Relevant bibliographies by topics / Raoult's law

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Academic literature on the topic 'Raoult's law'

Author: Grafiati
Published: 4 June 2021
Last updated: 14 February 2022

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Journal articles on the topic "Raoult's law"

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Hermans, J. J. "Raoult's law as a limiting law." Recueil des Travaux Chimiques des Pays-Bas 60, no. 5 (September 3, 2010): 370–72. http://dx.doi.org/10.1002/recl.19410600508.

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Hawkes, Stephen J. "Raoult's Law Is a Deception." Journal of Chemical Education 72, no. 3 (March 1995): 204. http://dx.doi.org/10.1021/ed072p204.

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Kovac, Jeffrey. "Molecular size and Raoult's Law." Journal of Chemical Education 62, no. 12 (December 1985): 1090. http://dx.doi.org/10.1021/ed062p1090.

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Wilson, Archie S. "A visual demonstration of Raoult's law." Journal of Chemical Education 67, no. 7 (July 1990): 598. http://dx.doi.org/10.1021/ed067p598.2.

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Hawkes, Stephen J. "Strategic Consequences from Errors in Raoult's Law Paper." Journal of Chemical Education 73, no. 1 (January 1996): 41. http://dx.doi.org/10.1021/ed073p41.

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YIN, Xuebo. "Relationship and Difference among Raoult's Law, Henry's Law, and Azeotropic Solution." University Chemistry 33, no. 5 (2018): 61–65. http://dx.doi.org/10.3866/pku.dxhx201801010.

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Ferrero Vallana, Federico M., Ricardo P. Girling, H. Q. Nimal Gunaratne, Lynette A. M. Holland, Pauline M. Mcnamee, Kenneth R. Seddon, Jonathan R. Stonehouse, and Oreste Todini. "Ionic liquids as modulators of fragrance release in consumer goods." New Journal of Chemistry 40, no. 12 (2016): 9958–67. http://dx.doi.org/10.1039/c6nj01626j.

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Bowler, Michael G., David R. Bowler, and Matthew W. Bowler. "Raoult's law revisited: accurately predicting equilibrium relative humidity points for humidity control experiments." Journal of Applied Crystallography 50, no. 2 (March 29, 2017): 631–38. http://dx.doi.org/10.1107/s1600576717003636.

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The humidity surrounding a sample is an important variable in scientific experiments. Biological samples in particular require not just a humid atmosphere but often a relative humidity (RH) that is in equilibrium with a stabilizing solution required to maintain the sample in the same state during measurements. The controlled dehydration of macromolecular crystals can lead to significant increases in crystal order, leading to higher diffraction quality. Devices that can accurately control the humidity surrounding crystals while monitoring diffraction have led to this technique being increasingly adopted, as the experiments become easier and more reproducible. Matching the RH to the mother liquor is the first step in allowing the stable mounting of a crystal. In previous work [Wheeler, Russi, Bowler & Bowler (2012). Acta Cryst. F68, 111–114], the equilibrium RHs were measured for a range of concentrations of the most commonly used precipitants in macromolecular crystallography and it was shown how these related to Raoult's law for the equilibrium vapour pressure of water above a solution. However, a discrepancy between the measured values and those predicted by theory could not be explained. Here, a more precise humidity control device has been used to determine equilibrium RH points. The new results are in agreement with Raoult's law. A simple argument in statistical mechanics is also presented, demonstrating that the equilibrium vapour pressure of a solvent is proportional to its mole fraction in an ideal solution: Raoult's law. The same argument can be extended to the case where the solvent and solute molecules are of different sizes, as is the case with polymers. The results provide a framework for the correct maintenance of the RH surrounding a sample.
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González, Juan Antonio, Isaiés Garcié de la Fuentá, and Jose Carlos Cobos. "Thermodynamics of mixtures with strongly negative deviations from Raoult's Law." Fluid Phase Equilibria 168, no. 1 (February 2000): 31–58. http://dx.doi.org/10.1016/s0378-3812(99)00326-x.

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Berka, Ladislav H., and Nicholas Kildahl. "Experiments for Modern Introductory Chemistry: Intermolecular Forces and Raoult's Law." Journal of Chemical Education 71, no. 7 (July 1994): 613. http://dx.doi.org/10.1021/ed071p613.

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Books on the topic "Raoult's law"

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Sherwood, Dennis, and Paul Dalby. Reactions in solution. Oxford University Press, 2018. http://dx.doi.org/10.1093/oso/9780198782957.003.0016.

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Another key chapter, examining reactions in solution. Starting with the definition of an ideal solution, and then introducing Raoult’s law and Henry’s law, this chapter then draws on the results of Chapter 14 (gas phase equilibria) to derive the corresponding results for equilibria in an ideal solution. A unique feature of this chapter is the analysis of coupled reactions, once again using first principles to show how the coupling of an endergonic reaction to a suitable exergonic reaction results in an equilibrium mixture in which the products of the endergonic reaction are present in much higher quantity. This demonstrates how coupled reactions can cause entropy-reducing events to take place without breaking the Second Law, so setting the scene for the future chapters on applications of thermodynamics to the life sciences, especially chapter 24 on bioenergetics.
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Book chapters on the topic "Raoult's law"

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Gooch, Jan W. "Raoult’s Law." In Encyclopedic Dictionary of Polymers, 609. New York, NY: Springer New York, 2011. http://dx.doi.org/10.1007/978-1-4419-6247-8_9765.

