Phases - Gas, liquid and solid Phases - Gas, liquid and solid Skills to develop Define and explain what a phase is. Interpret these phase transitions: melting and freezing. Phases - Gas, liquid and solid Gas, liquid, and solid are known as the three states of matter or material, but each of solid and liquid states may exist in one or more forms. Thus, another term is required to describe the various forms, and the term phase is used. Each distinct form is called a phase, but the concept of phase defined as a homogeneous portion of a system, extends beyond a single material, because a phase may also involve several materials. For example, a homogeneous solution of any number of substances is a one-phase system. Phase is a concept used to explain many physical and chemical changes (reactions). A solid has a definite shape and volume. A liquid has a definite volume but it takes the shape of a container whereas a gas fills the entire volume of a container. You already know that diamond and graphite are solids made up of the element carbon. They are two phases of carbon, but both are solids. Solids are divided into subclasses of amorphous (or glassy) solids and crystalline solids.

Arrangements of atoms or molecules in crystalline solids are repeated regularly over a very long range of millions of atoms, but their arrangements in amorphous solids are somewhat random or short range of say some tens or hundreds of atoms. In general, there is only one liquid phase of a material. However, there are two forms of liquid helium, each have some unique properties. Thus, the two forms are different (liquid) phases of helium. At a definite temperature and pressure, the two phases co-exist. So far, all gases behave alike as do mixtures of gases. Thus, a gas is usually considered as a phase. The Concept of Phase A phase is a distinct and homogeneous state of a system with no visible boundary separating it into parts. Conversion between these states is called a phase transition. Water, H2O, is the most common substance that its gas (steam), liquid (water), and solid (ice) phases are widely known. An ice water mixture has two phases, so are systems containing ice-and-vapour, and water-and-vapour. To recognize the vapour system in these system may require a keen observation, because the vapour usually blends with air, and is not detected directly.

You probably also know that several solids may exist for a substance, and each of the solid form is also called a phase. Diamond and graphite being the most quoted examples. Graphite and diamond are both solid carbon. They have different crystal shapes, colors, and structures. They represent two different phases of carbon. Under 1 atm, ice has hexagonal symmetry, while cubic ice is formed under high pressure. In fact, there are at least 8 different types of ice, each being a solid phase. When you mix water and alcohol, regardless of the relative amounts that you use, they are fully miscible. The resulting mixture has only one phase (a solution). However, water and oil are normally immiscible, and their boundary of separation is visible. They form a two-phase system. Sometimes you cannot "see" the boundary, and you will need scientific reasoning to realize the number of phases present in system. Well, there is so much concept packed into one term that we can not make the definition any simpler for you. However, the term is useful because it can be used to explain many phenomena. There is no substitute for it. Learn it and use it to explain physical changes.

Phase Transitions A state change of any material due to temperature or pressure change is a phase transition. A phase transition is a physical change (or reaction). The following diagram illustrates the key phase transitions: You should know the names of the process for these phase transitions. sublimation deposition SOLID ============> GAS ==============> SOLID melting freezing SOLID ============> LIQUID =============> SOLID (solidfication) condensation vaporization GAS =============> LIQUID =============> GAS The concepts of phase and phase transition introduce you into fields of materials studies. For example, if you search the internet with the phrase "phase transition", you get thousands of websites, some are related to the concept we have discussed here, but some may be using "phase transition" as a catchy phrase. For example, a Netscape Search gives a site for Studies of the phase transitions in calcite includes several papers published by a professor of Earth Sciences in UK whose career is based on the study of calcite, CaCO3, a very common mineral. The concept of phase transition is also applied to the study of nuclear matter such as protons and neutrons.

Confidence Building Questions Solid to Gas Phase Transition About Watch and Favorite Watch Watching this resources will notify you when proposed changes or new versions are created so you can keep track of improvements that have been made. Favorite Favoriting this resource allows you to save it in the “My Resources” tab of your account. There, you can easily access this resource later when you’re ready to customize it or assign it to your students. Solid to Gas Phase Transition Sublimation is the phase transition from the solid to the gaseous phase, without passing through an intermediate liquid phase. Learning Objective Discuss the process of sublimation. Key Points Sublimation is an endothermic phase transition in which a solid evaporates to a gas . Solids that sublimate have such high vapor pressures that heating leads to a substantial vaporization even before the melting point is reached. The enthalpy of sublimation (also called heat of sublimation) can be calculated as the sum of the enthalpy of fusion and the enthalpy of vaporization. Terms Full Text Phase Transition: Solid to Gas Sublimation is the process of transformation directly from the solid phase to the gaseous phase, without passing through an intermediate liquid phase.

