A molecule is defined as an electrically neutral group of at least two atoms held together by strong chemical bonds. Molecules can be made of a single element or different elements. The modern understanding of molecules evolved from early theories proposed by scientists like Dalton, Avogadro, and Cannizaro in the 19th century. A key concept is that equal volumes of gases contain equal numbers of molecules. Molecules have a definite mass and structure that determines their properties.

1. MOLECULE

2. MOLECULE A molecule is defined as an electrically neutral group of at least two atoms in a definite arrangement held together by very strong (covalent) chemical bonds.[1][2] Molecules are distinguished from polyatomic ions in this strict sense. In organic chemistry and biochemistry, the term molecule is used less strictly and also is applied to charged organic molecules and biomolecules. A molecule may consist of atoms of a single chemical element, as with oxygen (O2), or of different elements, as with water (H2O). Atoms and complexes connected by non-covalent bonds such as hydrogen bonds or ionic bonds are generally not considered single molecules.

3. Evolution of Molecular Theory The terms atom and molecule were used interchangeably until the early 19th cent. Initial experimental work with gases led to what is essentially the modern distinction. J. A. C. Charles and R. Boyle had shown that all gases exhibit the same relationship between a change in temperature or pressure and the corresponding change in volume. J. L. Gay-Lussac had shown that gases always combine in simple whole-number volume proportions and had rediscovered the earlier findings of Charles, which had not been published. Dalton's Theory One early theorist was John Dalton, best known for his atomic theory. Dalton believed that gases were made up of tiny particles, which he thought were atoms. He thought that these atoms were stationary and in contact with one another and that heat was a material substance, called caloric, that was contained in shells around the atom (these shells of caloric were actually what was in contact). When a gas was heated, the amount of caloric was increased, the shells became larger, and the gas expanded. Dalton did not accept Gay-Lussac's findings about combining volumes of gases, perhaps because it could not be explained by his theory.

4. Avogadro's Hypothesis A different theory from Dalton's that could explain the combining volumes of gases was proposed by the Italian physicist Amadeo Avogadro in 1811. According to his theory, under given conditions of temperature and pressure, a given volume of any gas contains a definite number of particles. From the earlier observation that one volume of hydrogen gas and one volume of chlorine gas react to form two volumes of hydrogen chloride gas he deduced that the particles in gaseous hydrogen or chlorine could not be single atoms, but must be some combination of atoms. He called this combination a molecule. Cannizaro's Compromise Many shortcomings of Dalton's theory were uncovered, and although a number of modifications were suggested, none were very successful. It was not until 1858 that the Italian chemist StanislaoCannizaro suggested a merging of Avogadro's and Dalton's theories. The acceptance of this revised theory was assisted by the acceptance by physicists at about the same time of the kinetic-molecular theory of gases that was first proposed in 1738 by Daniel Bernoulli.

5. CONCEPT Prior to the nineteenth century, chemists pursued science simply by taking measurements, before and after a chemical reaction, of the substances involved. This was an external approach, rather like a person reaching into a box and feeling of the contents without actually being able to see them. With the evolution of atomic theory, chemistry took on much greater definition: for the first time, chemists understood that the materials with which they worked were interacting on a level much too small to see. The effects, of course, could be witnessed, but the activities themselves involved the interactions of atoms in molecules. Just as an atom is the most basic particle of an element, a molecule is the basic particle of a compound. Whereas there are only about 90 elements that occur in nature, many millions of compounds are formed naturally or artificially. Hence the study of the molecule is at least as important to the pursuit of modern chemistry as the study of the atom. Among the most important subjects in chemistry are the ways in which atoms join to form molecules—not just the numbers and types of atoms involved, but the shape that they form together in the molecular structure.

6. HOW IT WORKS? Sucrose or common table sugar, of course, is grainy and sweet, yet it is made of three elements that share none of those characteristics. The formula for sugar is C12H22O11, meaning that each molecule is formed by the joining of 12 carbon atoms, 22 hydrogens, and 11 atoms of oxygen. Coal is nothing like sugar—for one thing, it is as black as sugar is white, yet it is almost pure carbon. Carbon, at least, is a solid at room temperature, like sugar. The other two components of sugar, on the other hand, are gases, and highly flammable ones at that. The question of how elements react to one another, producing compounds that are altogether unlike the constituent parts, is one of the most fascinating aspects of chemistry and, indeed, of science in general. Combined in other ways and in other proportions, the elements in sugar could become water (H2O), carbon dioxide (CO2), or even petroleum, which is formed by the joining of carbon and hydrogen.

