physical chemistry

Physical chemistry is the study of macroscopic, atomic, subatomic, and particulate phenomena in chemical systems in terms of laws and concepts of physics. It applies the principles, practices and concepts of physics such as motion, energy, force, time, thermodynamics, quantum chemistry, statistical mechanics and dynamics.
Physical chemistry, in contrast to chemical physics, is predominantly (but not always) 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, chemical equilibrium, and colloids.
Some of the relationships that physical chemistry strives to resolve include the effects of:

1. Intermolecular forces that act upon 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.
4. Surface chemistry and electrochemistry of membranes.[1]
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    Key concepts

   The key concepts of physical chemistry are the ways in which pure physics is applied to chemical problems.
   One of the key concepts in chemistry is that all chemical compounds can be described as groups of atoms bonded together and chemical reactions can be described as the making and breaking of those bonds. Predicting the properties of chemical compounds from a description of atoms and how they bond is one of the major goals of physical chemistry. To describe the atoms and bonds precisely, it is necessary to know both where the nuclei of the atoms are, and how electrons are distributed around them.^[2] Quantum chemistry, a subfield of physical chemistry especially concerned with the application of quantum mechanics to chemical problems, provides tools to determine how strong and what shape bonds are,^[2] how nuclei move, and how light can be absorbed or emitted by a chemical compound.^[3] Spectroscopy is the related sub-discipline of physical chemistry which is specifically concerned with the interaction of electromagnetic radiation with matter.
   Another set of important questions in chemistry concerns what kind of reactions can happen spontaneously and which properties are possible for a given chemical mixture. This is studied in chemical thermodynamics, which sets limits on quantities like how far a reaction can proceed, or how much energy can be converted into work in a combustion engine and which provides links between properties like the thermal expansion coefficient and change of rate in entropy with pressure for a gas or a liquid.^[4] It can frequently be used to assess whether a reactor or engine design is feasible, or to check the validity of experimental data. To a limited extent, quasi-equilibrium and non-equilibrium thermodynamics can describe irreversible changes.^[5] However, classical thermodynamics is mostly concerned with systems in equilibrium and reversible changes and not what actually does happen, or how fast, away from equilibrium.
   Which reactions do occur and how fast is the subject of chemical kinetics, another branch of physical chemistry. A key idea in chemical kinetics is that for reactants to react and form products, most chemical species must go through transition states which are higher in energy than either the reactants or the products and serve as a barrier to reaction.^[6] In general, the higher the barrier, the slower the reaction. A second is that most chemical reactions occur as a sequence of elementary reactions,^[7] each with its own transition state. Key questions in kinetics include how the rate of reaction depends on temperature and on the concentrations of reactants and catalysts in the reaction mixture, as well as how catalysts and reaction conditions can be engineered to optimize the reaction rate.
   The fact that how fast reactions occur can often be specified with just a few concentrations and a temperature, instead of needing to know all the positions and speeds of every molecule in a mixture, is a special case of another key concept in physical chemistry, which is that to the extent an engineer needs to know, everything going on in a mixture of millions of billions of billions of particles can often be described by just a few variables like pressure, temperature, and concentration. The precise reasons for this are described in statistical mechanics,^[8] a specialty within physical chemistry which is also shared with physics. Statistical mechanics also provides ways to predict the properties we see in everyday life from molecular properties without relying on empirical correlations based on chemical similarities.^[5]
   

    History

   See also: History of chemistry

   
   

   

   Fragment of M. Lomonosov's manuscript 'Physical Chemistry' (1752)

   The term "physical chemistry" was coined 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.^[9]
   Modern physical chemistry originated in the 1860s to 1880s with 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. This paper introduced several of the cornerstones of physical chemistry, such as Gibbs energy, chemical potentials, Gibbs phase rule.^[10] Other milestones include the subsequent naming and accreditation of enthalpy to Heike Kamerlingh Onnes and to macromolecular processes.^{[citation needed]}
   The first scientific journal 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. Together with Svante August Arrhenius,^[11] these were the leading figures in physical chemistry in the late 19th century and early 20th century. All three were awarded with the Nobel Prize in Chemistry between 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 World War II), more recent discoveries in astrochemistry,^[12] as well as the 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, Lydersen method, Joback method, Benson group increment theory, QSPR, QSAR

   Journals

   Some journals that deal with physical chemistry include:

   • Zeitschrift für Physikalische Chemie (1887)
   • Journal of Physical Chemistry A (from 1896 as Journal of Physical Chemistry, renamed in 1997)
   • Physical Chemistry Chemical Physics (from 1999, formerly Faraday Transactions with a history dating back to 1905)
   • Macromolecular Chemistry and Physics (1947)
   • Annual Review of Physical Chemistry (1950)
   • Molecular Physics (1957)
   • Journal of Physical Organic Chemistry (1988)
   • Journal of Physical Chemistry B (1997)
   • ChemPhysChem (2000)
   • Journal of Physical Chemistry C (2007)
   • Journal of Physical Chemistry Letters (from 2010, combined letters previously published in the separate journals)

   Historical journals that covered both chemistry and physics include Annales de chimie et de physique (started in 1789, published under the name given here from 1815–1914).Branches and related topics
   

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

    See also

   Physical chemistry portal

   • List of important publications in chemistry#Physical chemistry
   • List of unsolved problems in chemistry#Physical chemistry problems
   • Category:Physical chemists

   References

   1. ^ Torben Smith Sørensen (1999). Surface chemistry and electrochemistry of membranes. CRC Press. p. 134. ISBN 0-8247-1922-0. http://books.google.co.in/books?id=v1oU13xAl6AC&pg=RA1-PA134. 
   2. ^ ^{a b} Atkins, Peter and Friedman, Ronald (2005). Molecular Quantum Mechanics, p. 249. Oxford University Press, New York. ISBN 0-19-927498-3.
   3. ^ Atkins, Peter and Friedman, Ronald (2005). Molecular Quantum Mechanics, p. 342. Oxford University Press, New York. ISBN 0-19-927498-3.
   4. ^ Landau, L. D. and Lifshitz, E. M. (1980). Statistical Physics, 3rd Ed. p. 52. Elsevier Butterworth Heinemann, New York. ISBN 0-7506-3372-7.
   5. ^ ^{a b} Hill, Terrell L. (1986). Introduction to Statistical Thermodynamics, p. 1. Dover Publications, New York. ISBN 0-486-65242-4.
   6. ^ Schmidt, Lanny D. (2005). The Engineering of Chemical Reactions, 2nd Ed. p. 30. Oxford University Press, New York. ISBN 0-19-516925-5.
   7. ^ Schmidt, Lanny D. (2005). The Engineering of Chemical Reactions, 2nd Ed. p. 25, 32. Oxford University Press, New York. ISBN 0-19-516925-5.
   8. ^ Chandler, David (1987). Introduction to Modern Statistical Mechanics, p. 54. Oxford University Press, New York. ISBN 978-0-19-504277-1.
   9. ^ Alexander Vucinich (1963). Science in Russian culture. Stanford University Press. p. 388. ISBN 0-8047-0738-3. http://books.google.com/books?id=YoE1wsA6USQC&pg=PA388. 
   10. ^ Josiah Willard Gibbs, 1876, "On the Equilibrium of Heterogeneous Substances", Transactions of the Connecticut Academy of Sciences
   11. ^ Laidler, Keith (1993). The World of Physical Chemistry. Oxford: Oxford University Press. pp. 48. ISBN 0-19-855919-4. 
   12. ^ Herbst, Eric (May 12, 2005). "Chemistry of Star-Forming Regions". Journal of Physical Chemistry A 109 (18): 4017–4029. doi:10.1021/jp050461c. PMID 16833724. 

   External links

   Wikimedia Commons has media related to: Physical chemistry

   Wikibooks has more on the topic of: Physical chemistry

   At Wikiversity you can learn more and teach others about Physical chemistry at: The Department of Physical chemistry

   • Physical Chemistry (Keith J. Laidler, John H. Meiser and Bryan C. Sanctuary)
   • The World of Physical Chemistry (Keith J. Laidler, 1993)
   • Physical Chemistry from Ostwald to Pauling (John W. Servos, 1996)
   • 100 Years of Physical Chemistry (Royal Society of Chemistry, 2004)
   • Physical Chemistry: neither Fish nor Fowl? (Joachim Schummer, The Autonomy of Chemistry, Würzburg, Königshausen & Neumann, 1998, pp. 135–148)
   • Cathedrals of Science (Patrick Coffey, 2008)
   • The Cambridge History of Science: The modern physical and

   C2

Chemical kinetics

Reaction rate tends to increase with concentration - a phenomenon explained by collision theory.