It is an endothermic phase transition that occurs at temperatures and pressures below a substance 's triple point (the temperature and pressure at which all three phases coexist) in its phase diagram . At a given temperature, most chemical compounds and elements can possess one of the three different states of matter at different pressures. In these cases, the transition from the solid to the gaseous state requires an intermediate liquid state. But at temperatures below that of the triple point, a decrease in pressure will result in a phase transition directly from the solid to the gaseous. Also, at pressures below the triple point pressure, an increase in temperature will result in a solid being converted to gas without passing through the liquid region. Notice the triple point of the substance. At temperatures and pressures below those of the triple point, a phase change between the solid and gas phases can take place. For some substances, such as carbon and arsenic, sublimation is much easier than evaporation . This is because the pressure of their triple point is very high and it is difficult to obtain them as liquids. The solid has such high vapor pressures that heating leads to a substantial amount of direct vaporization even before the melting point is reached.

The process of sublimation requires additional energy and is therefore an endothermic change. The enthalpy of sublimation (also called heat of sublimation) can be calculated as the sum of the enthalpy of fusion and the enthalpy of vaporization. The reverse process of sublimation is deposition (i.e., gas to solid). For example, solid iodine, I2, is easily sublimed at temperatures around 100°C. Even ice has a measurable vapor pressure near its freezing point , as evidenced by the tendency of snow to evaporate in cold dry weather. There are other solids whose vapor pressure overtakes that of the liquid before melting can occur. Such substances sublime; a common example is solid carbon dioxide (dry ice) at 1 atm of atmospheric pressure. Heat of Vaporization - Chemistry LibreTexts Heat of Vaporization Contributors The Heat (or Enthalpy) of Vaporization is the quantity of heat that must be absorbed if a certain quantity of liquid is vaporized at a constant temperature. In a solution with both a vaporized and liquidized states, the kinetic energy of the vapor is higher than the kinetic energy of the liquid.

Temperature follows kinetic energy, showing a lower temperature in the remaining liquid. Introduction The Enthalpy of Vaporization is measured using constant temperature, meaning that something must be added to bring the lower temperature to the temperature of the solution prior to vaporization. To raise the temperature (and as a result, the kinetic energy), heat is added. The Enthalpy of Vaporization measures the amount of heat added, because it is the quantity that must be absorbed for a liquid to remain at constant temperature during vaporization. \(ΔH_{vap}\) is expressed in kJ/mol. \[ \Delta H_{vap} = H_{vapor} - H_{liquid}\] \( \Delta H_{vap}\) is the change in enthalpy of vaporization \(H_{vapor}\) is the absolute enthalpy of the gas state of a compound or element \(H_{liquid}\) is the absolute enthalpy of the liquid state of a compound or element Enthalpies is a state function under a specific set of conditions, but cannot be measured directly, but the given difference has a unique value that can be measured.

Because heat is added to the system to maintain the temperature, vaporization is an endothermic process, hence \( \Delta H_{vaporization}\) is always positive. The Enthalpy of Condensation Condensation is the opposite of vaporization, and therefore \( \Delta H_{condensation}\) is also the opposite of \( \Delta H_{vap}\). Because \( \Delta H_{vap}\) is an endothermic process, where heat is lost in a reaction and must be added into the system from the surroundings, \( \Delta H_{condensation}\) is an exothermic process, where heat is absorbed in a reaction and must be given off from the system into the surroundings. \[ΔH_{condensation} = H_{liquid} - H_{vapor} = -ΔH_{vap}\] Because \(ΔH_{condensation}\), also written as \(ΔH_{cond}\), is an exothermic process, its value is always negative. Moreover, \(ΔH_{cond}\) is equal in magnitude to \(ΔH_{vap}\), so the only difference between the two values for one given compound or element is the positive or negative sign.

