Wednesday, November 1, 2017

Plasma: The Fourth State Of Matter

Understanding what plasma is can be confusing, even though plasmas are part of our everyday lives. The Northern Lights, lightning, flames, fluorescent bulbs, neon signs and even the Sun are all examples of matter in this particular physical state. Although most of us are well acquainted with at least three states of matter - solid, liquid and gas - getting to know plasma, and how substances become plasma, can be daunting, but also fun and very useful.

Plasma Is Everywhere

It might surprise us to learn that plasma is by far the most abundant physical state of matter in the universe, much more so than our familiar three states of solid, liquid and gas. Perhaps even more surprising is the fact that plasma is just one of over 30 different states of matter currently listed on Wikipedia. More yet will undoubtedly be discovered in laboratories. Most of these exotic states are unknown beyond research circles because they tend not to be observed except under extraordinary conditions. Plasma, in contrast, is the most common state of ordinary matter under natural conditions in our universe.

Most of our visible universe consists of plasma, and it is a fortunate thing because that is the key reason why we can observe it with our telescopes. Stars are entirely made up of it. Most interstellar gas is plasma. Some sources estimate as much as 99% of visible matter in the universe is in a plasma state. By the end of this article we will understand how plasma can emit light. Even in our own solar system, over 99% of its mass is in a plasma state, thanks to the overwhelming mass contribution of our Sun. In our everyday life we regularly, but temporarily, witness plasma as lightning, fires, and the auroras, when energy is applied to matter, usually in its gaseous state. I did not explore plasma in high school and I had few occasions to explore it in my first years as an undergrad. One reason why plasma is not more thoroughly studied at these basic levels is that it can be difficult to jump from the fairly straightforward molecular theory needed to understand solids, liquids and gases into atomic and sub-atomic particle theories required to get a handle on the plasma state. Yet it's not only do-able; it's fascinating.

What Is a Phase Change?

A common theoretical thread to solids, liquids and gases is temperature. We can extend this temperature spectrum at the cold end by exploring the exotic-sounding Bose Einstein Condensate (BEC). In this state, mysterious quantum behaviours of matter are on a scale that we can observe. I wrote an article a few years ago that explores what happens when matter gets VERY cold and transforms into a BEC. Because stars are very hot, it is tempting to add plasma to the upper temperature end of this physical state spectrum. That, however, would be misleading and incorrect. Gases, when extremely hot, can break down into plasma but we do not need to heat a gas to transform it into plasma. Applying an electrical charge alone to a gas can do this. Think of a glowing neon tube light, for example. It is filled with gas in a plasma state but you can comfortably touch it with your hands.

There is an additional reason why temperature alone does not necessarily dictate the phase of a substance. We often first learn about phase changes by exploring water as a gas, liquid and solid, but when we talk about solids, liquids and gases, we really need to consider pressure in addition to temperature, which together determine what state matter is in. Under constant atmospheric pressure, temperature alone determines which state water will be in. This is how we encounter water in everyday life. We can also keep water at a constant temperature and change the pressure. Under enough pressure water will condense into liquid and even into exotic forms of ice. In a near-vacuum, water ice can sublimate directly into vapour.

Temperature, pressure and electrical potential are the three changing factors behind the phase changes we will focus on here. A phase change of any type relies on a boost or withdrawal of energy into/from matter. These aren't the only kinds of phase change that exist. The application of a magnetic field can change the physical state of magnetic material, for example. For all phase transitions (which is also called a change in physical state), the physical arrangement or ordering of atoms changes. In this article I'd like to start by exploring transitions between solids, liquids and gases in depth. Then I will compare those changes to transitions into a plasma state and explore what plasma is as a physical phase of matter.

Exploring Solids, Liquids and Gases Using Water as an Example

Matter changes from one physical state to another through a process called a phase transition. Every chemical compound and element undergoes a phase transition under a specific combination of temperature and pressure. The periodic stable below lists all the known elements. At 0°C and 1 atmospheric pressure, all elements in red type are gases, elements in green type are liquids, elements in black type are solids, and it is unknown what physical state those in grey type are. These "mystery" elements start at atomic number 100, fermium. These elements, plus several other smaller atomic number elements are synthetic or man-made. This is the highest atomic number element that can be created in a macroscopic amount, although a pure sample of this metal has not been created yet.

