The physical realm is made up of matter. On Earth, matter appears in three clearly defined forms—solid, liquid, and gas—whose visible and perceptible structure is a function of behavior that takes place at the molecular level. Though these are often referred to as "states" of matter, it is also useful to think of them as phases of matter. This terminology serves as a reminder that any one substance can exist in any of the three phases. Water, for instance, can be ice, liquid, or steam; given the proper temperature and pressure, it may be solid, liquid, and gas all at once! But the three definite earthbound states of matter are not the sum total of the material world: in outer space a fourth phase, plasma, exists—and there may be still other varieties in the physical universe.
Matter can be defined as physical substance that has mass; occupies space; is composed of atoms; and is ultimately convertible to energy. A significant conversion of matter to energy, however, occurs only at speeds approaching that of the speed of light, a fact encompassed in the famous statement formulated by Albert Einstein (1879-1955), E = mc2.
Einstein's formula means that every item possesses a quantity of energy equal to its mass multiplied by the squared speed of light. Given the fact that light travels at 186,000 mi (297,600 km) per second, the quantities of energy available from even a tiny object traveling at that speed are massive indeed. This is the basis for both nuclear power and nuclear weaponry, each of which uses some of the smallest particles in the known universe to produce results that are both amazing and terrifying.
The forms of matter that most people experience in their everyday lives, of course, are traveling at speeds well below that of the speed of light. Even so, transfers between matter and energy take place, though on a much, much smaller scale. For instance, when a fire burns, only a tiny fraction of its mass is converted to energy. The rest is converted into forms of mass different from that of the wood used to make the fire. Much of it remains in place as ash, of course, but an enormous volume is released into the atmosphere as a gas so filled with energy that it generates not only heat but light. The actual mass converted into energy, however, is infinitesimal.
The property of energy is, at all times and at all places in the physical universe, conserved. In physics, "to conserve" something means "to result in no net loss of" that particular component—in this case, energy. Energy is never destroyed: it simply changes form. Hence, the conservation of energy, a law of physics stating that within a system isolated from all other outside factors, the total amount of energy remains the same, though transformations of energy from one form to another take place.
Whereas energy is perfectly conserved, matter is only approximately conserved, as shown with the example of the fire. Most of the matter from the wood did indeed turn into more matter—that
The conservation of mass holds that total mass is constant, and is unaffected by factors such as position, velocity, or temperature, in any system that does not exchange any matter with its environment. This, however, is a qualified statement: at speeds well below c (the speed of light), it is essentially true, but for matter approaching c and thus, turning into energy, it is not.
Consider an item of matter moving at the speed of 100 mi (160 km)/sec. This is equal to 360,000 MPH (576,000 km/h) and in terms of the speeds to which humans are accustomed, it seems incredibly fast. After all, the fastest any human beings have ever traveled was about 25,000 MPH (40,000 km/h), in the case of the astronauts aboard Apollo 11 in May 1969, and the speed under discussion is more than 14 times greater. Yet 100 mi/sec is a snail's pace compared to c : in fact, the proportional difference between an actual snail's pace and the speed of a human walking is not as great. Yet even at this leisurely gait, equal to 0.00054 c, a portion of mass equal to 0.0001% (one-millionth of the total mass) converts to energy.
In his brilliant work Six Easy Pieces, American physicist Richard Feynman (1918-1988) asked his readers, "If, in some cataclysm, all of scientific knowledge were to be destroyed, and only one sentence passed on to the next generations of creatures, what statement would contain the most information in the fewest words? I believe it is the atomic hypothesis (or the atomic fact, or whatever you wish to call it) that all things are made of atoms—little articles that move around in perpetual motion, attracting each other when they are a little distance apart, but repelling upon being squeezed into one another. In that sentence, you will see, there is an enormous amount of information about the world, if just a little imagination and thinking are applied."
Feynman went on to offer a powerful series of illustrations concerning the size of atoms relative to more familiar objects: if an apple were magnified to the size of Earth, for instance, the atoms in it would each be about the size of a regular apple. Clearly atoms and other atomic particles are far too small to be glimpsed by even the most highly powered optical microscope. Yet, it is the behavior of particles at the atomic level that defines the shape of the entire physical world. Viewed from this perspective, it becomes easy to understand how and why matter is convertible to energy. Likewise, the interaction between atoms and other particles explains why some types of matter are solid, others liquid, and still others, gas.