Chemical kinetics, also known as reaction kinetics, is the study of rates of chemical processes. Chemical kinetics includes investigations of how different experimental conditions can influence the speed of a chemical reaction and yield information about the reaction's mechanism and transition states, as well as the construction of mathematical models that can describe the characteristics of a chemical reaction. In 1864, Peter Waage and Cato Guldberg pioneered the development of chemical kinetics by formulating the law of mass action, which states that the speed of a chemical reaction is proportional to the quantity of the reacting substances.
Chemical kinetics deals with the experimental determination of reaction rates from which rate laws and rate constants are derived. Relatively simple rate laws exist for zero-order reactions (for which reaction rates are independent of concentration), first-order reactions, and second-order reactions, and can be derived for others. In consecutive reactions, the rate-determining step often determines the kinetics. In consecutive first-order reactions, a steady state approximation can simplify the rate law. The activation energy for a reaction is experimentally determined through the Arrhenius equation and the Eyring equation. The main factors that influence the reaction rate include: the physical state of the reactants, the concentrations of the reactants, the temperature at which the reaction occurs, and whether or not any catalysts are present in the reaction.

[edit] Factors affecting reaction rate

[edit] Nature of the reactants

Depending upon what substances are reacting, the reaction rate varies. Acid/base reactions, the formation of salts, and ion exchange are fast reactions. When covalent bond formation takes place between the molecules and when large molecules are formed, the reactions tend to be very slow. Nature and strength of bonds in reactant molecules greatly influence the rate of its transformation into products.

[edit] Physical state

The physical state (solid, liquid, or gas) of a reactant is also an important factor of the rate of change. When reactants are in the same phase, as in aqueous solution, thermal motion brings them into contact. However, when they are in different phases, the reaction is limited to the interface between the reactants. Reaction can occur only at their area of contact; in the case of a liquid and a gas, at the surface of the liquid. Vigorous shaking and stirring may be needed to bring the reaction to completion. This means that the more finely divided a solid or liquid reactant the greater its surface area per unit volume and the more contact it makes with the other reactant, thus the faster the reaction. To make an analogy, for example, when one starts a fire, one uses wood chips and small branches — one does not start with large logs right away. In organic chemistry, on water reactions are the exception to the rule that homogeneous reactions take place faster than heterogeneous reactions.

 Concentration

Concentration plays a very important role in reactions, because, according to the collision theory of chemical reactions, molecules must collide in order to react together. As the concentration of the reactants increases, the frequency of the molecules colliding increases, striking each other more frequently by being in closer contact at any given point in time. One may think of two reactants being in a closed container: All the molecules contained within are colliding constantly. By increasing the amount of one or more of the reactants, these collisions happen more often, increasing the reaction rate.

Temperature

Temperature usually has a major effect on the rate of a chemical reaction. Molecules at a higher temperature have more thermal energy. Although collision frequency is greater at higher temperatures, this alone contributes only a very small proportion to the increase in rate of reaction. Much more important is the fact that the proportion of reactant molecules with sufficient energy to react (energy greater than activation energy: E > E_a) is significantly higher and is explained in detail by the Maxwell–Boltzmann distribution of molecular energies.
The 'rule of thumb' that the rate of chemical reactions doubles for every 10 °C temperature rise is a common misconception. This may have been generalized from the special case of biological systems, where the α (temperature coefficient) is often between 1.5 and 2.5.
A reaction's kinetics can also be studied with a temperature jump approach. This involves using a sharp rise in temperature and observing the relaxation time of the return to equilibrium.

 Catalysts

Generic potential energy diagram showing the effect of a catalyst in a hypothetical endothermic chemical reaction. The presence of the catalyst opens a different reaction pathway (shown in red) with a lower activation energy. The final result and the overall thermodynamics are the same.