We can compare the melting and boiling points of most of the elements above but we are all most familiar with water. Not only is water familiar but it is also exhibits some unique and fascinating properties unlike most other substances, and this will allow us to explore change of state in greater depth. It is a chemical compound of hydrogen and oxygen (H2O). We drink it in a liquid state. At 0°C, water freezes into ice, a solid state, and at 100°C it evaporates into a gas, water vapour. These transitions assume a pressure of one atmosphere (atm; the average air pressure at sea level)). To understand how pressure plays a role in the physical transition of a substance, imagine water vapour in a container fitted with a gas-tight plunger. The temperature is carefully maintained throughout the experiment at 130°C. The starting pressure is 1 atm. As the plunger lowers down into the container, the volume of gas decreases and the gas pressure increases. Soon the water vapour will be compressed into liquid water. 130°C is the boiling point at around 5 atm. Above 5 atm, water at 130°C will liquid. It will have to be compressed much further, to about 100,000 atm (because liquid water strongly resists compression) to solidify into scalding hot 130°C ice in the container.

Researchers at Sandia National Laboratories performed a similar experiment several years ago. They subjected liquid water (starting at 1 atm) to extremely rapid compression (to 70,000 atm). The water shrank abruptly into a dense phase of solid ice. When the pressure was relieved, it melted (and expanded) back into liquid water. The ice formed under these conditions is not the everyday ice we make in our freezers. Ordinary freezer ice, as most of us know, is less dense than liquid water. At pressures above 100,000 atm, however, water only exists as different kinds of very dense ice, even at temperatures of hundreds of degrees centigrade. Scientists think such hot ice exists in the deep interior of exoplanet GJ 436 b, a Neptune-size ice giant that orbits very close to its parent red dwarf star, GJ 436, which is 33 light-years away from Earth.

A Phase Change is a Molecular Rearrangement

Chemically, water in any physical state remains water. That is, it retains its chemical properties. Any matter undergoing a phase change from solid to liquid to gas and vice versa retains its chemical properties. Each water molecule consists of two hydrogen atoms covalently bonded to an oxygen atom. What changes during a phase transition is the molecular arrangement of these molecules (see below). At any particular temperature and pressure, water molecules will adopt the most thermodynamically stable arrangement possible for that environment. In order to transition into a new arrangement, the molecule-molecule interactions must change. Molecule-molecule, or in the case of elements, atom-atom interactions can shift and reorganize a substance, changing its physical properties in the process. The interactions between molecules in a substance tend to be attractive. In solids, these attractive forces are strong enough to be called chemical bonds, but these bonds are still much weaker than the chemical bonds that bind each molecule itself together (such as the two covalent hydrogen bonds in water). The simple images below highlight the general differences in atomic/molecular arrangement between solids, liquids and gases.

This is an actual atomic resolution image of the lattice-like arrangement of molecules in solid strontium titanate. Chemical bonds between strontium titanate molecules hold them close together in a tight regular arrangement. Above about 2000°C, strontium titanate crystals will melt, breaking these relatively weak intermolecular chemical bonds.

Kaneiderdaniel;German-language Wikipedia
This highly simplified two-dimensional representation (left) shows how atoms or molecules might be arranged in a typical liquid in a beaker. They have contact with their closest neighbours, where they experience very weak attractive inter-molecular forces, but there is no overall order to their arrangement, and the atoms/molecules can slide freely past one another.

As the diagram right simplistically shows, the spaces between molecules/atoms in a gas are vast compared to the size of the atoms or molecules themselves, many magnitudes greater than this simple diagram suggests. This is true even for a highly pressurized gas. The atoms/molecules can move freely in all directions. They experience no appreciable attractive forces between them so they move independently of one another. The ideal gas law assumes negligible molecular volume relative to gas volume and it assumes the absence of any molecular attraction. It quite accurately predicts how volume, temperature and pressure relate to one another in a gas. This means that most gases in reality act much like a theoretical ideal gas.

Water molecules bond with each other in a solid phase according to the Bernal-Fowler ice rules. This is where water gets quite interesting. Every oxygen atom can bond to 4 hydrogen atoms. Two bonds are strong. These are the covalent chemical bonds that hold each water molecule together, and they don't change during a phase transition. However, oxygen has 6 valence (or outer, chemically available) electrons. In water vapour, four electrons remain unbound as two lone pairs. As water vapour condenses into liquid water, some of the four unbound electrons take part in very weak molecule-molecule bonds with hydrogen atoms in adjacent molecules. These are called hydrogen bonds. Unlike the strong covalent hydrogen-oxygen bonds holding each molecule together, these bonds are weak and in the liquid state, they are transitory. The (proton) positive charges of hydrogen atoms are attracted to the negative charge zones of nearby oxygen lone electron pairs.