An atom is the smallest particle of a chemical element. It is not, however, the smallest particle in the universe; atoms are composed of subatomic particles, including protons, neutrons, and electrons. But at the subatomic level, it is meaningless to refer to, for instance, "an oxygen electron": electrons are just electrons. An atom, then, is the fundamental unit of matter. Most of the substances people encounter in the world, however, are not pure elements, such as oxygen or iron; they are chemical compounds, in which atoms of more than one element join together to form molecules.
One of the most well-known molecular forms in the world is water, or H2O, composed of two hydrogen atoms and one oxygen atom. The arrangement is extremely precise and never varies: scientists know, for instance, that the two hydrogen atoms join the oxygen atom (which is much larger than the hydrogen atoms) at an angle of 105° 3′. Other molecules are much more complex than those of water—some of them much, much more complex, which is reflected in the sometimes unwieldy names required to identify their chemical components.
The idea of atoms is not new. More than 24 centuries ago, the Greek philosopher Democritus (c. 470-380 B.C.) proposed that matter is composed of tiny particles he called atomos, or "indivisible." Democritus was not, however, describing matter in a concrete, scientific way: his "atoms" were idealized, philosophical constructs, not purely physical units.
Yet, he came amazingly close to identifying the fundamental structure of physical reality—much closer than any number of erroneous theories (such as the "four elements" of earth, air, fire, and water) that prevailed until modern times. English chemist John Dalton (1766-1844) was the first to identify what Feynman later called the "atomic hypothesis": that nature is composed of tiny particles. In putting forward his idea, Dalton adopted Democritus's word "atom" to describe these basic units.
Dalton recognized that the structure of atoms in a particular element or compound is uniform. He maintained that compounds are made up of compound atoms: in other words, that water, for instance, is a compound of "water atoms." Water, however, is not an element, and thus, it was necessary to think of its atomic composition in a different way—in terms of molecules rather than atoms. Dalton's contemporary Amedeo Avogadro (1776-1856), an Italian physicist, was the first scientist to clarify the distinction between atoms and molecules.
Obviously, it is impractical to weigh a single molecule, or even several thousand; what was needed, then, was a number large enough to make possible practical comparisons of mass. Hence, the mole, a quantity equal to "Avogadro's number." The latter, named after Avogadro though not derived by him, is equal to 6.022137 × 1023 (more than 600 billion trillion) molecules.
The term "mole" can be used in the same way that the word "dozen" is used. Just as "a dozen" can refer to twelve cakes or twelve chickens, so "mole" always describes the same number of molecules. A mole of any given substance has its own particular mass, expressed in grams. The mass of one mole of iron, for instance, will always be greater than that of one mole of oxygen. The ratio between them is exactly the same as the ratio of the mass of one iron atom to one oxygen atom. Thus, the mole makes it possible to compare the mass of one element or compound to that of another.
Contemporary to both Dalton and Avogadro was Scottish naturalist Robert Brown (1773-1858), who in 1827 stumbled upon a curious phenomenon. While studying pollen grains under a microscope, Brown noticed that the grains underwent a curious zigzagging motion in the water. At first, he assumed that the motion had a biological explanation—that is, it resulted from life processes within the pollen—but later he discovered that even pollen from long-dead plants behaved in the same way.
Brown never understood what he was witnessing. Nor did a number of other scientists, who began noticing other examples of what came to be known as Brownian motion: the constant but irregular zigzagging of particles in a puff of smoke, for instance. Later, however, Scottish physicist James Clerk Maxwell (1831-1879) and others were able to explain it by what came to be known as the kinetic theory of matter.
The kinetic theory, which is discussed in depth elsewhere in this book, is based on the idea that molecules are constantly in motion: hence, the water molecules were moving the pollen grains Brown observed. Pollen grains are many thousands of times as large as water molecules, but there are so many molecules in just one drop of water, and their motion is so constant but apparently random, that they are bound to move a pollen grain once every few thousand collisions.
Einstein, who was born the year Maxwell died, published a series of papers in which he analyzed the behavior of particles subjected to Brownian motion. His work, and the confirmation of his results by French physicist Jean Baptiste Perrin (1870-1942), finally put an end to any remaining doubts concerning the molecular structure of matter.
It may seem amazing that the molecular and atomic ideas were still open to question in the early twentieth century; however, the vast majority of what is known today concerning the atom emerged after World War I. At the end of the nineteenth century, scientists believed the atom to be indivisible, but growing evidence concerning electrical charges in atoms brought with it the awareness that there must be something smaller creating those charges.