Example 2.055 liters of steam at 100°C was collected and stored in a cooler container. What was the amount of heat involved in this reaction? The \(ΔH_{vap}\) of water = 44.0 kJ/mol. SOLUTION Liquid to Gas Phase Transition About Watch and Favorite Watch Watching this resources will notify you when proposed changes or new versions are created so you can keep track of improvements that have been made. Favorite Favoriting this resource allows you to save it in the “My Resources” tab of your account. There, you can easily access this resource later when you’re ready to customize it or assign it to your students. Liquid to Gas Phase Transition Vaporization of a sample of liquid is a phase transition from the liquid phase to the gas phase. Learning Objective Describe the process of vaporization. Key Points Evaporation is a phase transition from the liquid phase to the gas phase that occurs at temperatures below the boiling point at a given pressure . For molecules of a liquid to evaporate, they must be located near the surface, be moving in the proper direction, and have sufficient kinetic energy to overcome liquid-phase intermolecular forces .

Boiling is a phase transition from the liquid phase to the gas phase that occurs at or above the boiling temperature. Boiling is the rapid vaporization of a liquid and occurs when a liquid is heated to its boiling point. A liquid's boiling point is the temperature at which the vapor pressure of the liquid is equal to the pressure exerted on the liquid by the surrounding environment (air). Terms Boiling Boiling is the rapid vaporization of a liquid and occurs when a liquid is heated to its boiling point, or the temperature at which the vapor pressure of the liquid is equal to the pressure exerted on the surface of the liquid by the surrounding atmospheric gas (air). Full Text Phase Transition: Liquid to Gas Vaporization of a sample of liquid is a phase transition from the liquid phase to the gas phase. There are two types of vaporization: evaporation and boiling. Evaporation occurs at temperatures below the boiling point, and occurs on the liquid's surface. For molecules of a liquid to evaporate, they must be located near the surface, be moving in the proper direction, and have sufficient kinetic energy to overcome intermolecular forces present in the liquid phase.

Boiling, by contrast, is a rapid vaporization that occurs at or above the boiling temperature and at or below the liquid's surface. Vapor Pressure In much the same way that tea spreads out from a tea bag once the bag is immersed in water, molecules that are confined within a phase will tend to spread themselves (and the thermal energy they carry with them) as widely as possible. This fundamental law of nature is manifested in what is called the "escaping tendency" of the molecules from the phase. The escaping tendency is of fundamental importance in understanding all chemical equilibria and transformations. It is possible to observe the tendency of molecules to escape into the gas phase from a solid or liquid by placing the substance in a closed, evacuated container connected to a manometer for measuring gas pressure. The vapor pressure is a direct measure of the escaping tendency of molecules from a condensed state of matter. If this is done with water, the partial pressure of water Pw in the space above the liquid will initially be zero (step 1). Gradually, Pw will rise as molecules escape from the liquid phase and enter the vapor phase. At the same time, some of the vapor molecules will condense back into the liquid phase (step 2).

Because this latter process is less favorable (at the particular temperature represented here), Pw continues to rise as more water vapor forms. Eventually a balance is reached between the two processes (step 3), and Pw stabilizes at a ﬁxed, or equilibrium, value Pvap , which depends on the substance and temperature. Pvap is known as the "equilibrium vapor pressure", or simply as the "vapor pressure" of the liquid. The vapor pressure is a direct measure of the escaping tendency of molecules from a condensed state of matter . Boiling The boiling point is the temperature at which the vapor pressure of the liquid is equal to the pressure exerted on the liquid by the surrounding environment (air). The escaping tendency of molecules from a phase always increases with temperature; therefore, the vapor pressure of a liquid will be greater at higher temperatures. The variation of vapor pressure with temperature is not linear. Because the typical pressure at the earth's surface is 1 atm (760 torr), this is the vapor pressure that a liquid must equal in order for it to be at its normal boiling point. The variation of vapor pressure with temperature is not linear. The intercepts of each curve with the horizontal line at 1 atm (i.e.

760 torr) indicate the normal boiling point of each liquid, ranging from -25 °C for methyl chloride to over 80 °C for fluorobenzene and 2-heptene. The curve suggests that when the atmospheric pressure is lower than 1 atm (for instance, at higher altitudes), then the boiling point will occur at lower temperatures. This is because the liquid can be heated less in order for its vapor pressure to equal the atmospheric pressure. This has indeed been observed to be true. Related Phenomena Involving The Liquid Boundary Curvature The vapor pressure of a liquid is determined by the attractive forces that act on the molecules at the surface of a liquid. In a very small drop, the liquid surface is curved in such a way that each molecule experiences fewer nearest-neighbor attractions than is the case for the bulk liquid. The outermost molecules of the liquid are bound to the droplet less tightly, and the drop has a larger vapor pressure than does the bulk liquid. If the vapor pressure of the drop is greater than the partial pressure of vapor in the gas phase, the drop will evaporate. Non-polar molecules (gray) evaporate or boil more quickly than polar molecules (blue and red).