The attractive force behind hydrogen bonding is about 90% due to the attraction between opposite charges. This is an electrostatic phenomenon. 10% of the bonding force is due to electron sharing. This is a quantum mechanical phenomenon and it is also responsible for (much stronger) covalent bonding. In liquid water, the weak hydrogen attractions form and break very easily, allowing water molecules to slip past one another. The much stronger covalent intra-molecular bonds holding each water molecule together only break when water undergoes a chemical decomposition reaction called electrolysis, producing hydrogen and oxygen gases as products.

The diagram below models the hydrogen bonds between water molecules. Two lone electron pairs force every water molecule into a triangle shape, which makes it polar - each molecule has a positively charged region and a negatively charged region. The positive charge of hydrogen atoms is attracted to the negative charge zone of oxygen atoms. Hydrogen bonds are indicated by dotted lines.

User Qwerter at Czech Wikipedia    
The length and strength of hydrogen bonds is strongly dependent on temperature. As liquid water freezes, additional hydrogen bonds form between molecules. Every hydrogen atom is, in effect, bonded to two oxygen atoms - one bond is strong and one bond is weak. Under sufficient pressure and/or cold, water molecules come close enough together to bond into regular and stable lattice-like arrangements. As ordinary ice, the molecules form bonds that result in a hexagonal crystalline lattice arrangement. Under more extreme pressure/temperature regimes, water ice can transition into a variety of different lattice structures such as cubes, rhomboids and tetragons, which allow for denser molecular arrangements. Each transition into a denser crystalline lattice is a phase change. There are at least 17 known physical states of water ice alone. Changes in the number, length and strength of hydrogen bonds underlie each of these additional solid phase changes.

Most liquid substances solidify into denser molecular arrangements but there are a few exceptions, and they are always due to unusual bonding between the molecules. Water is an example. Ordinary water ice is actually less dense than liquid water. That is why it floats. This phase is called hexagonal ice, or ice lh. It is the only solid water phase encountered on Earth. Due to the unusual chemical bonds of water (responsible for its molecular shape and polarity), the hexagonal lattice arrangement of lh ice keeps molecules a bit further apart from one another than they are in the liquid state (yet the bonds themselves are stronger and more permanent). Ih ice has a density of 0.9167 g/cm3 compared to liquid water with a density of 1.18 g/cm3. The diagram below shows what the hexagonal lattice of Ih ice looks like. Gray dashed lines indicate hydrogen bonds.

In addition to a number of solid crystalline lattice states, water can also exist as a solid lacking any crystalline structure, called amorphous ice. On Saturn's icy moon, Enceladus, water ice strewn onto the surface from its many cryovolcanoes is amorphous. Liquid water flash-freezes in the vacuum of space (zero pressure) before it can organize into any structure. In contrast, along Enceladus's unique "tiger stripes," the ice is crystalline because here it is kept warm enough long enough from the heat of geothermal activity to arrange into a more thermodynamically favourable crystalline structure.

Below is Cassini's view of Enceladus's south pole, showing 4-5 "tiger stripes," which are tectonic fractures.

NASA/JPL/Space Science Institute
Critical Point and Triple Point

The temperature and pressure at which a phase change occurs is called a critical point. The temperature of the critical point depends on the pressure of the system and vice versa, so on a graph you have a line rather than a single point (see the phase graph below for water). However, there is a single point, a specific temperature and pressure, at which gas, liquid and solid states of water all have an identical free energy (or internal molecular energy). All three phases coexist at this point, which is called the triple point. A critical point is a line and a triple point is a point. In the graph below, two black lines represent the solid/liquid critical point and the liquid/gas critical point of water. In the lower center of the graph you can see the single triple for water. It is a single point on the graph found at 0.01°C (273.6°K) and 0.006 atm (611.657 pascals).

A note about pressure units on the above graph: 1 atm pressure is roughly equal to 1 bar pressure. 1 bar is equal to 100 kilopascals (kPa). The pascal is the internationally recognized SI unit for pressure. The bar and the atmosphere are older non-Si units. I am in the old habit of using atm units, where 1 atm is standard air pressure at sea level.