Eventually, physicists identified protons and electrons, but the neutron, with no electrical charge, was harder to discover: it was not identified until 1932. After that point, scientists were convinced that just three types of subatomic particles existed. However, subsequent activity among physicists—particularly those in the field of quantum mechanics—led to the discovery of other elementary particles, such as the photon. However, in this discussion, the only subatomic particles whose behavior is reviewed are the proton, electron, and neutron.
At the molecular level, every item of matter in the world is in motion. This may be easy enough to imagine with regard to air or water, since both tend to flow. But what about a piece of paper, or a glass, or a rock? In fact, all molecules are in constant motion, and depending on the particular phase of matter, this motion may vary from a mere vibration to a high rate of speed.
Molecular motion generates kinetic energy, or the energy of movement, which is manifested as heat or thermal energy. Indeed, heat is really nothing more than molecules in motion relative to one another: the faster they move, the greater the kinetic energy, and the greater the heat.
The movement of atoms and molecules is always in a straight line and at a constant velocity, unless acted upon by some outside force. In fact, the motion of atoms and molecules is constantly being interfered with by outside forces, because they are perpetually striking one another. These collisions cause changes in direction, and may lead to transfers of energy from one particle to another.
The behavior of molecules cannot be explained in terms of gravitational force. This force, and the motions associated with it, were identified by Sir Isaac Newton (1642-1727), and Newton's model of the universe seemed to answer most physical questions. Then in the late nineteenth century, Maxwell discovered a second kind of force, electromagnetism. (There are two other known varieties of force, strong and weak nuclear, which are exhibited at the subatomic level.) Electromagnetic force, rather than gravitation, explains the attraction between atoms.
Several times up to this point, the subatomic particles have been mentioned but not explained in terms of their electrical charge, which is principal among their defining characteristics. Protons have a positive electrical charge, while neutrons exert no charge. These two types of particles, which make up the vast majority of the atom's mass, are clustered at the center, or nucleus. Orbiting around this nucleus are electrons, much smaller particles which exert a negative charge.
Chemical elements are identified by the number of protons they possess. Hydrogen, first element listed on the periodic table of elements, has one proton and is thus identified as 1; carbon, or element 6, has six protons, and so on.
An atom usually has a neutral charge, meaning that it is composed of an equal number of protons or electrons. In certain situations, however, it may lose one or more electrons and thus acquire a net charge. Such an atom is called an ion. But electrical charge, like energy, is conserved, and the electrons are not "lost" when an atom becomes an ion: they simply go elsewhere.
Positive and negative charges interact at the molecular level in a way that can be compared to the behavior of poles in a pair of magnets. Just as two north poles or two south poles repel one another, so like charges—two positives, or two negatives—repel. Conversely, positive and negative charges exert an attractive force on one another similar to that of a north pole and south pole in contact.
In discussing phases of matter, the attraction between molecules provides a key to distinguishing between states of matter. This is not to say one particular phase of matter is a particularly good conductor of electrical current, however. For instance, certain solids—particularly metals such as copper—are extremely good conductors. But wood is a solid, too, and conducts electrical current poorly.
The properties of various forms of matter, viewed from the larger electromagnetic picture, are a subject far beyond the scope of this essay. In any case, the electromagnetic properties of concern in the present instance are not the ones demonstrated at a macroscopic level—that is, in view of "the big picture." Rather, the subject of the attractive force operating at the atomic or molecular levels has been introduced to show that certain types of material have a greater intermolecular attraction.
As previously stated, all matter is in motion. The relative speed of that motion, however, is a function of the attraction between molecules, which in turn defines a material according to one of the phases of matter. When the molecules in a material exert a strong attraction toward one another, they move slowly, and the material is called a solid. Molecules of liquid, by contrast, exert a moderate attraction and move at moderate speeds. A material substance whose molecules exert little or no attraction, and therefore, move at high speeds, is known as a gas.
These comparisons of molecular speed and attraction, obviously, are relative. Certainly, it is easy enough in most cases to distinguish between one phase of matter and another, but there are some instances in which they overlap. Examples of these will follow, but first it is necessary to discuss the phases of matter in the context of their behavior in everyday situations.
The attractions between particles have a number of consequences in defining the phases of matter. The strong attractive forces in solids cause its particles to be positioned close together. This means that particles of solids resist attempts to compress them, or push them together. Because of their close proximity, solid particles are fixed in an orderly and definite pattern. As a result, a solid usually has a definite volume and shape.
