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翻訳と辞書　辞書検索 [ 開発暫定版 ] スポンサード リンク Arrhenius equation ： ウィキペディア英語版 Arrhenius equation The Arrhenius equation is a formula for the temperature dependence of reaction rates. The equation was proposed by Svante Arrhenius in 1889, based on the work of Dutch chemist Jacobus Henricus van 't Hoff who had noted in 1884 that Van 't Hoff's equation for the temperature dependence of equilibrium constants suggests such a formula for the rates of both forward and reverse reactions. Arrhenius provided a physical justification and interpretation for the formula.〔Laidler, K. J. (1987) ''Chemical Kinetics'',Third Edition, Harper & Row, p.42〕 Currently, it is best seen as an empirical relationship.〔Kenneth Connors, Chemical Kinetics, 1990, VCH Publishers〕 It can be used to model the temperature variation of diffusion coefficients, population of crystal vacancies, creep rates, and many other thermally-induced processes/reactions. The Eyring equation, developed in 1935, also expresses the relationship between rate and energy. A historically useful generalization supported by Arrhenius' equation is that, for many common chemical reactions at room temperature, the reaction rate doubles for every 10 degree Celsius increase in temperature.〔Pauling, L.C. (1988) ''General Chemistry'', Dover Publications〕 ==Equation== Arrhenius' equation gives the dependence of the rate constant $k$ of a chemical reaction on the absolute temperature $T$ (in kelvins), where $A$ is the pre-exponential factor (or simply the ''prefactor''), $E\_a$ is the activation energy, and $R$ is the universal gas constant:〔〔〔 :$k\; =\; A\; e^$ Alternatively, the equation may be expressed as :$k\; =\; A\; e^$ The only difference is the energy units of $E\_a$: the former form uses energy per mole, which is common in chemistry, while the latter form uses energy per molecule directly, which is common in physics. The different units are accounted for in using either the gas constant $R$ or the Boltzmann constant $k\_B$ as the multiplier of temperature $T$. The units of the pre-exponential factor $A$ are identical to those of the rate constant and will vary depending on the order of the reaction. If the reaction is first order it has the units s−1, and for that reason it is often called the ''frequency factor'' or ''attempt frequency'' of the reaction. Most simply, $k$ is the number of collisions that result in a reaction per second, $A$ is the number of collisions (leading to a reaction or not) per second occurring with the proper orientation to react and $e^\backslash \; \backslash \; \backslash$ is the probability that any given collision will result in a reaction. It can be seen that either increasing the temperature or decreasing the activation energy (for example through the use of catalysts) will result in an increase in rate of reaction. Given the small temperature range kinetic studies occur in, it is reasonable to approximate the activation energy as being independent of the temperature. Similarly, under a wide range of practical conditions, the weak temperature dependence of the pre-exponential factor is negligible compared to the temperature dependence of the $\backslash exp(-E\_a/RT)\backslash \; \backslash$ factor; except in the case of "barrierless" diffusion-limited reactions, in which case the pre-exponential factor is dominant and is directly observable. 抄文引用元・出典: フリー百科事典『 ウィキペディア（Wikipedia）』 ■ウィキペディアで「Arrhenius equation」の詳細全文を読む スポンサード リンク 翻訳と辞書 : 翻訳のためのインターネットリソース Copyright(C) kotoba.ne.jp 1997-2016. All Rights Reserved.