This 6-minute video from Bergen University in Norway demonstrates how water behaves as it approaches its triple point.

At the triple point, all of the water in a system can be changed into vapour, liquid or solid just by tweaking the pressure or temperature a tiny bit. The triple point is also the lowest pressure at which liquid water can exist. At lower pressures, as on the surface of an icy moon (at near vacuum), water ice, perhaps warmed when facing its star, will sublimate directly into vapour, bypassing the liquid state altogether. For most substances, the triple point is also the lowest temperature at which the liquid state can exist but water, due to its unusual hydrogen bonds, is an exception. Notice the odd little horn on the green section in the graph above. This is an anomaly of water. If the temperature is just below the triple point, a steadily increasing pressure will transform ordinary solid water ice (Ih) into denser liquid water (at around 1000 atm) and then back into an even denser solid (now as ice VI) at around 10,000 atm.

A Brief Look at The Thermodynamics of a Phase Change

For the phase changes I've described so far, the transition process is abrupt, or more scientifically put, discontinuous. At a phase transition point (a critical point), such as the boiling point for water for example, two phases - liquid water and water vapour - co-exist at one specific temperature. In thermodynamic terms, they have the same average Gibbs free energy. In any system some molecules will randomly have slightly more energy and some will have slightly less. However, at critical point the substance cannot be distinguished as either gas or liquid.

All thermodynamic systems tend to assume the lowest possible Gibbs free energy. Systems also tend toward increasing entropy. Gibbs free energy can be looked at as a measure of potential chemical energy in a system that is available to do work in the system, while entropy can be thought of as a measure of the orderliness of a system. Highly ordered crystalline water ice has very low entropy compared to the high entropy of water vapour. Both states, however, are stable within their own temperature/pressure regime because each state has maximized its entropy according to the Gibbs free energy available to it.

Below the boiling point, the liquid phase is more thermodynamically stable (which means it has reached maximum possible entropy at a lower Gibbs free energy). As water vapour condenses into liquid water, it enters a more highly ordered state (lower entropy).  As this phase transition occurs, an energy exchange takes place. As the water enters a state of lower Gibbs free energy, it releases the extra free energy as latent heat. Conversely, if we wanted to melt ice into water we would have to apply heat to it. In other words we would have to add latent heat to the system so that we can increase the entropy of the water molecules.

So far we've focused on water as our example, but every chemical substance and element can exist as a solid, liquid and gas, depending on its temperature and pressure. Each substance has its own unique critical points and triple point, and it may have additional solid lattice states as well. These critical points are unique to each substance and depend on the substance's unique molecular bonding.

The Relationship Between Pressure and Phase Change

We've explored how substances and elements transition from solids to liquids to gases and vice versa. Increasing the pressure on a substance or increasing its temperature increases the average kinetic energy of that substance. Pressure and temperature determine which physical state a particular substance is most thermodynamically stable. The relationship between temperature and physical state is fairly straightforward as we've seen. However, the relationship between pressure and physical state depends on the starting physical state of the substance. A gas responds to pressure by shrinking in volume and increasing kinetic energy. A gas resists compression through thermal pressure: Even though a container full of gas appears static, the molecules in it are in constant motion. Collisions between those molecules and the sides of the container are detected as a force per unit area, or pressure. If a gas is compressed, it will heat up because the total thermal energy of that gas, which doesn't change, is now concentrated into a smaller volume. If heat is continually removed from the system as it's compressed, the gas will remain at the original temperature it was even though it will eventually condense into a liquid and then further solidify into a solid.

Liquids, unlike gases, tend to be fairly incompressible. This means that liquids don't experience much increase in kinetic energy when they are compressed, but they will compress under extreme pressure. The pressure that resists compression is not thermal pressure, as with gases,  but another kind of outward pressure (which gases also express but is masked by thermal pressure). This pressure is the result of the Pauli exclusion principle. This quantum mechanical principle means that when electron orbitals in adjacent atoms are forced into very close proximity, they will powerfully resist overlapping one another into the same atomic orbital. To note, this is not the same thing as orbital sharing, in which an electron will occupy an empty orbital in an adjacent atom, and through such sharing create a chemical bond. I will explain this more accurately in a moment.