Attractions between molecules are shown with dotted lines. Run the model, then heat the liquids. What does boiling look like at the molecular level? A bubble is a hole in a liquid; molecules at the liquid boundary are curved inward, so they experience stronger nearest-neighbor attractions. As a consequence, the vapor pressure Pw of the liquid facing into a bubble is always less than that of the bulk liquid Pw at the same temperature. When the bulk liquid is at its normal boiling point (that is, when its vapor pressure is 1 atm), the pressure of the vapor within the bubble will be less than 1 atm, so the bubble will tend to collapse. Also, since the bubble is formed within the liquid, the hydrostatic pressure of the overlaying liquid will add to this eﬀect. For both of these reasons, a liquid will not boil until the temperature is raised slightly above the boiling point, a phenomenon known as superheating. Once the boiling begins, it will continue to do so at the liquid's proper boiling point. States of matter: introduction 1 Solids, liquids and gases What distinguishes solids, liquids, and gases– the three major states of matter — from each other?

Let us begin at the microscopic level, by reviewing what we know about gases, the simplest state in which matter can exist. At ordinary pressures, the molecules of a gas are so far apart that intermolecular forces have an insignificant effect on the random thermal motions of the individual particles. As the temperature decreases and the pressure increases, intermolecular attractions become more important, and there will be an increasing tendency for molecules to form temporary clusters. These are so short-lived, however, that even under extreme conditions, gases cannot be said to possess “structure” in the usual sense. The contrast at the microscopic level between solids, liquids and gases is most clearly seen in the simplified schematic views above. The molecular units of crystalline solids tend to be highly ordered, with each unit occupying a fixed position with respect to the others. In liquids, the molecules are able to slip around each other, introducing an element of disorder and creating some void spaces that decrease the density. Gases present a picture of almost total disorder, with practically no restrictions on where any one molecule can be.

Solids, liquids and gases: how to tell them apart Having lived our lives in a world composed of solids, liquids, and gases, few of us ever have any difficulty deciding into which of these categories a given sample of matter falls. Our decision is most commonly based on purely visual cues: a gas is transparent and has no definite boundaries other than those that might be imposed by the walls of a confining vessel. Liquids and solids possess a clearly delineated phase boundary that gives solids their definite shapes and whose light-reflecting properties enable us to distinguish one phase from another. Solids can have any conceivable shape, and their surfaces are usually too irregular to show specular (mirror-like) reflection of light. Liquids, on the other hand, are mobile; except when in the form of tiny droplets, liquids have no inherent shape of their own, but assume the shape of their container and show an approximately flat upper surface. Our experience also tells us that these categories are quite distinct; a phase , which you will recall is a region of matter having uniform intensive properties, is either a gas, a liquid, or a solid.

Thus the three states of matter are not simply three points on a continuum; when an ordinary solid melts, it usually does so at a definite temperature, without apparently passing through any states that are intermediate between a solid and a liquid. Although these common-sense perceptions are usually correct, they are not infallible, and in fact there are gases that are not transparent, there are solids such as glasses and many plastics that do not have sharp melting points, but instead undergo a gradual transition from solid to liquid known as softening, and when subject to enough pressure, solids can exhibit something of the flow properties of liquids (glacial ice, for example). Some solids can flow — slowly! Kluane Glacier in Canada's Yukon Territory. Typical glacial flow rates are 10-200 m/yr. More on glacial flow . [ image ] Macroscopic physical properties A more scientific approach would be to compare the macroscopic physical properties of the three states of matter, but even here we run into difficulty.

It is true, for example, that the density of a gas is usually about a thousandth of that of the liquid or solid at the same temperature and pressure; thus one gram of water vapor at 100°C and 1 atm pressure occupies a volume of 1671 mL; when it condenses to liquid water at the same temperature, it occupies only 1.043 mL. solid 13.9 cm3/mol The table at the left compares the molar volumes of neon in its three states. For the gaseous state, P = 1 atm and T = 0°C. The excluded volume is the volume actually taken up by the neon atoms according to the van der Waals model . It is this extreme contrast with the gaseous states that leads to the appellation “ condensed states of matter ” for liquids and solids. However, gases at very high pressures can have densities that exceed those of other solid and liquid substances, so density alone is not a sufficiently comprehensive criterion for distinguishing between the gaseous and condensed states of matter.