7. MOLECULAR STRUCTURE It is not enough, however, to know that a certain combination of atoms forms a certain molecule, because molecules may have identical formulas and yet be quite different substances. In English, for instance, there is the word "rose." Simply seeing the word, however, does not tell us whether it is a noun, referring to a flower, or a verb, as in "she rose through the ranks." Similarly, the formula of a compound does not necessarily tell what it is, and this can be crucial. For instance, the formula C2H6O identifies two very different substances. One of these is ethyl alcohol, the type of alcohol found in beer and wine. Note that the elements involved are the same as those in sugar, though the proportions are different: in fact, some aspects of the body's reaction to ethyl alcohol are not so different from its response to sugar, since both lead to unhealthy weight gain. In reasonable small quantities, of course, ethyl alcohol is not toxic, or at least only mildly so; yet methyl ether—which has an identical formula—is a toxin.

8. MOLECULES AND COMPOUND A molecule can be most properly defined as a group of atoms joined in a specific structure. A compound, on the other hand, is a substance made up of more than one type of atom—in other words, more than one type of element. Not all compounds are composed of discrete molecules, however. For instance, table salt (NaCl) is an ionic compound formed by endlessly repeating clusters of sodium and chlorine that are not, in the strictest sense of the word, molecules. Salt is an example of a crystalline solid, or a solid in which the constituent parts are arranged in a simple, definite geometric pattern repeated in all directions. There are three kinds of crystalline solids, only one of which has a truly molecular structure. In an ionic solid such as table salt, ions (atoms, or groups of atoms, with an electric charge) bond a metal to a nonmetal—in this case, the metal sodium and the nonmetal chlorine. Another type of crystalline solid, an atomic solid, is formed by atoms of one element bonding to one another. A diamond, made of pure carbon, is an example. Only the third type of crystalline solid is truly molecular in structure: a molecular solid—sugar, for example—is one in which the molecules have a neutral electric charge.

9. MOLECULAR MASS Just as the atoms of elements have a definite mass, so do molecules—a mass equal to that of the combined atoms in the molecule. The figures for the atomic mass of all elements are established, and can be found on the periodic table; therefore, when one knows the mass of a hydrogen atom and an oxygen atom, as well as the fact that there are two hydrogens and one oxygen in a molecule of water, it is easy to calculate the mass of a water molecule. BONDING BETWEEN MOLECULES Note that the mass of an atom in a molecule does not change; nor, indeed, do the identities of the individual atoms. An oxygen atom in water is the same oxygen atom in sugar, or in any number of other compounds. With regard to compounds, it should be noted that these are not the same thing as a mixture, or a solution. Sugar or salt can be dissolved in water at the appropriate temperatures, but the resulting solution is not a compound; the substances are joined physically, but they are not chemically bonded.

10. Chemical Bonding Chemical bonding is the joining, through electromagnetic force, of atoms representing different elements. Each atom possesses a certain valency, which determines its ability to bond with atoms of other elements. Valency, in turn, is governed by the configuration of valence electrons at the highest energy level (the shell) of the atom.

11. Electronegativity Not all elements bond covalently in the same way. Each has a certain value of electronegativity—the relative ability of an atom to attract valence electrons. Elements capable of bonding are assigned an electronegativity value ranging from a minimum of 0.7 for cesium to a maximum of 4.0 for fluorine. The greater the electronegativity value, the greater the tendency of an element to attract valence electrons. Attractions Between Molecules The energy required to pull apart a molecule is known as bond energy. Covalent bonds that involve hydrogen are among the weakest bonds between atoms, and hence it is relatively easy to separate water into its constituent parts, hydrogen and oxygen. (This is sometimes done by electrolysis, which involves the use of an electric current to disperse atoms.) Double and triple covalent bonds are stronger, but strongest of all is an ionic bond. The strength of the bond energy in salt, for instance, is reflected by its melting point of 1,472°F (800°C), much higher than that of water, at 32°F (0°C).

