physical chemistry

Physical chemistry (also called physicochemistry) is the explanation of macroscopic, microscopic, atomic, subatomic, and particulate phenomena in chemical systems in terms of physical concepts; sometimes using the principles, practices and concepts of physics like thermodynamics, quantum chemistry, statistical mechanics and dynamics.^[1]

Physical chemistry in contrast to chemical physics is still predominantly a macroscopic or supra-molecular science, as the majority of the principles on which physical chemistry was founded are concepts related to the bulk rather than on molecular/atomic structure alone; for example, colloids.^[2], chemical equilibrium etc. Some of the relationships that physical chemistry has lately tried to resolve include the effects of:

1. Intermolecular forces on the physical properties of materials (plasticity, tensile strength, surface tension in liquids).
2. Reaction kinetics on the rate of a reaction.
3. The identity of ions on the electrical conductivity of materials.

There is a large body of knowledge embodied in this very old branch of chemistry that can not be explained just by using the concepts of physics.

History

The term "physical chemistry" was probably first introduced by Mikhail Lomonosov in 1752, when he presented a lecture course entitled "A Course in True Physical Chemistry" (Russian: «Курс истинной физической химии») before the students of Petersburg University.

The foundation of modern physical chemistry is thought to have been laid in the 1860s to 1880s by work on chemical thermodynamics, electrolytes in solutions, chemical kinetics and other subjects. One milestone was the publication in 1876 by Josiah Willard Gibbs of his paper, On the Equilibrium of Heterogeneous Substances, which contained several of the cornerstones of physical chemistry, such as Gibbs energy, chemical potentials, Gibbs phase rule ^[3] and subsequent naming and accreditation of enthalpy to Heike Kamerlingh Onnes and to macromolecular processes.^{[citation needed]}

The first scientific journal for publications specifically in the field of physical chemistry was the German journal, Zeitschrift für physikalische Chemie, founded in 1887 by Wilhelm Ostwald and Jacobus Henricus van 't Hoff, which were two of the other leading figures of physical chemistry in the late 19th century and early 20th century together with Svante August Arrhenius. All three were awarded with the Nobel Prize in Chemistry in the period 1901-1909.

Developments in the following decades include the application of statistical mechanics to chemical systems and work on colloids and surface chemistry, where Irving Langmuir made many contributions. Another important step was the development of quantum mechanics into quantum chemistry from the 1930s, where Linus Pauling was one of the leading names. Theoretical developments have gone hand in hand with developments in experimental methods, where the use of different forms of spectroscopy, such as infrared spectroscopy, microwave spectroscopy, EPR spectroscopy and NMR spectroscopy, is probably the most important 20th century development.

Further development in physical chemistry may be attributed to discoveries in nuclear chemistry, especially in isotope separation (before and during WW II), and after, by development of calculation algorithms in field of "additive physicochemical properties" (practically all of physicochemical properties, as: boiling point, critical point, surface tension, vapor pressure etc. - more than 20 in all, can be precisely calculated from chemical structure, even if such chemical molecule is still non existent), and in this area is concentrated practical importance of contemporary physical chemistry. -> see, Group contribution method, Joback method

Properties of gases

Kinetic theory

One charateristic of physical chemistry (much like other branches of physicla science) is that, in order to develop workable theories to explain the natural world, models are adopted for each particular system under consideration. One example is the kinetic theory of gases, whereby a gas is treated as a collection of independent particles in ceaseless, chaotic motion. Thus, kinetic theory attempts to explain the macroscopic properties of gases, such as pressure, temperature, or volume, by considering the motion of their primary particles (atoms or molecules).

Essentially, the theory assumes that pressure is due to collisions between molecules moving at different velocities, while temperature is a macroscopic assessment of the average speed or velocities of the particles. Thus, pressure arises from the force exerted by particles (atoms or molecules) impacting on the walls of the container.

Clearly, the more particles that are present in the container, the greater the number of collisions with the walls of the container, and thus the greater the pressure. However, adding more particles will limit the motion of the particles since their paths (or trajectories) will be limited due the interference of collisions. I.E. They will move just as fast, but not as far. For example, due the interference of collisions, mixing will take longer when more particles are present.