Similarly, the density of a solid is usually greater than that of the corresponding liquid at the same temperature and pressure, but not always: you have certainly seen ice floating on water! Problem example 1 Compare the density of gaseous xenon (molar mass 131 g) at 100 atm and 0°C with that of a hydrocarbon liquid for which ρ = 0.104 g/mL at the same temperature. Solution: For simplicity, we will pretend that xenon approximates an ideal gas under these conditions, which it really does not. The ideal molar volume at 0° C and 1 atm is 22.4 L; at 100 atm, this would be reduced to .22 L or 220 mL, giving a density ρ = (131 g) / (224 mL) = 0.58 g/mL. In his autobiographical Uncle Tungsten, the physician/author Oliver Sacks describes his experience with xenon-filled balloons of "astonishing density — as near to 'lead balloons" as could be [imagined].

If one twirled these xenon balloons in one's hand, then stopped, the heavy gas, by its own momentum, would continue rotating for a minute, almost as if it were a liquid." Other physical properties, such as the compressibility , surface tension, and viscosity, are somewhat more useful for distinguishing between the different states of matter. Even these, however, provide no well-defined dividing lines between the various states. Rather than try to develop a strict scheme for classifying the three states of matter, it will be more useful to simply present a few generalizations. If you don't know the meaning of one of the terms in the "property" column in the table, just move your mouse over it to bring up a brief definition. Relative magnitudes of some properties of the three states of matter property thermal expansion coefficient Thermal expansion coefficient large (= R/P) kinetic energy per molecule Kinetic energy per molecule large medium small Some of these deal with macroscopic properties (that is, properties such as the density that relate to bulk matter), and others with microscopic properties that refer to the individual molecular units.

Condensed states of matter Even the most casual inspection of the above table shows that solids and liquids possess an important commonality that distinguishes them from gases: in solids and liquids, the molecules are in contact with their neighbors. As a consequence, these condensed states of matter generally possess much higher densities than gases. Equations of state In our study of gases, we showed that the macroscopic properties of a gas (the pressure, volume, and temperature) are related through an equation of state , and that for the limiting case of an ideal gas, this equation of state can be derived from the relatively small set of assumptions of the kinetic molecular theory. To the extent that a volume of gas consists mostly of empty space, all gases have very similar properties. Equations of state work for gases because gases consist mostly of empty space, so intermolecular interactions can be largely neglected. In condensed matter, these interactions dominate, and they tend to be unique to each particular substance, so there is no such thing as a genrally useful equation of state of liquids and solids. Is there a somewhat more elaborate theory that can predict the behavior of the other two principal states of matter, liquids and solids?

Very simply, the answer is "no"; despite much effort, no one has yet been able to derive a general equation of state for condensed states of matter. The best one can do is to construct models based on the imagined interplay of attractive and repulsive forces, and then test these models by computer simulation. Nevertheless, the very factors that would seem to make an equation of state for liquids and solids impossibly complicated also give rise to new effects that are easily observed, and which ultimately define the distinguishing characteristics of the gaseous, liquid, and solid states of matter. In this unit, we will try to learn something about these distinctions, and how they are affected by the chemical constitution of a substance. 2 Liquids Crystalline solids and gases stand at the two extremes of the spectrum of perfect order and complete chaos. Liquids display elements of both qualities, and both in limited and imperfect ways. Liquids and solids share most of the properties of having their molecular units in direct contact as discussed in the previous section on condensed states of matter. At they same time, liquids, like gases, are fluids, meaning that their molecular units can move more or less independently of each other.

But whereas the volume of a gas depends entirely on the pressure (and thus generally on the volume within which it is confined), the volume of a liquid is largely independent of the pressure. We discuss the properties of liquids in some detail in another lesson . Here we offer just enough to help you see how they relate to the other major states of matter. 3 Solids Of the four ancient elements of "fire, air, earth and water", it is the many forms of solids ("earths") that we encounter in daily life and which give form, color and variety to our visual world. The solid state, being the form of any substance that prevails at lower temperatures, is one in which thermal motion plays an even smaller role than in liquids. The thermal kinetic energy that the individual molecular units do have at temperatures below their melting points allows them to oscillate around a fixed center whose location is determined by the balance between local forces of attraction and repulsion due to neighboring units, but only very rarely will a molecule jump out of the fixed space alloted to it in the lattice. Thus solids, unlike liquids, exhibit long-range order, cohesiveness and rigidity, and possess definite shapes.