12. Nature of Molecules Molecules are made up of two or more atoms, either of the same element or of two or more different elements, joined by one or more covalent chemical bonds. According to the kinetic-molecular theory, the molecules of a substance are in constant motion. The state (solid, liquid, or gaseous) in which matter appears depends on the speed and separation of the molecules in the matter. Substances differ according to the structure and composition of their molecules. A molecular compound is represented by its molecular formula; for example, water is represented by the formula H2O. A more complex structural formula is sometimes used to show the arrangement of atoms in the molecule. Molecules differ in size and molecular weight as well as in structure. In a chemical reaction between molecular substances, the molecules are often broken apart into atoms or radicals that recombine to form other molecules, i.e., other substances. In other cases two or more molecules will combine to form a single larger molecule, or a large molecule will be broken up into several smaller molecules.

13. COLLOIDS

14. COLLOIDS Colloids were originally defined by Thomas Graham in 1861 as substances, such as starch or gelatin, which will not diffuse through a membrane. He distinguished them from crystalloids (e.g. inorganic salts), which would pass through membranes. Later it was recognized that colloids were distinguished from true solutions by the presence of particles that were too small to be observed with a normal microscope yet were much larger than normal molecules. Colloids are now regarded as systems in which there are two or more phases, with one (the dispersed phase) distributed in the other (the continuous phase). Moreover, at least one of the phases has small dimensions (in the range 10−9–10−6 m).

16. Types of colloids Colloids are common in everyday life. Some examples include whipped cream, mayonnaise, milk, butter, gelatin, jelly, muddy water, plaster, colored glass, and paper.

17. Parts of Colloids Colloidal particleis a small amount of matter having size typical for colloids and with a clear phase boundary (phase colloids). A group of such particles (aggregate, agglomerate) or being a macromolecule (eg. solution of polymermolecules is a molecular colloid) or a molecular aggregate (e.g. micelle). Dispersing medium is the substance in which the colloidal particles are distributed. In muddy water, for example, the colloidal particles are tiny grains of sand, silt, and clay. The dispersing medium is the water in which these particles are suspended.

19. Properties of Colloids Each type of mixture has special properties by which it can be identified. For example, a suspension always settles out after a certain period of time. That is, the particles that make up the suspension separate from the medium in which they are suspended and fall to the bottom of a container. In contrast, colloidal particles typically do not settle out. Like the particles in a solution, they remain in suspension within the medium that contains them. Colloids also exhibit Brownian movement. Brownian movement is the random zigzag motion of particles that can be seen under a microscope. The motion is caused by the collision of molecules with colloid particles in the dispersing medium. In addition, colloids display the Tyndall effect. When a strong light is shone through a colloidal dispersion, the light beam becomes visible, like a column of light. A common example of this effect can be seen when a spotlight is turned on during a foggy night. You can see the spotlight beam because of the fuzzy trace it makes in the fog (a colloid).

20. Light shining through a solution of sodium hydroxide (left) and a colloidal mixture. The size of colloidal particles makes the mixture, which is neither a solution nor a suspension, appear cloudy.

21. Interaction between colloid particles The following forces play an important role in the interaction of colloid particles: Excluded volume repulsion: This refers to the impossibility of any overlap between hard particles. Electrostatic interaction: Colloidal particles often carry an electrical charge and therefore attract or repel each other. The charge of both the continuous and the dispersed phase, as well as the mobility of the phases are factors affecting this interaction. van der Waals forces: This is due to interaction between two dipoles that are either permanent or induced. Even if the particles do not have a permanent dipole, fluctuations of the electron density gives rise to a temporary dipole in a particle. This temporary dipole induces a dipole in particles nearby. The temporary dipole and the induced dipoles are then attracted to each other. This is known as van der Waals force, and is always present (unless the refractive indexes of the dispersed and continuous phases are matched), is short-range, and is attractive.

22. Entropic forces: According to the second law of thermodynamics, a system progresses to a state in which entropy is maximized. This can result in effective forces even between hard spheres. Steric forcesbetween polymer-covered surfaces or in solutions containing non-adsorbing polymer can modulate interparticle forces, producing an additional steric repulsive force (which is predominantly entropic in origin) or an attractive depletion force between them. Such an effect is specifically searched for with tailor-made superplasticizers developed to increase the workability of concrete and to reduce its water content.