A crystal is a type of solid in which the constituent parts are arranged in a simple, definite geometric arrangement that is repeated in all directions. Metals, for instance, are crystalline solids. Other solids are said to be amorphous, meaning that they possess no definite shape. Amorphous solids—clay, for example—either possess very tiny crystals, or consist of several varieties of crystal mixed randomly. Still other solids, among them glass, do not contain crystals.
Because of their strong attractions to one another, solid particles move slowly, but like all particles of matter, they do move. Whereas the particles in a liquid or gas move fast enough to be in relative motion with regard to one another, however, solid particles merely vibrate from a fixed position.
This can be shown by the example of a singer hitting a certain note and shattering a glass. Contrary to popular belief, the note does not have to be particularly high: rather, the note should be on the same wavelength as the vibration of the glass. When this occurs, sound energy is transferred directly to the glass, which shatters because of the sudden net intake of energy.
As noted earlier, the attraction and motion of particles in matter has a direct effect on heat and temperature. The cooler the solid, the slower and weaker the vibrations, and the closer the particles are to one another. Thus, most types of matter contract when freezing, and their density increases. Absolute zero, or 0K on the Kelvin scale of temperature—equal to −459.67°F (−273°C)—is the point at which vibration virtually stops. Note that the vibration virtually stops, but does not stop entirely. In any event, the lowest temperature actually achieved, at a Finnish nuclear laboratory in 1993, is 2.8 · 10−10 K, or 0.00000000028K—still above absolute zero.
The behavior of water at the freezing/melting point is interesting and exceptional. Above 39.2°F (4°C) water, like most substances, expands when heated. But between 32°F (0°C) and that temperature, however, it actually contracts. And whereas most substances become much denser with lowered temperatures, the density of water reaches its maximum at 39.2°F. Below that point, it starts to decrease again.
Not only does the density of ice begin decreasing just before freezing, but its volume increases. This is the reason ice floats: its weight is less than that of the water it has displaced, and therefore, it is buoyant. Additionally, the buoyant qualities of ice atop very cold water explain why the top of a lake may freeze, but lakes rarely freeze solid—even in the coldest of inhabited regions.
Instead of freezing from the bottom up, as it would if ice were less buoyant than the water, the lake freezes from the top down. Furthermore, ice is a poorer conductor of heat than water, and, thus, little of the heat from the water below escapes. Therefore, the lake does not freeze completely—only a layer at the top—and this helps preserve animal and plant life in the body of water. On the other hand, the increased volume of frozen water is not always good for humans: when water in pipes freezes, it may increase in volume to the point where the pipe bursts.
When heated, particles begin to vibrate more and more, and, therefore, move further apart. If a solid is heated enough, it loses its rigid structure and becomes a liquid. The temperature at which a solid turns into a liquid is called the melting point, and melting points are different for different substances. For the most part, however, solids composed of heavier particles require more energy—and, hence, higher temperatures—to induce the vibrations necessary for freezing. Nitrogen melts at −346°F (−210°C), ice at 32°F (0°C), and copper at 1,985°F (1,085°C). The melting point of a substance, incidentally, is the same as its freezing point: the difference is a matter of orientation—that is, whether the process is one of a solid melting to become a liquid, or of a liquid freezing to become a solid.
The energy required to change a solid to a liquid is called the heat of fusion. In melting, all the heat energy in a solid (energy that exists due to the motion of its particles) is used in breaking up the arrangement of crystals, called a lattice. This is why the water resulting from melted ice does not feel any warmer than when it was frozen: the thermal energy has been expended, with none left over for heating the water. Once all the ice is melted, however, the absorbed energy from the particles—now moving at much greater speeds than when the ice was in a solid state—causes the temperature to rise.
The particles of a liquid, as compared to those of a solid, have more energy, more motion, and less attraction to one another. The attraction, however, is still fairly strong: thus, liquid particles are in close enough proximity that the liquid resists compression.
On the other hand, their arrangement is loose enough that the particles tend to move around one another rather than merely vibrating in place, as solid particles do. A liquid is therefore not definite in shape. Both liquids and gases tend to flow, and to conform to the shape of their container; for this reason, they are together classified as fluids.
Owing to the fact that the particles in a liquid are not as close in proximity as those of a solid, liquids tend to be less dense than solids. The liquid phase of substance is thus inclined to be larger in volume than its equivalent in solid form. Again, however, water is exceptional in this regard: liquid water actually takes up less space than an equal mass of frozen water.