Classification of solids Most people who have lived in the world long enough to read this have already developed a rough way of categorizing sollds on the basis of macroscopic properties they can easily observe; everyone knows that a piece of metal is fundamentally different from a rock or a chunk of wood. Unfortunately, nature's ingenuity is far too versatile to fit into any simple system of classifying solids, especially those composed of more than a single chemical substance. Classification according to bond type The most commonly used classification is based on the kinds of forces that join the molecular units of a solid together. We can usually distinguish four major categories on the basis of properties such as general appearance, hardness, and melting point. type of solid low-to-moderate mp, low hardness Classifications should not be taken too seriously! It's important to understand that these four categories are in a sense idealizations that fail to reflect the diversity found in nature. The triangular diagram shown here illustrates this very nicely by displaying examples of binary compounds whose properties suggest that they fall somewhere other than at a vertex of the triangle. For an excellent discussion of diagrams of this kind with many examples, see this page from Mark Leach's Chemogenesis site.

The triangle shown above excludes what is probably the largest category: molecular solids that are bound by van der Waals forces (which are described in the next lesson ). One way of including these is to expand the triangle to a tetrahedron (the so-called Laing tetrahedron). Although this illustrates the concept, it is visually awkward to include many examples of the intermediate cases. Classification by type of molecular unit Solids, like the other states of matter, can be classified according to whether their fundamental molecular units are atoms, electrically-neutral molecules, or ions. But solids possess an additional property that gases and liquids do not: an enduring structural arrangement of their molecular units. Over-simplifying only a bit, we can draw up a rough classification of solids according to the following scheme: structure 4 Crystalline solids In a solid comprised of identical molecular units, the most favored (lowest potential energy) locations occur at regular intervals in space. If each of these locations is actually occupied, the solid is known as a perfect crystal. Much more on crystals and their structures can be found in these units that deal with lattices and external shapes , cubic crystals and packing of spheres , ionic solids , and determination of crystal structures.

What really defines a crystalline solid is that its structure is composed of repeating unit cells each containing a small number of molecular units bearing a fixed geometric relation to one another. The resulting long-range order defines a three-dimensional geometric framework known as a lattice. [ image ] Geometric theory shows that only fourteen different types of lattices are possible in three dimensions, and that just six different unit cell arrangements can generate these lattices. The regularity of the external faces of crystals, which in fact correspond to lattice planes, reflects the long-range order inherent in the underlying structure. Perfection is no more attainable in a crystal than in anything else; real crystals contain defects of various kinds, such as lattice positions that are either vacant or occupied by impurities, or by abrupt displacements or dislocations of the lattice structure. Most pure substances, including the metallic elements, form crystalline solids. But there are some important exceptions. Metallic solids The details of metallic bonding are covered in this Chemical Bonding unit . In metals the valence electrons are free to wander throughout the solid, instead of being localized on one atom and shared with a neighboring one. The valence electrons behave very much like a mobile fluid in which the fixed lattice of atoms is immersed.

This provides the ultimate in electron sharing, and creates a very strong binding effect in solids composed of elements that have the requisite number of electrons in their valence shells. The characteristic physical properties of metals such as their ability to bend and deform without breaking, their high thermal and electrical conductivities and their metallic sheen are all due to the fluid-like behavior of the valence electrons. Molecular solids Recall that a "molecule" is defined as a discrete aggregate of atoms bound together sufficiently tightly (that is, by directed covalent forces) to allow it to retain its individuality when the substance is dissolved, melted, or vaporized. The two words italicized in the preceding sentence are important; covalent bonding implies that the forces acting between atoms within the molecule are much stronger than those acting between molecules, and the directional property of covalent bonding confers on each molecule a distinctive shape which affects a number of its properties. Most compounds of carbon — and therefore, most chemical substances, fall into this category. Many simpler compounds also form molecules; H2O, NH3, CO2, and PCl5 are familiar examples.

