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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 Rate equation ： ウィキペディア英語版 Rate equation The rate law or rate equation for a chemical reaction is an equation that links the reaction rate with concentrations or pressures of reactants and constant parameters (normally rate coefficients and partial reaction orders).〔(IUPAC Gold Book definition of rate law ). See also: According to IUPAC Compendium of Chemical Terminology.〕 For many reactions the rate is given by a power law such as :$r\backslash ;\; =\backslash ;\; k()^x()^y$ where () and () express the concentration of the species A and B, respectively (usually in moles per liter (molarity, M)). The exponents ''x'' and ''y'' are the partial reaction orders and must be determined experimentally; they are often not equal to the stoichiometric coefficients. The constant ''k'' is the ''rate coefficient'' or ''rate constant'' of the reaction. The value of this coefficient ''k'' may depend on conditions such as temperature, ionic strength, surface area of an adsorbent or light irradiation. For elementary reactions, which consist of a single step, the order equals the molecularity as predicted by collision theory. For example, a bimolecular elementary reaction A + B → products will be second order overall and first order in each reactant, with rate equation $r\backslash ;\; =\backslash ;\; k()\; ()$. For multistep reactions, the order of each step equals the molecularity, but this is ''not'' generally true for the overall rate. The rate equation of a reaction with an assumed multi-step mechanism can often be derived theoretically using quasi-steady state assumptions from the underlying elementary reactions, and compared with the experimental rate equation as a test of the assumed mechanism. The equation may involve a fractional order, and may depend on the concentration of an intermediate species. The rate equation is a differential equation, and it can be integrated to obtain an integrated rate equation that links concentrations of reactants or products with time. ==Zero order reactions== A Zero order reaction has a rate that is independent of the concentration of the reactant(s). Increasing the concentration of the reacting species will not speed up the rate of the reaction i.e. the amount of substance reacted is proportional to the time. Zero order reactions are typically found when a material that is required for the reaction to proceed, such as a surface or a catalyst, is saturated by the reactants. The rate law for a zero order reaction is :$\backslash \; r\; =\; k$ where r is the reaction rate and k is the reaction rate coefficient with units of concentration or time. If, and only if, this zeroth order reaction 1) occurs in a closed system, 2) there is no net build-up of intermediates, and 3) there are no other reactions occurring, it can be shown by solving a mass balance equation for the system that: :$r\; =\; -\backslash frac=k$ If this differential equation is integrated it gives an equation often called the integrated zero order rate law. :$\backslash \; ()\_t\; =\; -kt\; +\; ()\_0$ where $\backslash \; ()\_t$ represents the concentration of the chemical of interest at a particular time, and $\backslash \; ()\_0$ represents the initial concentration. A reaction is zero order if concentration data are plotted versus time and the result is a straight line. A plot of $\backslash \; ()\_t$ vs. time t gives a straight line with a slope of $-k$. The half-life of a reaction describes the time needed for half of the reactant to be depleted (same as the half-life involved in nuclear decay, which is a first order reaction). For a zero order reaction the half-life is given by : $\backslash \; t\_\; \backslash frac\; =\; \backslash frac$ ;Example of a zero order reaction * Reversed Haber process: $2NH\_3\; (g)\; \backslash rightarrow\; \backslash ;\; 3H\_2\; (g)\; +\; N\_2\; (g)$ The order of a reaction cannot be deduced from the chemical equation of the reaction. 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「Rate 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.