When a liquid experiences an increase in temperature, its particles take on energy and begin to move faster and faster. They collide with one another, and at some point the particles nearest the surface of the liquid acquire enough energy to break away from their neighbors. It is at this point that the liquid becomes a gas or vapor.
As heating continues, particles throughout the liquid begin to gain energy and move faster, but they do not immediately transform into gas. The reason is that the pressure of the liquid, combined with the pressure of the atmosphere above the liquid, tends to keep particles in place. Those particles below the surface, therefore, remain where they are until they acquire enough energy to rise to the surface.
The heated particle moves upward, leaving behind it a hollow space—a bubble. A bubble is not an empty space: it contains smaller trapped particles, but its small weight relative to that of the liquid it disperses makes it buoyant. Therefore, a bubble floats to the top, releasing its trapped particles as gas or vapor. At that point, the liquid is said to be boiling.
As they rise, the particles thus have to overcome atmospheric pressure, and this means that the boiling point for any liquid depends in part on the pressure of the surrounding air. This is why cooking instructions often vary with altitude: the greater the distance from sea level, the less the air pressure, and the shorter the required cooking time.
Atop Mt. Everest, Earth's highest peak at about 29,000 ft (8,839 m) above sea level, the pressure is approximately one-third normal atmospheric pressure. This means the air is one-third as dense as it is as sea level, which explains why mountain-climbers on Everest and other tall peaks must wear oxygen masks to stay alive. It also means that water boils at a much lower temperature on Everest than it does elsewhere. At sea level, the boiling point of water is 212°F (100°C), but at 29,000 ft it is reduced by one-quarter, to 158°F (70°C).
Of course, no one lives on the top of Mt. Everest—but people do live in Denver, Colorado, where the altitude is 5,577 ft (1,700 m) and the boiling point of water is 203°F (95°C). Given the lower boiling point, one might assume that food would cook faster in Denver than in New York, Los Angeles, or some other city close to sea level. In fact, the opposite is true: because heated particles escape the water so much faster at high altitudes, they do not have time to acquire the energy needed to raise the temperature of the water. It is for this reason that a recipe may contain a statement such as "at altitudes above XX feet, add XX minutes to cooking time."
If lowered atmospheric pressure means a lowered boiling point, what happens in outer space, where there is no atmospheric pressure? Liquids boil at very, very low temperatures. This is one of the reasons why astronauts have to wear pressurized suits: if they did not, their blood would boil—even though space itself is incredibly cold.
Note that the process of a liquid changing to a gas is similar to what occurs when a solid changes to a liquid: particles gain heat and therefore energy, begin to move faster, break free from one another, and pass a certain threshold into a new phase of matter. And just as the freezing and melting point for a given substance are the same temperature, the only difference being one of orientation, the boiling point of a liquid transforming into a gas is the same as the condensation point for a gas turning into a liquid.
The behavior of water in boiling and condensation makes possible distillation, one of the principal methods for purifying seawater in various parts of the world. First, the water is boiled, then, it is allowed to cool and condense, thus forming water again. In the process, the water separates from the salt, leaving it behind in the form of brine. A similar separation takes place when salt water freezes: because salt, like most solids, has a much lower freezing point than water, very little of it remains joined to the water in ice. Instead, the salt takes the form of a briny slush.
Having reached the gaseous state, a substance takes on characteristics quite different from those of a solid, and somewhat different from those of a liquid. Whereas liquid particles exert a moderate attraction to one another, particles in a gas exert little to no attraction. They are thus free to move, and to move quickly. The shape and arrangement of gas is therefore random and indefinite—and, more importantly, the motion of gas particles give it much greater kinetic energy than the other forms of matter found on Earth.
The constant, fast, and random motion of gas particles means that they are always colliding and thereby transferring kinetic energy back and forth without any net loss. These collisions also have the overall effect of producing uniform pressure in a gas. At the same time, the characteristics and behavior of gas particles indicate that they will tend not to remain in an open container. Therefore, in order to maintain any pressure on a gas—other than the normal atmospheric pressure exerted on the surface of the gas by the atmosphere (which, of course, is also a gas)—it is necessary to keep it in a closed container.