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"Raoult's law." In Encyclopedic Dictionary of Polymers, 817. New York, NY: Springer New York, 2007. http://dx.doi.org/10.1007/978-0-387-30160-0_9585.

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Anderson, Greg M., and David A. Crerar. "Aqueous Electrolyte Solutions." In Thermodynamics in Geochemistry. Oxford University Press, 1993. http://dx.doi.org/10.1093/oso/9780195064643.003.0021.

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In dealing with the thermodynamic properties of ions we have one difficulty in addition to those encountered in dealing with compounds and elements. For compounds and elements we found that although we could measure absolute values for some properties, others such as enthalpy and the other energy terms contained an undetermined constant. We got around this by using the concept of "formation from the elements." It would of course be very convenient to also have thermodynamic properties of individual ions, but because positively and negatively charged ions cannot be separated from each other to any significant extent, their individual properties cannot be measured. To get around this, we need an additional convention, while retaining the formation from the elements convention. In addition we have certain problems in dealing with the activities and activity coefficients of electrolytes and individual ions. In the following section we discuss the problems of activities of ionic species. We follow the presentation of Klotz (1964), and include the HC1 example used by Pitzer and Brewer (1961), and an expanded consideration of the choice of solute components. Following that we discuss the conventions used to obtain numerical values for the state variables of individual ions. We begin by demonstrating that the basic approach is not arbitrarily chosen by chemists with a view to confusing students, nor is it dictated by the electrically charged nature of ions. It is dictated by the algebraic consequence of the fact that when neutral solute molecules dissociate into charged particles, the number of solute particles is increased. For example, when one mole of the undissociated solute AB(aq), which can be treated using Henry's Law, Raoult's Law, and the rest of the equations developed in previous chapters, becomes instead one mole of A(aq) plus one mole of B(aq), certain consequences develop that have nothing to do with whether A(aq) and B(aq) are electrically charged or not.
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Oriakhi, Christopher O. "Ideal Solutions and Colligative Properties." In Chemistry in Quantitative Language. Oxford University Press, 2009. http://dx.doi.org/10.1093/oso/9780195367997.003.0019.

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Colligative properties of solutions are those that depend only on the number of solute particles (molecules or ions) in the solution rather than on their chemical or physical properties. The colligative properties that can be measured experimentally include: • Vapor pressure depression • Boiling point elevation • Freezing point depression • Osmotic pressure Noncolligative properties, on the other hand, depend on the identity of the dissolved species and the solvent. Examples include solubility, surface tension, and viscosity. The addition of a solute to a solvent typically causes the vapor pressure of the solvent (above the resulting solution) to be lower than the vapor pressure above the pure solvent. As the concentration of the solute in the solution changes, so does the vapor pressure of the solvent above a solution. The vapor pressure of a solution of a nonvolatile solute is always lower than that of the pure solvent. For example, an aqueous solution of NaCl has a lower vapor pressure than pure water at the same temperature. The addition of solute to a pure solvent depresses the vapor pressure of the solvent. This observation, first made by Raoult, is now commonly known as Raoult’s law. The law states that the lowering of vapor pressure of a solution containing non-volatile solute is proportional to the mole fraction of the solute.
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Bokstein, Boris S., Mikhail I. Mendelev, and David J. Srolovitz. "Introduction to statistical thermodynamics of gases." In Thermodynamics and Kinetics in Materials Science. Oxford University Press, 2005. http://dx.doi.org/10.1093/oso/9780198528036.003.0014.

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As we discussed earlier in this book, thermodynamics provides very general relations between the properties of a system. On the other hand, thermodynamics is unable to predict any of the individual properties without the addition of either empirical or microscopic information. For example, we used thermodynamics to obtain Raoult’s law from Henry’s law, but we cannot derive Henry’s law from thermodynamic principles. Statistical mechanics provides an approach to determine individual thermodynamic properties from microscopic considerations. When applied in the realm of physical chemistry, we refer to this approach as statistical thermodynamics. In this chapter, we provide a simplified derivation of the Gibbs distribution, which is the basis of much of statistical thermodynamics. We then use statistical mechanics to show how the properties of an ideal gas can be obtained from a small number of properties of the molecules in the gas. This will allow us to determine such quantities as the equilibrium and rate constants of gas phase chemical reactions. As a result, we will gain new insight into the phenomena which we have already considered on the basis of phenomenological thermodynamics or formal kinetics. This approach will also show how to determine some of the parameters we previously introduced as input data in our thermodynamic considerations. As we have already seen, a finite system will eventually come into equilibrium with its surroundings. We even showed that when thermodynamic equilibrium is established, the temperatures, pressures, and chemical potentials of the system and its surroundings are equal (see Section 1.5.2). However, we never discussed what equilibrium actually is. For example, does this mean that the energy of the system is truly constant or is it only constant on average? When the system has a particular energy, does this mean that it is in a unique physical state or can it be in any one of several states that have exactly the same energy? In the latter case, can we simply talk about the probability the system is in each of these states? If the energy can fluctuate, what is the probability that the system has a particular energy? A very general approach to these types of questions was suggested by Gibbs and is now known as Gibbs statistics.
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Guedes, A., G. Platt, and F. Neto. "Inversion of functions from the plane to the plane to solve nonlinear algebraic systems: Calculating of double azeotrope using the modified Raoult’s Law in the mixture benzene + hexafluorobenzene." In Engineering Optimization 2014, 365–69. CRC Press, 2014. http://dx.doi.org/10.1201/b17488-65.