Under increasing pressure, the liquid will likely first become denser and more viscous and then it will transition into a solid, in which molecules will pack more closely together but in a very ordered arrangement that respects the exclusion principle. If you recall my earlier example of liquid water being subjected to sudden extreme pressure and condensing into a dense form of ice, in both cases it eventually becomes more thermodynamically favourable for the water molecules to transition into a high-density highly ordered form of ice than to remain as a highly pressurized liquid. That "decision" made by a substance is abrupt, which means it happens all at once.

Water Under Very Intense Pressure

What happens if you continue to increase the pressure on a solid such as water ice? Solids, like liquids, resist compression. If we take a look at the water phase graph above once again, we see that above 100,000 atm (or bars), water exists only as a solid no matter what the temperature is, at least up to 350°C. Cubic ice VII will transform into an even denser cubic lattice called ice X as the pressure is increased. Its crystalline lattice will transition into increasingly denser molecular arrangements, each transition being a phase change. At the highest pressure on the graph (around 6-10 million bars) we see a phase of ice called high-pressure ice XI. This extremely dense XI hexagonal ice is hypothetical at this time and should not be confused with ice XI orthorhombic, which likely forms at temperatures below -200°C at zero pressure, such as on the surface of Pluto's moon, Hydra, for example.

A 2010 article by Burkhard Militzer and Hugh Wilson explains what researchers think happens when the pressure on ice increases even further than our graph goes. In this theoretical case, water bypasses hypothetical high-pressure ice XI. They suggest instead that above about 3 million bar, Ice X could transition into a different more complex lattice that contains 12 hydrogen atoms per unit. If the temperature is also increased to about 2700°C, they think the lattice might transition into yet another new arrangement. The hydrogen atoms may become mobile within a stable lattice consisting only of oxygen atoms. This would create a super-ionic phase. If the temperature is increased even further under this intense pressure regime, the oxygen atoms themselves may also become mobile and the lattice itself might melt into a new extremely dense unstructured phase of matter that no longer consists of water as we know it. The atoms are no longer chemically bound to each other. This is an exception to the simpler rule we started with - that the chemistry of substances does not change during a phase transition. As one goes deeper into almost any theory, rules based on simpler understandings tend to hit the roadside.

At temperatures below 2700°C and pressures above 48 million bar, a transition to a metallic ice phase is possible. Such intense pressure is expected to exist deep within icy Uranus or Neptune. In this phase, the structure of the ice would resemble stacked corrugated sheets of oxygen and hydrogen atoms, which would be electrically conductive.

Even more extreme pressures can be forced on matter, for example, inside the core of a rapidly collapsing star. What would happen to water under such conditions? As pressure is increased within a solid, individual atoms resist being pressed closer together. To explain this, we need to revisit the electron orbital. Hydrogen is a simple example to illustrate this. This atom (at ground state) has one electron in its 1s electron orbital. Any orbital can hold a maximum of two electrons, so 1s could also accommodate a second electron if there was one. An electron in another atom's orbital can temporarily occupy that role. If it is another hydrogen atom, then a covalent electron-sharing chemical bond is created, turning atomic hydrogen into hydrogen gas, H2.

An atom is a system, and like all physical systems, it tends toward the lowest free energy state possible. There are many forces going on in a hydrogen molecule. The two electrons repulse each other. The two protons repulse each other. Each electron is attracted to both proton nuclei and vice versa. There is an optimum distance between electrons and protons where all of these electrostatic forces add up to a lowest possible energy state. This is the most stable state. When atoms are squeezed closer together (or pulled further part), they resist and a force must be applied. Although perhaps easier to visualize, the electrostatic interactions that I've just described here are insignificant compared to much more powerful quantum interactions that respect the exclusion principle. Two or more electrons cannot share identical quantum numbers. If two electrons share an orbital they must be of opposite spin (spin is a quantum number). This exclusion principle translates into a repulsive quantum force that very powerfully resists additional pressure put onto an already extremely dense solid phase.

As pressure increases, the atomic orbital structure itself is forced to shift. Normally, electrons fill only a few energy levels in an atom. Many energy levels are unoccupied. Under extreme pressure, all the electrons are forced into the lowest energy level orbitals, as close to the nucleus as possible. This is ultra-dense electron-degenerate matter that is expected to exist in the core of a white dwarf stellar remnant. The exclusion principle means that two same-spin electrons won't share an orbital no matter how strongly they are forced together. Electrons, experiencing such intense repulsive quantum forces, respond by moving faster and faster in these lowest orbitals. Under enough pressure, they approach the speed of light, which is the limit of this arrangement. Protons begin to absorb electrons, creating a degenerate phase of matter, called neutron matter which is the densest state of matter possible. It is thought to exist in super nova stellar core fragments called neutron stars. Such transitions into extremely dense matter are technically phase changes as well, as is the final transition of matter into a black hole, in which matter as we know it collapses entirely.