There are a number of gas laws (examined in another essay in this book) describing the response of gases to changes in pressure, temperature, and volume. Among these is Boyle's law, which holds that when the temperature of a gas is constant, there is an inverse relationship between volume and pressure: in other words, the greater the pressure, the less the volume, and vice versa. According to a second gas law, Charles's law, for gases in conditions of constant pressure, the ratio between volume and temperature is constant—that is, the greater the temperature, the greater the volume, and vice versa.
In addition, Gay-Lussac's law shows that the pressure of a gas is directly related to its absolute temperature on the Kelvin scale: the higher the temperature, the higher the pressure, and vice versa. Gay-Lussac's law is combined, along with Boyle's and Charles's and other gas laws, in the ideal gas law, which makes it possible to find the value of any one variable—pressure, volume, number of moles, or temperature—for a gas, as long as one knows the value of the other three.
Principal among states of matter other than solid, liquid, and gas is plasma, which is similar to gas. (The term "plasma," when referring to the state of matter, has nothing to do with the word as it is often used, in reference to blood plasma.) As with gas, plasma particles collide at high speeds—but in plasma, the speeds are even greater, and the kinetic energy levels even higher.
The speed and energy of these collisions is directly related to the underlying property that distinguishes plasma from gas. So violent are the collisions between plasma particles that electrons are knocked away from their atoms. As a result, plasma does not have the atomic structure typical of a gas; rather, it is composed of positive ions and electrons. Plasma particles are thus electrically charged, and, therefore, greatly influenced by electrical and magnetic fields.
Formed at very high temperatures, plasma is found in stars and comets' tails; furthermore, the reaction between plasma and atomic particles in the upper atmosphere is responsible for the aurora borealis or "northern lights." Though not found on Earth, plasma—ubiquitous in other parts of the universe—may be the most plentiful among the four principal states of matter.
Among the quasi-states of matter discussed by physicists are several terms that describe the structure in which particles are joined, rather than the attraction and relative movement of those particles. "Crystalline," "amorphous," and "glassy" are all terms to describe what may be individual states of matter; so too is "colloidal."
A colloid is a structure intermediate in size between a molecule and a visible particle, and it has a tendency to be dispersed in another medium—as smoke, for instance, is dispersed in air. Brownian motion describes the behavior of most colloidal particles. When one sees dust floating in a ray of sunshine through a window, the light reflects off colloids in the dust, which are driven back and forth by motion in the air otherwise imperceptible to the human senses.
The number of states or phases of matter is clearly not fixed, and it is quite possible that more will be discovered in outer space, if not on Earth. One intriguing candidate is called dark matter, so described because it neither reflects nor emits light, and is therefore invisible. In fact, luminous or visible matter may very well make up only a small fraction of the mass in the universe, with the rest being taken up by dark matter.
If dark matter is invisible, how do astronomers and physicists know it exists? By analyzing the gravitational force exerted on visible objects when there seems to be no visible object to account for that force. An example is the center of our galaxy, the Milky Way, which appears to be nothing more than a dark "halo." In order to cause the entire galaxy to revolve around it in the same way that planets revolve around the Sun, the Milky Way must contain a staggering quantity of invisible mass.
Dark matter may be the substance at the heart of a black hole, a collapsed star whose mass is so great that its gravitational field prevents light from escaping. It is possible, also, that dark matter is made up of neutrinos, subatomic particles thought to be massless. Perhaps, the theory goes, neutrinos actually possess tiny quantities of mass, and therefore in huge groups—a mole times a mole times a mole—they might possess appreciable mass.
In addition, dark matter may be the deciding factor as to whether the universe is infinite. The more mass the universe possesses, the greater its overall gravity, and if the mass of the universe is above a certain point, it will eventually begin to contract. This, of course, would mean that it is finite; on the other hand, if the mass is below this threshold, it will continue to expand indefinitely. The known mass of the universe is nowhere near that threshold—but, because the nature of dark matter is still largely unknown, it is not possible yet to say what effect its mass may have on the total equation.
Physicists at the Joint Institute of Laboratory Astrophysics in Boulder, Colorado, in 1995 revealed a highly interesting aspect of atomic behavior at temperatures approaching absolute zero. Some 70 years before, Einstein had predicted that, at extremely low temperatures, atoms would fuse to form one large "superatom." This hypothesized structure was dubbed the Bose-Einstein Condensate (BEC) after Einstein and Satyendranath Bose (1894-1974), an Indian physicist whose statistical methods contributed to the development of quantum theory.
