. They also exclude all magnetic fields from their interior, a phenomenon known as the Meissner effect or perfect diamagnetism. Superconducting magnets are used as electromagnets in magnetic resonance imaging machines.

The phenomenon of superconductivity was discovered in 1911, and for 75 years was only known in some metals and metallic alloys at temperatures below 30 K. In 1986 so-called high-temperature superconductivity was discovered in certain ceramic oxides, and has now been observed in temperatures as high as 164 K.[12]


Close to absolute zero, some liquids form a second liquid state described as superfluid because it has zero viscosity or infinite fluidity. This was discovered in 1937 for helium which forms a superfluid below the lambda temperature of 2.17 K. In this state it will attempt to 'climb' out of its container.[13] It also has infinite thermal conductivity so that no temperature gradient can form in a superfluid.

These properties are explained by the theory that the common isotope helium-4 forms a Bose–Einstein condensate (see next section) in the superfluid state. More recently, Fermionic condensate superfluids have been formed at even lower temperatures by the rare isotope helium-3 and by lithium-6.[14]

Bose-Einstein condensates

In 1924, Albert Einstein and Satyendra Nath Bose predicted the "Bose-Einstein condensate," sometimes referred to as the fifth state of matter.

In the gas phase, the Bose-Einstein condensate remained an unverified theoretical prediction for many years. In 1995 the research groups of Eric Cornell and Carl Wieman, of JILA at the University of Colorado at Boulder, produced the first such condensate experimentally. A Bose-Einstein condensate is "colder" than a solid. It may occur when atoms have very similar (or the same) quantum levels, at temperatures very close to absolute zero (–273 °C).

Rydberg molecules

One of the metastable states of strongly non-ideal plasma is Rydberg matter, which forms upon condensation of excited atoms. These atoms can also turn into ions and electrons if they reach a certain temperature. In April 2009, Nature reported the creation of Rydberg molecules from a Rydberg atom and a ground state atom,[15] confirming that such a state of matter could exist.[16] The experiment was performed using ultracold rubidium atoms.

High-energy states

Plasma (ionized gas)

Plasmas or ionized gases can exist at temperatures starting at several thousand degrees C. Two examples of plasma are the charged air produced by lightning, and a star such as our own sun.

As a gas is heated, electrons begin to leave the atoms, resulting in the presence of free electrons, which are not bound to an atom or molecule, and ions, which are chemical species that contain unequal number of electrons and protons, and therefore possess an electrical charge. The free electric charges make the plasma electrically conductive so that it responds strongly to electromagnetic fields. At very high temperatures, such as those present in stars, it is assumed that essentially all electrons are "free," and that a very high-energy plasma is essentially bare nuclei swimming in a sea of electrons. Plasma is the most common state of non-dark matter in the universe.

A plasma can be considered as a gas of highly ionized particles, but the powerful interionic forces lead to distinctly different properties, so that it is usually considered as a different phase or state of matter.

Quark-gluon plasma

This is a state of matter discovered at the CERN in 2000, in which the quarks that would normally make up protons and neutrons are freed and can be observed individually, similar to splitting molecules into atoms. This state of matter allows scientists to observe the properties of individual quarks, and not just theorize. See also Strangeness production.

Other proposed states

Degenerate matter

Under extremely high pressure, ordinary matter undergoes a transition to a series of exotic states of matter collectively known as degenerate matter. These are of great interest to astrophysics, because these high-pressure conditions are believed to exist inside stars that have used up their nuclear fusion "fuel", such as white dwarves and neutron stars.


A supersolid is a spatially ordered material (that is, a solid or crystal) with superfluid properties. A supersolid is a solid, but exhibits so many other properties that many argue it is another state of matter.[17]

String-net liquid

When in a normal solid state, the atoms of matter align themselves in a grid pattern, so that the spin of any electron is the opposite of the spin of all electrons touching it. But in a string-net liquid, atoms are arranged in some pattern which would require some electrons to have neighbors with the same spin. This gives rise to some curious properties, as well as supporting some unusual proposals about the fundamental conditions of the universe itself.


A superglass is a phase of matter which is characterized at the same time by superfluidity and a frozen amorphous structure.

See also

Notes and references

  1. ^ D.L. Goodstein (1985). States of Matter. Dover Phoenix. ISBN 978-0486495064. 
  2. ^ A.P. Sutton (1993). Electronic Structure of Materials. Oxford Science Publications. pp. 10–12. ISBN 978-0198517542. 
  3. ^ M. Chaplin (20 August 2009). "Water phase Diagram". Water Structure and Science. Retrieved 2010-02-23. 
  4. ^ M.A. Wahab (2005). Solid State Physics: Structure and Properties of Materials. Alpha Science. pp. 1–3. ISBN 1842652184. 
  5. ^ F. White (2003). Fluid Mechanics. McGraw-Hill. p. 4. ISBN 0-07-240217-2. 
  6. ^ G. Turrell (1997). Gas Dynamics: Theory and Applications. John Wiley & Sons. pp. 3–5. ISBN 0471975737. 
  7. ^ S.V. Nemilov (1994). Thermodynamic and Kinetic Aspects of the Vitreous State. CRC Press. ISBN 0849337828. 
  8. ^ J. Zarzycki (1991). Glasses and the Vitreous State. Cambridge University Press. ISBN 0521355826. 
  9. ^ R. Zallen (1998). The Physics of Amorphous Solids. Wiley. pp. 3–5. ISBN 0471299413. 
  10. ^ J.C. Phillips (1982). "The Physics of Glass". Physics Today 35: 27. doi:10.1063/1.2914932. 
  11. ^ Shao, Y.; Zerda, T. W. (1998). "Phase Transitions of Liquid Crystal PAA in Confined Geometries". Journal of Physical Chemistry B 102 (18): 3387–3394. doi:10.1021/jp9734437. 
  12. ^ M. Tinkham (2004). Introduction to Superconductivity. Courier Dover. pp. 17–23. ISBN 0486435032. 
  13. ^ J.R. Minkel (20 February 2009). "Strange but True: Superfluid Helium Can Climb Walls". Scientific American. Retrieved 2010-02-23. 
  14. ^ L. Valigra (22 June 2005). "MIT physicists create new form of matter". MIT News. Retrieved 2010-02-23. 
  15. ^ V. Bendkowsky et al. (2009). "Observation of Ultralong-Range Rydberg Molecules". Nature 458: 1005. doi:10.1038/nature07945. 
  16. ^ V. Gill (23 April 2009). "World First for Strange Molecule". BBC News. Retrieved 2010-02-23. 
  17. ^ G. Murthy et al. (1997). "Superfluids and Supersolids on Frustrated Two-Dimensional Lattices". Physical Review B 55: 3104. doi:10.1103/PhysRevB.55.3104. 