23. GAS

24. GAS Gas is one of three classical states of matter.Near absolute zero, a substance exists as a solid. As heat is added to this substance it melts into a liquid at its melting point (see phase change), boils into a gas at its boiling point, and if heated high enough would enter a plasma state in which the electrons are so energized that they leave their parent atoms from within the gas. A pure gas may be made up of individual atoms (e.g. a noble gas or atomic gas like neon), elemental molecules made from one type of atom (e.g. oxygen), or compound molecules made from a variety of atoms (e.g. carbon dioxide). A gas mixture would contain a variety of pure gases much like the air. What distinguishes a gas from liquids and solids is the vast separation of the individual gas particles. This separation usually makes a colorless gas invisible to the human observer. The interaction of gas particles in the presence of electric and gravitational fields are considered negligible. The gaseous state of matter is found between the liquid and plasma states[2], the latter of which provides the upper temperature boundary for gases. Bounding the lower end of the temperature scale lie degenerative quantum gases[3] which are gaining increased attention these days.[4] High-density atomic gases super cooled to incredibly low temperatures are classified by their statistical behavior as either a Bose gas or a Fermi gas.

25. Physical characteristics Drifting smoke particles provide clues to the movement of the surrounding gas. As most gases are difficult to observe directly with our senses, they are described through the use of four physical properties or macroscopic characteristics: the gas’s pressure, volume, number of particles (chemists group them by moles), and temperature. These four characteristics were repeatedly observed by men such as Robert Boyle, Jacques Charles, John Dalton, Joseph Gay-Lussac and Amedeo Avogadro for a variety of gases in a great many settings. Their detailed studies ultimately led to a mathematical relationship among these properties expressed by the ideal gas law. Gas particles are widely separated from one another, and as such do not influence adjacent particles to the same degree as liquids or solids. This influence (intermolecular forces) results from the magnetic charges that these gas particles carry. Like charges repel, while oppositely charged particles attract one another. Gases made from ions carry permanent charges, as do compounds with their polar covalent bonds. These polar covalent bonds produce permanent charge concentrations within the molecule while the compound's net charge remains neutral.

26. Compared to the other states of matter, gases have an incredibly low density and viscosity. Pressure and temperature influence the particles within a certain volume. This variation in particle separation and speed is referred to as compressibility. This particle separation and size influences optical properties of gases as can be found in the following list of refractive indices. Finally, gas particles spread apart or diffuse in order to homogeneously distribute themselves throughout any container. Drifting smoke particles provide clues to the movement of the surrounding gas.

27. MACROSCOPIC When observing a gas, it is typical to specify a frame of reference or length scale. A larger length scale corresponds to a macroscopic or global point of view of the gas. This region (referred to as a volume) must be sufficient in size to contain a large sampling of gas particles. The resulting statistical analysis of this sample size produces the "average" behavior (i.e. velocity, temperature or pressure) of all the gas particles within the region. By way of contrast, a smaller length scale corresponds to a microscopic or particle point of view. Pressure The symbol used to represent pressure in equations is "p" or "P" with SI units of pascals. When describing a container of gas, the term pressure (or absolute pressure) refers to the average force the gas exerts on the surface area of the container. Within this volume, it is sometimes easier to visualize the gas particles moving in straight lines until they collide with the container (see diagram at top of the article). The force imparted by a gas particle into the container during this collision is the change in momentum of the particle.

28. Temperature The symbol used to represent temperature in equations is T with SI units of kelvins. The speed of a gas particle is proportional to its absolute temperature. The volume of the balloon in the image to the right shrinks when the trapped gas particles slow down with the addition of extremely cold nitrogen. The temperature of any physical system is related to the motions of the particles (molecules and atoms) which make up the [gas] system.[ Specific Volume The symbol used to represent specific volume in equations is "v" with SI units of cubic meters per kilogram. The symbol used to represent volume in equations is "V" with SI units of cubic meters. When performing a thermodynamic analysis, it is typical to speak of intensive and extensive properties. Properties which depend on the amount of gas (either by mass or volume) are called extensive properties, while properties that do not depend on the amount of gas are called intensive properties.

29. Specific volume is an example of an intensive property because it is the ratio of volume occupied by a unit of mass of a gas that is identical throughout a system at equilibrium.[11] 1000 atoms of protactinium as a gas occupy the same space as any other 1000 atoms for any given temperature and pressure. This concept is easier to visualize for solids such as iron which are incompressible compared to gases. When the seat ejection is initiated in the image above the specific volume increases with the expanding gases, while mass is conserved. Since a gas fills any container in which it is placed, volume is an extensive property. Density The symbol used to represent density in equations is ρ (pronounced rho) with SI units of kilograms per cubic meter. This term is the reciprocal of specific volume. Since gas molecules can move freely within a container, their mass is normally characterized by density. Density is the mass per volume of a substance or simply, the inverse of specific volume. For gases, the density can vary over a wide range because the particles are free to move closer together when constrained by pressure or volume or both. This variation of density is referred to as compressibility. Like pressure and temperature, density is a state variable of a gas and the change in density during any process is governed by the laws of thermodynamics