Cooling about 2,000 atoms of the element rubidium to a temperature just 170 billionths of a degree Celsius above absolute zero, the physicists succeeded in creating an atom 100 micrometers across—still incredibly small, but vast in comparison to an ordinary atom. The superatom, which lasted for about 15 seconds, cooled down all the way to just 20 billionths of a degree above absolute zero. The Colorado physicists won the Nobel Prize in physics in 1997 for their work.
In 1999, researchers in a lab at Harvard University also created a superatom of BEC, and used it to slow light to just 38 MPH (60.8 km/h)—about 0.02% of its ordinary speed. Dubbed a "new" form of matter, the BEC may lead to a greater understanding of quantum mechanics, and may aid in the design of smaller, more powerful computer chips.
At places throughout this essay, references have been made variously to "phases" and "states" of matter. This is not intended to confuse, but rather to emphasize a particular point. Solids, liquids, and gases are referred to as "phases," because substances on Earth—water, for instance—regularly move from one phase to another. This change, a function of temperature, is called (aptly enough) "change of phase."
There is absolutely nothing incorrect in referring to "states of matter." But "phases of matter" is used in the present context as a means of emphasizing the fact that most substances, at the appropriate temperature and pressure, can be solid, liquid, or gas. In fact, a substance may even be solid, liquid, and gas.
The phases of matter can be likened to the phases of a person's life: infancy, babyhood, childhood, adolescence, adulthood, old age. The transition between these stages is indefinite, yet it is
At the transition point between adolescence and adulthood—say, at seventeen years old—a young person may say that she is an adult, but her parents may insist that she is still an adolescent or a child. And indeed, she might qualify as either. On the other hand, when she is thirty, it would be ridiculous to assert that she is anything other than an adult.
At the same time, a person at a certain age may exhibit behaviors typically associated with another age. A child, for instance, may behave like an adult, or an adult like a baby. One interesting example of this is the relationship between age two and late adolescence. In both cases, the person is in the process of individualizing, developing an identity separate from that of his or her parents—yet clearly, there are also plenty of differences between a two-year-old and a seventeen-year-old.
As with the transitional phases in human life, in the borderline pressure levels and temperatures for phases of matter it is sometimes difficult to say, for instance, if a substance is fully a liquid or fully a gas. On the other hand, at a certain temperature and pressure level, a substance clearly is what it is: water at very low temperature and pressure, for instance, is indisputably ice—just as an average thirty-year-old is obviously an adult. As for the second observation, that a person at one stage in life may reflect characteristics of another stage, this too is reflected in the behavior of matter.
A liquid crystal is a substance that, over a specific range of temperature, displays properties both of a liquid and a solid. Below this temperature range, it is unquestionably a solid, and above this range it is just as obviously a liquid. In between, however, liquid crystals exhibit a strange solid-liquid behavior: like a liquid, their particles flow, but like a solid, their molecules maintain specific crystalline arrangements.
Long, wide, and placed alongside one another, liquid crystal molecules exhibit interesting properties in response to light waves. The speed of light through a liquid crystal actually varies, depending on whether the light is traveling along the short or long sides of the molecules. These differences in light speed may lead to a change in the direction of polarization, or the vibration of light waves.
The cholesteric class of liquid crystals is so named because the spiral patterns of light through the crystal are similar to those which appear in cholesterols. Depending on the physical properties of a cholesteric liquid crystal, only certain colors may be reflected. The response of liquid crystals to light makes them useful in liquid crystal displays (LCDs) found on laptop computer screens, camcorder views, and in other applications.
In some cholesteric liquid crystals, high temperatures lead to a reflection of shorter visible light waves, and lower temperatures to a display of longer visible waves. Liquid crystal thermometers thus show red when cool, and blue as they are warmed. This may seem a bit unusual to someone who does not understand why the thermometer displays those colors, since people typically associate red with heat and blue with cold.
A liquid crystal exhibits aspects of both liquid and solid, and thus, at certain temperatures may be classified within the crystalline quasi-state of matter. On the other hand, the phenomenon known as the triple point shows how an ordinary substance, such as water or carbon dioxide, can actually be a liquid, solid, and vapor—all at once.
Again, water—the basis of all life on Earth—is an unusual substance in many regards. For instance, most people associate water as a gas or vapor (that is, steam) with very high temperatures. Yet, at a level far below normal atmospheric pressure, water can be a vapor at temperatures as low as −4°F (−20 °C). (All of the pressure values in the discussion of water at or near the triple point are far below atmospheric norms: the pressure at which water would turn into a vapor at −4°F, for instance, is about 1/1000 normal atmospheric pressure.)