Where Plasma Fits In As A Phase Transition

Where does plasma fit into this progression of phase transitions? The temperature of a substance or element is the measure of the average vibrational energy of its constituent molecules or atoms. Gas molecules generally have a great deal of vibrational energy as well as kinetic energy. They move around in all directions at great velocities, widely separated from each other. They occasionally bump into one another, and the force of these collisions translates into gas pressure. If a gas is cooled, it will eventually condense into a liquid at its critical point. Pressurizing a gas adjusts that critical point to a higher temperature. If you take another look at the phase diagram of water, you notice that water will boil into water vapour at just 50°C, rather than at 100°C if the pressure is 1/10th that of atmospheric pressure. At the other extreme, the temperature of water streaming from a deep sea hydrothermal vent can reach over 400°C but it does not boil because the pressure is over 300 times atmospheric pressure at depths over 3000 m, where most hydrothermal vents are located.

One way to make plasma is to heat gas. What happens if we heat a container of gas rather than compress it? We increase the total kinetic energy of the molecules. If we heat a sample of hydrogen gas enough, we will eventually supply enough energy to break the chemical bonds and dissociate the gas into hydrogen atoms. The bond dissociation energy of H2 is about 50,000°C but this temperature depends on the density of the gas. As density increases, the bond dissociation energy increases (due to two particles taking up more space than one). Now a mono-atomic gas, the hydrogen atoms can be further energized into ions. The ionization energy of hydrogen is about 150,000°C at 1 atm. This is also pressure-dependent. As density increases, the ionization energy decreases. As atoms are forced together, the gaps between electrons energy levels get shorter, which means less energy is required to ionize an atom.

Ionization means that the atom loses one or more electrons to create a gas that contains positive and negative ions. If we ionize hydrogen gas, we create a cloud of electrons and protons, called plasma. A very dense liquid-like plasma state in which protons are surrounded by a sea of mobile electrons, called metallic hydrogen, might exist deep inside Jupiter and Saturn. Because electrons are not confined to orbitals, this is also called a degenerate state of matter.

The process of ionization from a neutral gas into charged plasma is considered by many researchers to be the fourth physical phase after gases, liquids and solids. This phase transition, however, is a different and gradual, rather than abrupt, process. Atoms with more than one electron tend to lose electrons gradually as increasing energy is applied to a gas. The other three transitions - solid to liquid to gas - are each marked by an abrupt change in the arrangement of the molecules in a substance. In a transition from a gas into plasma, unlike the other three phase transitions, the substance also chemically changes because at least some molecular bonds are broken. For example, water vapour can ionize into hydrogen/oxygen plasma within a lightning bolt. The process of ionization into plasma is reversible, a characteristic in common with the three other changes of state. If energy is removed from the plasma, the ions will recombine into neutral atoms and molecules. However, in a complex mixture of gases, some new molecules might form as different ions react with one another.

The process of ionization takes place over a series of steps. We will go back to hydrogen as our example. As energy is pumped into atomic hydrogen gas, the kinetic energy of the atoms increases and some atoms become excited. An atom is excited through two possible processes: collisions with other atoms and the absorption of electromagnetic energy (light).

The electron in a hydrogen atom at ground state is located near the nucleus in a lowest possible energy orbital, which is actually a cloud of possible locations rather than a defined circular orbit. The electron in each hydrogen atom can absorb either kinetic or photon energy and move to a higher energy orbital. Electrons can even move into higher energy orbitals that are not ordinarily occupied by electrons. The atom is now in an excited state. If energy is removed from this system, the electron will return to its lowest energy (ground) state by releasing exactly the same energy as the energy difference between the orbitals. The energy is quantized. It is released as photons of specific wavelengths. For hydrogen, some of these photons are in the visible range. An excited hydrogen atom commonly emits red or aqua blue photons depending on the orbital drop, but it can also emit higher energy ultraviolet photons if it absorbs and releases more energy. As energy is added to the system, the electron continues to move up into higher energy orbitals until it eventually breaks free from the atom altogether. The proton nucleus can no longer hold onto the electron so it is now a completely ionized nucleus without any electrons. If an atom with more than one electron loses some of its electrons, but not all,  it is partially ionized. Our hydrogen atom is now a dissociated electron and proton. A gas of these ions is called completely ionized plasma, shown in the simple diagram below.