30. GAS LAW English chemist Robert Boyle (1627-1691), who made a number of important contributions to chemistry—including his definition and identification of elements—seems to have been influenced by Torricelli. If so, this is an interesting example of ideas passing from one great thinker to another: Torricelli, a student of Galileo Galilei (1564-1642), was no doubt influenced by Galileo's thermoscope. The gas laws are not a set of government regulations concerning use of heating fuel; rather, they are a series of statements concerning the behavior of gases in response to changes in temperature, pressure, and volume. These were derived, beginning with Boyle's law, during the seventeenth, eighteenth, and nineteenth centuries by scientists whose work is commemorated through the association of their names with the laws they discovered. In addition to Boyle, these men include fellow English chemists John Dalton (1766-1844) and William Henry (1774-1836); French physicists and chemists J. A. C. Charles (1746-1823) and Joseph Gay-Lussac (1778-1850); and Italian physicist Amedeo Avogadro (1776-1856).

31. BOYLE’S AND CHARLES LAW Boyle's law holds that in isothermal conditions (that is, a situation in which temperature is kept constant), an inverse relationship exists between the volume and pressure of a gas. (An inverse relationship is a situation involving two variables, in which one of the two increases in direct proportion to the decrease in the other.) In this case, the greater the pressure, the less the volume and vice versa. Therefore, the product of the volume multiplied by the pressure remains constant in all circumstances. Charles's law also yields a constant, but in this case the temperature and volume are allowed to vary under isobarometric conditions—that is, a situation in which the pressure remains the same. As gas heats up, its volume increases, and when it cools down, its volume reduces accordingly. Hence, Charles established that the ratio of temperature to volume is constant.

32. GAY LUSSAC’S LAW From Boyle's and Charles's law, a pattern should be emerging: both treat one parameter (temperature in Boyle's, pressure in Charles's) as unvarying, while two other factors are treated as variables. Both, in turn, yield relationships between the two variables: in Boyle's law, pressure and volume are inversely related, whereas in Charles's law, temperature and volume are directly related. In Gay-Lussac's law, a third parameter, volume, is treated as a constant, and the result is a constant ratio between the variables of pressure and temperature. According to Gay-Lussac's law, the pressure of a gas is directly related to its absolute temperature. AVOGADRO’S LAW In 1811, Amedeo Avogadro verified that equal volumes of pure gases contain the same number of particles. His theory was not generally accepted until 1858 when another Italian chemist StanislaoCannizzaro was able to explain non-ideal exceptions. For his work with gases a century prior, the number that bears his name Avogadro's constant represents the number of atoms found in 12 grams of elemental carbon-12 (6.022×1023 mol-1). This specific number of gas particles, at standard temperature and pressure (ideal gas law) occupies 22.40 liters and is referred to as the molar volume.

33. DALTON’S LAW In 1801, John Dalton published the Law of Partial Pressures from his work with ideal gas law relationship: The pressure of a mixture of gases is equal to the sum of the pressures of all of the constituent gases alone. Mathematically, this can be represented for n species as: Pressuretotal = Pressure1 + Pressure2 + ... + Pressuren Among his key journal observations upon mixing unreactive "elastic fluids" (gases) were the following.[20]: Unlike liquids, heavier gases did not drift to the bottom upon mixing. Gas particle identity played no role in determining final pressure (they behaved as if their size was negligible).

34. KINETIC THEORY OF GASES From the preceding gas laws, a set of propositions known collectively as the kinetic theory of gases has been derived. Collectively, these put forth the proposition that a gas consists of numerous molecules, relatively far apart in space, which interact by colliding. These collisions are responsible for the production of thermal energy, because when the velocity of the molecules increases—as it does after collision—the temperature increases as well. There are five basic postulates to the kinetic theory of gases: Gases consist of tiny molecular or atomic particles. The proportion between the size of these particles and the distances between them is so small that the individual particles can be assumed to have negligible volume. These particles experience continual random motion. When placed in a container, their collisions with the walls of the container constitute the pressure exerted by the gas. The particles neither attract nor repel one another. The average kinetic energy of the particles in a gas is directly related to absolute temperature.