As everyone knows, at relatively low temperatures, water is a solid—ice. But if the pressure of ice falls below a very low threshold, it will turn straight into a gas (a process known as sublimation) without passing through the liquid stage. On the other hand, by applying enough pressure, it is possible to melt ice, and thereby transform it from a solid to a liquid, at temperatures below its normal freezing point.
The phases and changes of phase for a given substance at specific temperatures and pressure levels can be plotted on a graph called a phase diagram, which typically shows temperature on the x-axis and pressure on the y-axis. The phase diagram of water shows a line between the solid and liquid states that is almost, but not quite, exactly perpendicular to the x-axis: it slopes slightly upward to the left, reflecting the fact that solid ice turns into water with an increase of pressure.
Whereas the line between solid and liquid water is more or less straight, the division between these two states and water vapor is curved. And where the solid-liquid line intersects the vaporization curve, there is a place called the triple point. Just below freezing, in conditions equivalent to about 0.7% of normal atmospheric pressure, water is a solid, liquid, and vapor all at once.
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The smallest particle of a chemical element. An atom can exist either alone or in combination with other atoms in a molecule. Atoms are made up of protons, neutrons, and electrons. In most cases, the electrical charges in atoms cancel out one another; but when an atom loses one or more electrons, and thus has a net charge, it becomes an ion.
A substance made up of atoms of more than one chemical element. These atoms are usually joined in molecules.
A substance made up of only one kind of atom.
A law of physics which holds that within a system isolated from all other outside factors, the total amount of energy remains the same, though transformations of energy from one form to another take place.
A physical principle which states that total mass is constant, and is unaffected by factors such as position, velocity, or temperature, in any system that does not exchange any matter with its environment. Unlike the other conservation laws, however, conservation of mass is not universally applicable, but applies only at speeds significant lower than that of light—186,000 mi (297,600 km) per second. Close to the speed of light, mass begins converting to energy.
In physics, "to conserve" something means "to result in no net loss of" that particular component. It is possible that within a given system, the component may change form or position, but as long as the net value of the component remains the same, it has been conserved.
Negatively charged particles in an atom. Electrons, which spin around the nucleus of protons and neutrons, constitute a very small portion of the atom's mass. In most atoms, the number of electrons and protons is the same, thus canceling out one another. When an atom loses one or more electrons, however—thus becoming an ion—it acquires a net electrical charge.
The force that resists motion when the surface of one object comes into contact with the surface of another.
Any substance, whether gas or liquid, that tends to flow, and that conforms to the shape of its container. Unlike solids, fluids are typically uniform in molecular structure for instance, one molecule of water is the same as another water molecule.
A phase of matter in which molecules exert little or no attraction toward one another, and therefore move at high speeds.
An atom that has lost or gained one or more electrons, and thus has a net electrical charge.
A phase of matter in which molecules exert moderate attractions toward one another, and therefore move at moderate speeds.
Physical substance that has mass; occupies space; is composed of atoms; and is ultimately (at speeds approaching that of light) convertible to energy. There are several phases of matter, including solids, liquids, and gases.
A unit equal to 6.022137 × 1023 (more than 600 billion trillion) molecules. Their size makes it impossible to weigh molecules in relatively small quantities; hence the mole facilitates comparisons of mass between substances.
A group of atoms, usually of more than one chemical element, joined in a structure.
A subatomic particle that has no electrical charge. Neutrons are found at the nucleus of an atom, alongside protons.
The various forms of material substance (matter), which are defined primarily in terms of the behavior exhibited by their atomic or molecular structures. On Earth, three principal phases of matter exist, namely solid, liquid, and gas. Other forms of matter include plasma.
One of the phases of matter, closely related to gas. Plasma apparently does not exist on Earth, but is found, for instance, in stars and comets' tails. Containing neither atoms nor molecules, plasma is made up of electrons and positive ions.
A positively charged particle in an atom. Protons and neutrons, which together form the nucleus around which electrons orbit, have approximately the same mass—a mass that is many times greater than that of an electron.
A phase of matter in which molecules exert strong attractions toward one another, and therefore move slowly.
In physics, the term "system" usually refers to any set of physical interactions isolated from the rest of the universe. Anything outside of the system, including all factors and forces irrelevant to a discussion of that system, is known as the environment.