In this simple diagram (left) of completely ionized plasma, all electrons have been stripped from their nuclei, creating an electrically conductive "electron sea."

Under real conditions, low-energy plasmas often contain a mixture of excited, partially ionized and neutral atoms, which can be used to emit a beautiful glow in a gas-discharge lamp, for example. The tube below, filled with diffuse low-energy hydrogen plasma, glows pink - a mixture of red and aqua blue photon emissions.; Wikipedia
The plasma in the tube above contains mostly neutral and excited atoms and only a few ionized atoms. The degree of ionization in plasma, sometimes called plasma density or electron density, is the number of free electrons in a volume of plasma. Even a partially ionized gas which contains only 1% ionized atoms, can be considered plasma if it exhibits plasma behaviours such as responding to magnetic fields and conducting electricity.

The plasma in the tube above was not created by heating hydrogen gas until it ionized; this tube is at room temperature. Extreme heat is one way to create plasma. Hydrogen in the interior of the Sun and other stars is extremely high-energy completely ionized plasma because it is an incredibly hot environment. Any substance, if hot enough, will transition into plasma. Even the exotic solid states of water, which remain solid even up to 400°C, would eventually break down into plasma as temperature is increased. I should note here that this process is not chemical ionization of water, in which water (H20) self-ionizes into hydroxide (0H-) and hydronium (H30+) ions. That is a chemical equilibrium reaction, and no change of phase occurs.

Rather than through heat, the hydrogen plasma in the tube above was created by applying an electric potential to hydrogen gas. This is another way that energy is applied to a gas in order to ionize it. Instead of electrons being "shaken off" of an atom with high kinetic energy, electrons (which, remember, are charge-carrying particles) are drawn off the atom by a powerful electrical force. It is maybe analogous to ducklings being swept away from their mom down a stream with a powerful current. It can take tremendous energy to completely ionize a gas. It all depends on which molecules and elements the gas consists of and what kind of energy is applied. When we focus in on what happens to an ionizing atom, it is more accurate to talk about energy in term s of electron volts (eV) than in temperature, which is an average of particle energies. The most tightly bound (and stable) atom is helium. It has two electrons, and therefore two ionization energies - it takes 25 eV to remove just one electron (this partially ionizes the atom). It takes much more energy (about 54 eV) to strip both electrons off the helium nucleus (now the atom is completely ionized). If a collision with another atom strikes with over 54 eV of energy, the target atom can be completely ionized.

It takes far less energy to excite atoms rather than ionize them. Even a relatively small 120V household circuit can light up a small neon lamp, which is a vacuum tube filled with mostly neon gas and a little argon gas.

In gases, ionization is reversible. If the energy is removed from the system, such as turning off the applied voltage gradient (unplugging the hydrogen tube above), the protons recombine with free electrons, the atoms return to ground state, and molecules recombine into neutral gas (H2). In such a closed system of only a single atomic gas, the gas can be ionized and recombined, transitioned into liquid and solid states and back again, over and over, demonstrating an entirely reversible process, but under real conditions, the extreme energy of some plasmas can trigger various chemical reactions as well, particularly combustion reactions. These reactions are non-reversible and add a non-reversible component to the phase change.

In solids, in particular, the process is not easy to describe in terms of a phase change. An example here might be a fulgurite. Lightning strikes sand and leaves behind a hollow tube of glass buried in the ground. Lightning is an extremely powerful electric potential, of up to 500,000 volts. The melting point of silicon dioxide (pure sand) is 1710°C and the boiling point is 2230°C. The interior of a lightning bolt can reach a temperature of 28,000°C, far higher than the boiling point of sand, enough to chemically break it apart and at least partially ionize its silicon and oxygen atoms. When lightning strikes sand, some sand is explosively vapourized into gas and plasma, leaving a hollow tube, where the bolt struck, surrounded by a layer of molten sand that quickly solidifies into glass. If there were impurities in the sand, such as soil and plant debris, these components would have combusted and reacted with each other, resulting in new chemical compounds in the fulgurite's glass. Air molecules in the vicinity are also broken down into plasma temporarily. In this case, a series of very rapid phase changes has occurred from solid to liquid to gas to plasma, but there was also opportunity for chemical reactions to take place, which are irreversible. Even the air itself, which is a mixture of gases, does not transition to a plasma state and then back into neutral gases without some chemical changes taking place. Some highly reactive oxygen ions and excited oxygen molecules will recombine into ozone molecules, for example, creating the fresh-air smell after a thunderstorm.