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Conference papers on the topic "Raoult's law"

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Daniel Gomes Ribeiro, João Flávio Vieira de Vasconcellos, and Gustavo Mendes Platt. "Comparison between different Artificial Neural Network architectures in VLE prediction models using experimental and simulated data (via Modified Raoult's Law) of binary systems." In 23rd ABCM International Congress of Mechanical Engineering. Rio de Janeiro, Brazil: ABCM Brazilian Society of Mechanical Sciences and Engineering, 2015. http://dx.doi.org/10.20906/cps/cob-2015-1311.

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Stirniman, Michael, and Jing Gui. "Polydispersity Effects in Evaporation of Perfluoropolyether Thin Films." In STLE/ASME 2001 International Joint Tribology Conference. American Society of Mechanical Engineers, 2001. http://dx.doi.org/10.1115/trib-nano2001-107.

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Abstract The evaporation rates of bulk liquid and thin films of an alcohol-derivatized perfluoropolyether have been studied experimentally and computationally. We find that the time dependence of the evaporation rate in both cases is dominated by the polydispersity, and can be described very well by a model that incorporates the molecular weight distribution, molecular-weight-dependent Arrhenius parameters of evaporation, and Raoult’s law of vapor pressures. Minor corrections to the model that account for surface interactions are necessary in the case of thin film evaporation.
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Howlett, Larry D. "The Theory of Osmosis." In ASME 2003 International Mechanical Engineering Congress and Exposition. ASMEDC, 2003. http://dx.doi.org/10.1115/imece2003-55040.

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A fresh view for explaining the process of osmosis and the phenomenon of osmotic pressure is presented. The process of osmosis was identified and modeled more than 100 years ago. Others have unsuccessfully challenged the original model developed by J.H. van’t Hoff. We revisit the basic equations and assumptions used in the thermodynamic derivation of the osmosis model. And, we propose a small but significantly different view of the traditional theory of osmosis. From this new view of osmosis and the osmosis experiment, we conclude that osmosis occurs at atmospheric pressure. In cellular membranes, flow from the solvent to the solution is related to the vapor pressure difference determined from the concentration difference with Raoult’s law. Furthermore, we suggest that osmotic pressure as determined from the osmosis experiment is related to both the solution properties and the membrane characteristics. We suggest that the difference between experimental and theoretical determination of osmotic pressure can be attributed to capillary action that may occur in some man made membranes.
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Gerken, William J., and Matthew A. Oehlschlaeger. "Nanofluid Pendant Droplet Evaporation." In ASME 2013 Heat Transfer Summer Conference collocated with the ASME 2013 7th International Conference on Energy Sustainability and the ASME 2013 11th International Conference on Fuel Cell Science, Engineering and Technology. American Society of Mechanical Engineers, 2013. http://dx.doi.org/10.1115/ht2013-17537.

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Nanofluids, stable colloidal suspensions of nanoparticles in a base fluid, have potential applications in the heat transfer, combustion and propulsion, manufacturing, and medical fields. Experiments were conducted to determine the evaporation rate of room temperature, millimeter-sized pendant droplets of ethanol laden with varying (0–3%) weight percentages of 40–60 nm aluminum nanoparticles (nAl). High-resolution droplet images were collected as a function of time for the determination of D-square law evaporation rates. Results show an asymptotic decrease in droplet evaporation rate with increasing nAl loading. The evaporation rate decreases by approximately 15% at around 1% to 3% nAl loading relative to the evaporation rate of pure ethanol, a reduction greater than can be explained by reduction in the vapor pressure of an ideal nanofluid mixture by Raoult’s law. It is hypothesized that the reduction in evaporation rate could be due to two phenomena: 1) the reduction in the ethanol volume fraction available for evaporation due to an interfacial layer on the immersed nanoparticle surface and 2) the aggregation of nanoparticles within the droplet and at the droplet surface, reducing the liquid diffusion rate to the surface and the liquid volume fraction at the surface available for evaporation.
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