A catalyst is a substance that accelerates the rate of a chemical reaction but remains chemically unchanged afterwards. The catalyst increases rate reaction by providing a different reaction mechanism to occur with a lower activation energy. In autocatalysis a reaction product is itself a catalyst for that reaction leading to positive feedback. Proteins that act as catalysts in biochemical reactions are called enzymes. Michaelis-Menten kinetics describe the rate of enzyme mediated reactions. A catalyst does not affect the position of the equilibria, as the catalyst speeds up the backward and forward reactions equally.
In certain organic molecules, specific substituents can have an influence on reaction rate in neighbouring group participation.
Agitating or mixing a solution will also accelerate the rate of a chemical reaction, as this gives the particles greater kinetic energy, increasing the number of collisions between reactants and, therefore, the possibility of successful collisions.

pressure Increasing the pressure in a gaseous reaction will increase the number of collisions between reactants, increasing the rate of reaction. This is because the activity of a gas is directly proportional to the partial pressure of the gas. This is similar to the effect of increasing the concentration of a solution.
A reaction's kinetics can also be studied with a pressure jump approach. This involves making fast changes in pressure and observing the relaxation time of the return to equilibrium.

Equilibrium
While chemical kinetics is concerned with the rate of a chemical reaction, thermodynamics determines the extent to which reactions occur. In a reversible reaction, chemical equilibrium is reached when the rates of the forward and reverse reactions are equal and the concentrations of the reactants and products no longer change. This is demonstrated by, for example, the Haber–Bosch process for combining nitrogen and hydrogen to produce ammonia. Chemical clock reactions such as the Belousov–Zhabotinsky reaction demonstrate that component concentrations can oscillate for a long time before finally attaining the equilibrium.

Free energy

In general terms, the free energy change (ΔG) of a reaction determines whether a chemical change will take place, but kinetics describes how fast the reaction is. A reaction can be very exothermic and have a very positive entropy change but will not happen in practice if the reaction is too slow. If a reactant can produce two different products, the thermodynamically most stable one will in general form, except in special circumstances when the reaction is said to be under kinetic reaction control. The Curtin–Hammett principle applies when determining the product ratio for two reactants interconverting rapidly, each going to a different product. It is possible to make predictions about reaction rate constants for a reaction from free-energy relationships.
The kinetic isotope effect is the difference in the rate of a chemical reaction when an atom in one of the reactants is replaced by one of its isotopes.
Chemical kinetics provides information on residence time and heat transfer in a chemical reactor in chemical engineering and the molar mass distribution in polymer chemistry.

Applications

The mathematical models that describe chemical reaction kinetics provide chemists and chemical engineers with tools to better understand and describe chemical processes such as food decomposition, microorganism growth, stratospheric ozone decomposition, and the complex chemistry of biological systems. These models can also be used in the design or modification of chemical reactors to optimize product yield, more efficiently separate products, and eliminate environmentally harmful by-products. When performing catalytic cracking of heavy hydrocarbons into gasoline and light gas, for example, kinetic models can be used to find the temperature and pressure at which the highest yield of heavy hydrocarbons into gasoline will occur. Kinetics is also a basic aspect of chemistry.

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Reaction mechanism

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In chemistry, a reaction mechanism is the step by step sequence of elementary reactions by which overall chemical change occurs.^[1]
Although only the net chemical change is directly observable for most chemical reactions, experiments can often be designed that suggest the possible sequence of steps in a reaction mechanism. Recently, electrospray ionization mass spectrometry^[2] has been used to corroborate the mechanism of several organic reaction proposals.

Description

A chemical mechanism describes in detail exactly what takes place at each stage of an overall chemical reaction (transformation). It also describes each reactive intermediate, activated complex, and transition state, and which bonds are broken (and in what order), and which bonds are formed (and in what order). A complete mechanism must also account for all reactants used, the function of a catalyst, stereochemistry, all products formed and the amount of each. It must also describe the relative rates of the reaction steps and the rate equation for the overall reaction. Reaction intermediates are chemical species, often unstable and short-lived, which are not reactants or products of the overall chemical reaction, but are temporary products and reactants in the mechanism's reaction steps. Reaction intermediates are often free radicals or ions. Transition states can be unstable intermediate molecular states even in the elementary reactions. Transition states are commonly molecular entities involving an unstable number of bonds and/or unstable geometry. They correspond to maxima on the reaction coordinate, and to saddle points on the potential energy surface for the reaction.

S_N2 reaction mechanism. Note the negatively-charged transition state in brackets in which the central carbon atom in question shows five bonds, an unstable condition.

The electron or arrow pushing method is often used in illustrating a reaction mechanism; for example, see the illustration of the mechanism for benzoin condensation in the following examples section.
A reaction mechanism must also account for the order in which molecules react. Often what appears to be a single step conversion is in fact a multistep reaction.