The Gas Laws

Boyle's Law

Boyle’s Law states that when the pressure of a gas increases, then the volume of the gas decreases proportionally. In other words, the product of the pressure and volume is fixed or constant, or

PV = constant

Or if the subscript 1 represents the initial conditions, and 2 represents the final conditions, then the following relationship holds:

P₁V₁=P₂ V₂

For this law to apply, the temperature must remain fixed, and must always be given in Kelvins, in order to use these formulas.

Charles' Law

Charles’ Law states that If you increase the temperature, then the volume will increase proportionally. In other words, the volume of a sample divided by the temperature is equal to a constant, or:

V / T = constant

Another way of expressing this relationship is by stating that:

V₁ / T₁ = V₂ / T₂

Again, T must always be expressed in Kelvins. Also, for this law to apply, the pressure must remain fixed.

Gay Lusaac’s Law

Gay Lusaac’s Law states that the pressure is proportional to the temperature. Thus, the ratio of the pressure to the absolute temperature is fixed or constant.

P / T = constant

Another way of expressing this relationship is by stating that:

P₁ / T₁ = P₂ / T₂

Again, T must always be expressed in Kelvins. Also, for this law to apply, the volume must remain fixed.

Combined Gas Law

We can combine these laws to get a combined gas law, which states that:

PV / T = constant

or

P₁ x ( V₁ / T₁ ) = P₂ x ( V₂ / T₂ )

for a fixed amount of gas (or a fixed number of moles, n ).

For example, let’s start with 2.37 liters of a gas @ 25.0 degrees Celsius (298 K) and 1 atmosphere. Then heat the gas to 297 degrees Celsius (570 K). Finally, increase the pressure to 10 atmospheres. What is the final volume?

Note that while heating it, the volume is going to go up. But in compressing it, the volume is going to go down. So there are two forces working against each other. We start with 2.37 liters of a gas @ 25.0 degrees Celsius (298 K) and 1 atmosphere. We heat it to 297 degrees Celsius (570 K). At the same time, we’re going to increase the pressure to 10 atmospheres.

In the next step, heat is added. Then, rearranging the combined gas law algebraically, and isolating the variable V₂ yields the following relationship:

V₂ = P₁ x V₁ / T₁ x T₂ / P₂. This same expression could just as easily be written as a product of pressure ratios and temperature ratios as follows:

V₂ = V₁ x ( P₁ / P₂ ) x ( T₂ / T₁ )

V2 = (2.37)(1 / 10)(570 / 298)

Note that the ratio of pressures is less than 1 and the ratio of temps is greater than 1, so that they offset each other. But 1 / 10 is a much greater decrease than 1.91 is an increase. Thus, the volume will decrease by (0.1)(1.91) = ( 0.19 )

V2 = (2.37)(0.19) = 0.453L

Ideal Gas Law

The Ideal Gas Law states that:

PV = nRT

where n = # of moles of gas, and the Gas Constant R is given by 0.0821 liter atmospheres per mole Kelvins

In other words, PV / nT is a constant value, so that the following relationship holds:

(P₁ x V₁) / (n₁ x T₁) = (P₂ x V₂ ) / (n₂ x T₂)

So how can we use the Ideal Gas Law? One of the things that we can do is we can calculate the volume of 1 mole of gas. Let’s calculate it at 0.00 degrees C and 1.00 atmosphere. The reason why I chose 0.00 degrees C and 1.00 atmosphere is because chemists have decided to give these two conditions a special name. We call that special set of conditions STP for standard temperature and pressure. And standard temperature and pressure is just 1.00 atmospheres and 0.00 degrees C ( = 273 K ).

Let: P V = n R T. Then rearranging algebraically and isolating the variable yields the following relationship:

V = nRT / P = 22.4liters

This figure represents the volume of a mole of any gas at STP ( about the size of a standard toilet water tank).

Dalton's Law

The pressure exerted by a mixture of ideal gases is equal to the sum of the pressures exerted by the individual gases occupying the same volume alone. For example:

An 11.2 liter tank of gas is found in the coldest part of the refrigerator (at standard temperature, 0ºC = 273 K). It contains 4 moles of gas: 1 mole of oxygen and 3 moles of neon. What is the pressure in the tank? For an ideal gas:

P = nRT / V = 7.91atm

Note that the molar volume of all gases (including oxygen, O₂ and neon, Ne) are equal.