Why isn't a neon sign or lamp hot? Neon lights might get warm to the touch but they never get hot. Yet we can think of the ionized atoms in plasma as hot. The atoms have at least enough energy to become excited and lose some outermost electrons. The plasma itself contains some free fast moving electrons with significant kinetic energy, but in the case of the low-density plasma in a neon lamp, it also contains mostly neutral atoms with far less kinetic and thermal energy. This means that its average temperature will be simply warm to the touch. A neon light contains very diffuse plasma and most of the neon gas remains in a neutral gas state that absorbs the excess energy of the electrons that collide with them. It doesn't take much energy to create glowing plasma in which only a few outermost electrons of atoms are excited and fewer still are stripped off, creating a small electrical current in the tube that sustains the excited-atom glow. Ordinary air is an exception. It is actually a powerful electrical insulator and would make a lousy plasma light. It will ionize and glow (this is a lightning bolt) only when it is subjected to a very powerful electrical potential of more than approximately 100,000V. The dielectric strength of air, its ability to withstand a potential gradient before breaking down or ionizing, is 3.0 MV/m (million volts/metre), which is much higher than neon's, which is 0.02 MV/m. Dielectric strength is an intrinsic property of a material.

Understanding Plasma is Essential To Understanding Our Universe

Plasmas, no matter how they are created, have unique physical properties. They act quite differently from neutral gases. Like gases, plasmas do not have a definite shape or volume. They can be compressed fairly easily. Unlike gases however, which are electrically neutral, plasmas respond to electric and magnetic fields. Even though the charges are usually balanced overall in plasma, they are separate and free to move. This means they can produce electric currents and magnetic fields and they respond to them as well. Electromagnetic forces exerted on plasmas act on them across very long distances. This gives plasma special coherent behaviours that gases never display. The Sun and other stars are examples of plasma created under extreme heat. Their interiors consist of extremely dense, energetic, and completely ionized plasma. Powerful plasma currents exchange heat released from ongoing nuclear fusion in the stellar core. These physical currents are also powerful electrical currents because they are moving charges. The moving charges set up incredibly intense magnetic fields that can interfere with one another and snap. These are the mechanisms behind the violent solar weather that can damage communications, satellites and electrical systems 150 million miles away on Earth. While stars are made of dense plasma, most of the visible matter in the universe consists of very diffuse highly ionized plasma. Both kinds of plasma are the key reason why we can see distant stars and various gas clouds. They glow because they are plasmas that contain atoms that are continually being excited and returning to ground state, emitting photons of light in the process. Intergalactic space, interstellar space and interplanetary space are also all (extremely diffuse) plasma. Solar wind is plasma and Earth's ionosphere in the upper atmosphere consists of diffuse atmospheric gases ionized by solar radiation, or plasma in other words. At one point, prior to recombination, the entire universe consisted of completely ionized plasma. It was too energetic to support intact atoms. Expansion (and therefore cooling) allowed that primordial plasma to settle into the first and most abundant simple atoms such as hydrogen and helium. For these reasons, some theorists think that plasma theory should play a more dominant role in understanding the cosmology of the universe, alongside general relativity, high-energy astronomy, mechanics and dynamics.

Plasma behaviour can be extremely complex. Understanding the complicated dynamics of stars, for example, requires computer modeling based on plasma theory, much of that theory belonging to the field of magnetohydrodynamics, a field of study that only got under way in last few decades. Understanding the electrodynamics of plasma (the behaviour of an electrical fluid) on the largest scales might also help to explain the evolution of galaxies as well as their birth from the collapse of interstellar clouds into stars, which is not yet completely understood. Plasma theory might also shine some light on the mysterious nature of dark matter, required to explain the rotation curves of galaxies. Black holes, quasars and active galactic nuclei must all incorporate plasma dynamics in order for us to fully understand how they work. Extreme plasma dynamics could even represent missing pieces in our understanding of cosmic inflation and the accelerating expansion of the universe, the latter of which is now called dark energy. An intuitive understanding of plasma could be a useful key to understanding the way the universe works.