 Examples

Consider the following reaction:

CO + NO₂ → CO₂ + NO

In this case, it has been experimentally determined that this reaction takes place according to the rate law . This form suggests that the rate-determining step is a reaction between two molecules of NO₂. A possible mechanism for the overall reaction which explains the rate law is:

2 NO₂ → NO₃ + NO (slow)

NO₃ + CO → NO₂ + CO₂ (fast)

Each step is called an elementary step, and each has its own rate law and molecularity. The elementary steps should add up to the original reaction.
When determining the overall rate law for a reaction, the slowest step is the step that determines the reaction rate. Because the first step (in the above reaction) is the slowest step, it is the rate-determining step. Because it involves the collision of two NO₂ molecules, it is a bimolecular reaction with a rate law of . If we were to cancel out all the molecules that appear on both sides of the reaction, we would be left with the original reaction.
Other reactions may have mechanisms of several consecutive steps. In organic chemistry, one of the first reaction mechanisms proposed was that for the benzoin condensation, put forward in 1903 by A. J. Lapworth.

Benzoin condensation reaction mechanism. Cyanide ion (CN^-) acts as a catalyst here, entering at the first step and leaving in the last step. Proton (H⁺) transfers occur at (i) and (ii). The arrow pushing method is used in some of the steps to show where electron pairs go.

There are also more complex mechanisms such as chain reactions, in which the propagation steps of the chain form a closed cycle.

 Modeling

A correct reaction mechanism is an important part of accurate predictive modeling. For many combustion and plasma systems, detailed mechanisms are not available or require development.
Even when information is available, identifying and assembling the relevant data from a variety of sources, reconciling discrepant values and extrapolating to different conditions can be a difficult process without expert help. Rate constants or thermochemical data are often not available in the literature, so computational chemistry techniques or group-additivity methods must be used to obtain the required parameters.
At the different stages of a reaction mechanism's elaboration, appropriate methods must be used.

 Molecularity

Main article: molecularity

Molecularity in chemistry is the number of colliding molecular entities that are involved in a single reaction step.

• A reaction involving one molecular entity is called unimolecular.
• A reaction involving two molecular entities is called bimolecular.
• A reaction involving three molecular entities is called termolecular.

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PHYSICAL CHEMISTRY MENU Kinetic Theory Kinetic Theory . . . Basic kinetic theory ideas about solids, liquids and gases, and changes of state. Ideal and real gases. The ideal gas equation. Boyle's Law and Charles' Law. Chemical energetics Enthalpy changes during reactions . . . An explanation of the various important kinds of enthalpy change, and a limited look at the calculations which go with them. Rates of Reaction Basic ideas about rates of reaction . . . Covers collision theory, and describes and explains the effects of surface area, concentration, pressure, temperature and catalysts on reaction rates. Also takes an introductory look at the rate equation for a reaction - particularly order of reaction and the rate constant. Catalysis . . . A detailed look at various types of catalytic action with a wide range of examples taken from A level syllabuses. Equilibria Chemical equilibria . . . 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Instrumental analysis Explains how you can analyse substances using machines - mass spectrometry, infra-red spectroscopy, NMR, UV-visible absorption spectrometry and chromatography. Basic Organic Chemistry Includes help on bonding, naming and isomerism, and a discussion of organic acids and bases. Properties of organic compounds Covers the physical and chemical properties of compounds on UK A level chemistry syllabuses, and includes a limited amount of biochemistry. Organic Reaction Mechanisms Covers all the mechanisms required by the current UK A level chemistry syllabuses. About this site Includes a contact address if you have found any difficulties with the site. Questions and comments A selection of questions that I have been asked lots of times about Chemguide together with a few general comments. There are also a number of chemistry questions that I have been asked and which I haven't been able to find good answers for! Chemistry Calculations A description of the author's book on calculations at UK A level chemistry standard. Textbook suggestions Suggestions for textbooks and revision guides covering the UK AS and A level chemistry syllabuses, with links to Amazon.co.uk if you want to follow them up. Download syllabuses For UK students and international students using UK exams (e.g. Cambridge International). Download a copy of your current syllabus from your examiners. Links A random collection of links to sites that I have found interesting or useful. You will find it is a fairly quirky collection - that's deliberate.