Some of the elements, such as H2, O2, O3, P4 and S8 also occur as discrete molecules. Liquids and solids that are composed of molecules are held together by van der Waals forces, and many of their properties reflect this weak binding. Thus molecular solids tend to be soft or deformable, have low melting points, and are often sufficiently volatile to evaporate (sublime) directly into the gas phase; the latter property often gives such solids a distinctive odor. Iodine is a good example of a volatile molecular crystal. The solid (mp 114° C , bp 184°) consists of I2 molecules bound together only by dispersion forces. If you have ever worked with solid iodine in the laboratory, you will probably recall the smell and sight of its purple vapor which is easily seen in a closed container. [images: left , right: Wikimedia commons ] Because dispersion forces and the other van der Waals forces increase with the number of atoms, larger molecules are generally less volatile, and have higher melting points, than do the smaller ones.

Also, as one moves down a column in the periodic table, the outer electrons are more loosely bound to the nucleus, increasing the polarisability polarisability of the atom and thus its susceptibility to van der Waals-type interactions. This effect is particularly apparent in the progression of the boiling points of the successively heavier noble gas elements. Covalent solids These are a class of extended-lattice compounds (see Section 6 below) in which each atom is covalently bonded to its nearest neighbors. This means that the entire crystal is in effect one super-giant “molecule”. The extraordinarily strong binding forces that join all adjacent atoms account for the extreme hardness of such substances; these solids cannot be broken or abraded without cleaving a large number of covalent chemical bonds. Similarly, a covalent solid cannot “melt” in the usual sense, since the entire crystal is its own giant molecule. When heated to very high temperatures, these solids usually decompose into their elements. ↑ diamond lattice structure [ source ] The Hard Materials Web site has a lot of interesting information on some of the substances described in this section.

Diamond Diamond is the hardest material known, defining the upper end of the 1-10 scale known as Moh's hardness . Diamond cannot be melted; above 1700°C it is converted to graphite , the more stable form of carbon. The diamond unit cell is face-centered cubic and contains 8 carbon atoms. The four darkly shaded ones are contained within the cell and are completely bonded to other members of the cell. The other carbon atoms (6 in faces and 4 at corners) have some bonds that extend to atoms in other cells. (Two of the carbons nearest the viewer are shown as open circles in order to more clearly reveal the bonding arrangement.) Wikipedia has a very good article on the the occurrence, properties, and uses of diamonds. Other covalent solids Boron nitride BN is similar to carbon in that it exists as a diamond-like cubic polymorph polymorph as well as in a hexagonal form analogous to graphite. Cubic BN is the second hardest material after diamond, and finds use in industrial abrasives and cutting tools. Recent interest in BN has centered on its carbon-like ability to form nanotubes and related nanostructures.

( A-Zmaterials article , Nanotechnology article , Wikipedia article ) Silicon carbide is an extremely rare mineral on the earth, and comes mostly from meteorites which are believed to have their origins in carbonaceous stars. The first synthetic SiC was made accidently by E.G. Acheson in 1891 who immediately recognized its industrial prospects and founded the Carborundum Co. Silicon carbide SiC is also known as carborundum. Its structure is very much like that of diamond with every other carbon replaced by silicon. On heating at atmospheric pressure, it decomposes at 2700°C, but has never been observed to melt. Structurally, it is very complex; at least 70 crystalline forms have been identified. Its extreme hardness and ease of synthesis have led to a diversity of applications — in cutting tools and abrasives, high-temperature semiconductors, and other high-temperature applications, manufacture of specialty steels, jewelry, and many more. For much more about the history, properties and uses of this versatile material, see this Wikipedia article .

Tungsten carbide WC is probably the most widely-encountered covalent solid owing to its use in "carbide" cutting tools and as the material used to make the rotating balls in ball-point pens . It's high-melting (2870°C) form has a structure similar to that of diamond and is only slightly less hard. In many of its applications it is embedded in a softer matrix of cobalt or coated with titanium compounds. ( Wikipedia article ) 5 Amorphous solids In some solids there is so little long-range order that the substance cannot be considered crystalline at all; such a solid is said to be amorphous. Amorphous solids possess short-range order but are devoid of any organized structure over longer distances; in this respect they resemble liquids. However, their rigidity and cohesiveness allow them to retain a definite shape, so for most practical purposes they can be considered to be solids. Glasses This term refers generally to solids formed from their melts that do not return to their crystalline forms on cooling, but instead form hard, and often transparent amorphous solids.