External links

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Study guide

Up to date as of January 14, 2010

From Wikiversity

The States of Matter are laws of physics the bring the form of a material into what it's properties are,thus an example would be:

                              common form of

when a gas returns to it's liquid form, a procedure called CONDENSATION, takes place. e.g:


however, when a liguid transforms into a gas, a procedure called EVAPORATION, will continue. e.g:


finally, when a solid turns itself into a gas without passing through the liquid state, we have a term called SUBLIMATION in its tracks.e.g:

                                 cooled Carbon Dioxide
                                     Carbon Dioxide

Now we will focus on properties of materials, where we should notice the "jobs" of objects. For example, a bar of steel would be used the most common material for holding the structureof a household property, while elastic on the otherhand is too dangerous to set a foot in.

Simple English

There are four common states or phases of matter on earth: solid, liquid, gas, and plasma. The state of matter tends to determine its relative density, viscosity (how well it flows), and malleability (how easy it is to bend).


Phases of Matter


In a solid the positions of atoms are fixed relative to each other over long time. That is due to the cohesion or "friction" between molecules. This cohesion is provided by metallic, covalent or ionic bonds. Only solids can be pushed on by a force without changing shape, which means that they can be resistant to deformation. Solids also tend to be strong enough to hold their own shape in a container. Solids are generally denser than liquids.


In a liquid, molecules are attracted to other molecules strong enough to keep molecules in contact, but not strong enough to fix a particular structure. The molecules can continually move with respect to each other. This means that liquids can flow smoothly, but not as smoothly as gases. Liquids will tend to take the shape of a container that they are in. Liquids are generally less dense than solids, but denser than gases.


In a gas, the chemical bonds are not strong enough to hold atoms or molecules together, and from this a gas is a collection of independent, unbonded molecules which interact mainly by collision. Gases tend to take the shape of their container, and are less dense than both solids and liquids.


Plasmas are gases that have so much energy that the electrons that orbit an atom nucleus will launch away from the nucleus, like a rocket ship launching away from earth. Since the atom nuclei still "want" electrons, they are extremely good at conducting electricity. For example, air is not good at conducting electricity. However, in a lightning bolt, the atoms in air get so much energy that they no longer can hold on to their electrons, and become a plasma for a brief time. In this time, a current of electrons called lightning are able to travel down the air that is now a plasma. This will only occur in the direct path of a lightning bolt, but is why lightning can actually hit a car through its rubber tires.

It is said that the most common state of matter in the universe is plasma. This is partly due to the fact that space–which is where the freely-moving matter in the universe is–has a low pressure, which can also cause a change of phase.

Quark-Gluon Plasmas

Quark-gluon plasmas are a relatively newly discovered phase of matter that happen at about 2 trillion degrees Kelvin. Basically, scientists believe that protons and neutrons are held together by tiny things called quarks (which are "glued" together by things conveniently called "gluons"). At an incredibly high temperature only achievable by the Large Hadron Collider at CERN, quarks and gluons begin to separate into a new state of matter. Little is known about quark-gluon plasmas because of the sheer amount of energy needed to make them.

Bose-Einstein Condensates

Bose-Einstein condensates and fermionic condensates are phases of matter that apply to particles called bosons and fermions, respectively. (More than one boson can exist in the same spot at the same time, while only one fermion can exist in the same spot at the same time). Bose-Einstein condensates and fermionic condensates occur at incredibly low temperatures (about 4° Kelvin, which is the same as -452° Fahrenheit). Little is known about either state because of the sheer amount of energy needed to be taken away to create them. Inside of them, all of the particles begin to act like one big quantum state. That is, they have near-zero electrical resistance, and have almost no friction.


Many other states of matter exist, including strange matter, superfluids, and supersolids, and possibly string-net liquids. Scientists will continue to find more about phases of matter.

Phase Changes

Phases of matter can be changed by a number of things. This includes pressure and temperature.

When a solid becomes a liquid, it is called melting. When a solid becomes a gas, it is called sublimation. When a liquid becomes a gas, it is called vaporization. When a gas becomes a liquid, it is called condensation. When a liquid becomes a solid, it is called freezing. The freezing point and the melting point are said to be the same, because any increase in temperature will cause it to melt and any drop in temperature will cause it to freeze. This is also the reason that the vaporizing and condensation point are the same.

File:Phase change -
Phase changes

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