Molar quantities

Another significant premise of physical chemisty relates to the measurement of chemcial quantities. Since the number of atoms contained in 1 gram of matter of any element or compound is on the order of 10²² to 10²³, it is necessary to implement a new standard of measurement (like the dozen) which places a specific large number of particles (atoms or molcules) within one unit. That unit is expressed numnerically by Avogadro's number, which is equal to 6.02 x 10²³. The formal definition of a mole is that it is the amount of substance that ocntains as many particles as there are atoms in exactly 12.0 grams of Carbon-12.

A distinction is made in chemistry between extensive properties and intensive properties. The former depend on the size (or the extent) of the sample, while the latter are independent of it. Examples of intensive properties are temperature, density and pressure. Molar quantities such as the molar mass M (mass per mole) and the molar volume V_M are also intensive because the property and the amount are both proportional to the size of the sample, so that the ratio is independent of its extent.

Concentration

It is often necessary to consider the amounts of matter present in chemical solutions of a wide range of types and compositions. The molar concentration of a a solute in solution is equal ot the amount of substance of the solute (given in moles om illimoles) divded by the volume of solution. Molar concentrations are typicall ymeasured in molarity, or moles per liter, M, such a 1 M aqueous solution of hydrochloric acid contains 1 mole of HCl in a total solution volume of 1 liter. The term molality refers to the number of moles of solute per mass amount of solvent used to prepare the solution. The units are typitcally expressed in moles per kilogram (mol/kg).

Energy

The basic concept at the core of all theories and explanations in the field of physical chemistry is the same as that of many others: energy, which. Energy can be loosely defined as the capacity to do work. We shall often make user of the universal law of nature that energy can be neither created nor destroyed; but rather it is conserved. Thus, the total amount of energy availabe in the universe at any given point in time is fixed, or constant. (We will ultimately speak of open and closed systems).

Kinetic energy

The kinitic energy of a body (or particle) is the energy which it possesses aa result of its motion -- whether that motion be associated with linear velocity, angular velocity, rotational velocity, or vibrational frequency. For a body of mass m traveling at a sped v, the kinetic enegy is detrmeined as follows:

K.E. = 1/2 mv²

In other words, a heavy body traveling rapidly has a high kinetic energy. A static (or stationary) body has no kinetic energy.

Potential energy

The potential energy of a body is the energy it possesses as a result of its position. Consider, for example, the potential energy of a body of mass m in the gravitational field of the earth. If a body is at a height h above hte earth's surface, then its potential energy is given as follows:

P.E. = mgh

where the gravitational constant g is given as follows:

g = 9.81 m / s²

Electric field potential

In the interactions between electrons, atomic nuclei, and electrically charged ions (cations and ions), one must consider the potential energy of a charged body in the vicintiy of another charged body. Note that there that there will be an electrical field acting on the charges carried by any two charged bodies or particles. If a particle of charge q₁ is at a distance r from another particle of charge q₂, then their Coulomb potential energy is expressed as folloows:

V = q₁ x q₂ / 4πr

Note that V = 0 at the point of infinite separation.

Spectrum of light

The energy of an electromaganetic field figures prominently in physical chemistry due to the intimate relationship between the structure of the atom, the nature of the chemical bond, and the absorption and emission of light over a wide range of frequencies and wavelengths -- from the ultraviolet (UV) portion to the visible (or white light) portion into the far Inrared (IR) and microwave portions of the spectrum. Not only does the analytical technique of pectroscopy depend critically on the nature of light, but osme species acquire the energy they need to participate in vital biochemical reactions via photochemistry (or photosynthesis).

An electromagetic field consits of two components: an electric field that acts on elctrically charged particles anda magnetic field which acts only on charged particles which are in motion. Each field produces a force sufficient to accelerate the particle on which it is acting. An electromagnetic filed is generated when the charged particles are in motion. That is the basic principle of a radio transmitter, where electrons are moved back and forth in an antenna, and hence generates an electromagnetic disturbance that propagates through space. Another example is illustrated by the massive amounts of ferrous iron stored deep within the core of the earth, which is in constant rotation about its central axis, generating a geomagnetic field about the surface of the earth.