Although some organic substances such as sugar can form glasses ("rock candy"), the term more commonly describes inorganic compounds, especially those based on silica, SiO2. Natural silica-based glasses, known as obsidian , are formed when certain volcanic magmas cool rapidly. The difference between crystalline silica and silica glass is shown in these simplified two-dimensional projections. It is readily apparent that much of the ordered arrangement of crystalline silica is lost in the glassy form. But the glass retains enough Si–O bonds to form a hard, rigid material. Ordinary glassis composed mostly of SiO2, which usually exists in nature in a crystalline form known as quartz. If quartz (in the form of sand) is melted and allowed to cool, it becomes so viscous that the molecules are unable to move to the low potential energy positions they would occupy in the crystal lattice, so that the disorder present in the liquid gets “frozen into” the solid. In a sense, glass can be regarded as a supercooled liquid. Glasses are transparent because the distances over which disorder appears are small compared to the wavelength of visible light, so there is nothing to scatter the light and produce cloudiness.

Some excellent references on glass and its composition: Wikipedia article Answers.com article Ordinary glass is made by melting silica sand to which has been added some calcium and sodium carbonates. These additives reduce the melting point and make it more difficult for the SiO2 molecules to arrange themselves into crystalline order as the mixture cools. See here for a brief discussion of this important topic. Glass is believed to have first been made in the Middle East at least as early as 3000 BCE. Its workability and ease of coloring has made it one of mankind's most important and versatile materials. But even after all these years, "The Nature of Glass Remains Anything but Clear", according to this interesting 2008 NY Times article . Layer or sheet-like structures Solid cadmium chloride is a good example of a layer structure. The Cd and Cl atoms occupy separate layers; each of these layers extends out in a third dimension to form a sheet. The CdCl2 crystal is built up from stacks of these layers held together by van der Waals forces. Each Cd atom is covalently bonded to two chlorine atoms in adjacent layers; all other "bonds" are weak van der Waals attractions.

It's worth pointing out that although salts such as CuCl2 and CdCl2 are dissociated into ions when in aqueous solution, the solids themselves should not be regarded as "ionic solids". See also this section of the lesson on ionic solids. Graphite Graphite is a polymorph polymorph of carbon and its most stable form. It consists of sheets of fused benzene rings stacked in layers. The spacing between layers is sufficient to admit molecules of water vapor and other atmospheric gases which become absorbed in the interlamellar spaces and act as lubricants, allowing the layers to slip along each other. Thus graphite itself often has a flake-like character and is commonly used as a solid lubricant, although it loses this property in a vacuum. As would be expected from its anisotropic anisotropic structure, the electric and thermal conductivity of graphite are much greater in directions parallel to the layers than across the layers. The melting point of 4700-5000°C makes graphite useful as a high-temperature refractory material. Graphite is the most common form of relatively pure carbon found in nature.

Its name comes from the same root as the Greek word for "write" or "draw", reflecting its use as pencil "lead" since the 16th century. (The misnomer, which survives in common use, is due to its mis-identification as an ore of the metallic element of the same name at a time long before modern chemistry had developed.) Some interesting Wikipedia pages: Graphite ; Pencils, their history and manufacture Graphene Graphene is a two-dimensional material consisting of a single layer of graphite — essentially "chicken wire made of carbon" that was discovered in 2004. Small fragments of graphene can be obtained by several methods; one is to attach a piece of Scotch Tape™ to a piece of graphite and then carefully pull it off (a process known as exfoliation.) Fragments of graphene are probably produced whenever one writes with a pencil. Graphene has properties that are uniquely different from all other solids. It is the strongest known material, and it exhibits extremely high electrical conductivity due to its massless electrons which are apparently able to travel at relativistic velocities through the layer.

Some excellent articles on graphene: Scientific American (April 2007) - Wikipedia - Lawrence Berkeley Laboratory What you should be able to do Make sure you thoroughly understand the following essential ideas which have been presented above. It is especially imortant that you know the precise meanings of all the green-highlighted terms in the context of this topic. State the major feature that characterizes a condensed state of matter. Describe some of the major observable properties that distinguish gases, liquid and solids, and state their relative magnitudes in these three states of matter. Describe the dominant forces and the resulting physical properties that distinguish ionic, covalent, metallic, and molecular solids. Explain the difference between

Question: What is the name of the transition from a liquid state directly to a gas of an element or compound?
Answer: Vaporisation