The alternating field can be characterized as a sine wave whose wavelength (or inverse frequency) is givne by the distance between the neighboring peaks (or valleys) of the wave. The energy of the filed is directly proportional to the amplitude of the eletromagnetic disturbance -- just as the intensity of a sound wave is determined by its amplitude. The disturbance propagates as a wave through empoty space, or a vacuum, at a fixed speed called the speed of light, which is about 3 x 10⁸ meters per sec (m/s).

Emission spectra

When the electrons in an atom or element are excited, the additional energy pushes the electrons to higher energy orbitals whic hare futrhte fomr the aomtic nucleus (i.e. outer shells). When the electrons leave the excited state, this energy is re-emitted in the form of a photon. The wavelength (or, equivalently, frequency) of the photon is determined by the difference in energy between the two states. These emitted photons form the element's emission spectrum. The fact that only certain colors appear in an element's atomic emission spectrum means that only certain frequencies (or wavelengths) of light are emitted. Each of these frequencies are related to energy by the formula:

E_photon = hv

where E is the energy of the photon, v is its frequency, and h is Planck's constant. Thus, only photons having specific quanta (or amounts) of energy are emitted by the atom. (These discrete energy levels form the basis of quantum mechanics).

Emission spectrum of Hydrogen

Emission spectrum of Iron

The emission spectrum of a chemical element or chemical compound is the relative intensity of each frequency of electromagnetic radiation emitted by the element's atoms or the compound's molecules when they are returned to a ground state. Each element's emission spectrum is unique. Therefore, spectroscopy can be used to the identify the elements in matter of unknown composition. Similarly, the emission spectra of molecules can be used in forensic investigations for identification, chemical analysis and determination of chemical composition.

Molecular radiation

As well as the electronic transitions discussed above, the energy of a molecule can also change via rotational, vibrational and vibronic (combined vibrational and electronic) transitions. These energy transitions often lead to closely-spaced groups of many different spectral lines, known as spectral bands. Unresolved band spectra may appear as a spectral continuum.

Molecular emission is the mechanism behind the sulfur lamp and the deuterium arc lamp.

Notes

1. ^ Quantum Chemistry (p3 - "PHYSICAL CHEMISTRY"), states that "We can divide physical chemistry into four areas: thermodynamics, quantum chemistry, statistical mechanics and kinetics".
2. ^ Physical Chemistry of Macromolecules (p1 - "INTRODUCTION"), defines the formation of physical chemistry as being between macromolecules and colloids in modern physical chemistry. Also defines the "fierce battles" in the 1900s between the inclusion of colloids AS macromolecules.
3. ^ Josiah Willard Gibbs, 1876, "On the Equilibrium of Heterogeneous Substances", Transactions of the Connecticut Academy of Sciences

References

1. Levine, I. N. (1978). Physical Chemistry McGraw-Hill publishing ISBN 0-07-037418-X
2. Atkins, P.W. (1978). Physical Chemistry Oxford University Press ISBN 0-7167-3539-X
3. Berry, S. R., Rice, S. A, Ross, J. (2000). Physical Chemistry 2nd ed. Oxford University Press. ISBN 0-19-510589-3
4. Hunter, R. J. (1993) Introduction to Modern Colloid Science Oxford University Press. ISBN 0-19-855386-2
5. Hiemenz, P. C., Rajagopalan, R., (1997). Principles of Colloid and Surface Chemistry Marcel Dekker Inc., New York. ISBN 0-8247-9397-8
6. Moore, W.J. (1963). Physical Chemistry 4th ed. Longman publishers/London/Prentice Hall, NJ.

See also

Physical chemistry portal

Sub-topics

• Photochemistry
• Thermochemistry
• Chemical kinetics
• Quantum chemistry
• Electrochemistry
• Surface chemistry
• Solid-state chemistry
• Spectroscopy
• Biophysical chemistry
• Materials science
• Physical organic chemistry
• Membrane formation
• Micromeritics

Publications

• Important publications in physical chemistry(chemistry),
• Important publications in physical chemistry(physics)