{"id":6035,"date":"2020-10-13T13:58:46","date_gmt":"2020-10-13T08:28:46","guid":{"rendered":"http:\/\/astan.lk\/al_virtualclassroom\/?p=6035"},"modified":"2020-10-13T14:00:03","modified_gmt":"2020-10-13T08:30:03","slug":"chemical-reactions-rate","status":"publish","type":"post","link":"https:\/\/astan.lk\/al_virtualclassroom\/chemical-reactions-rate\/","title":{"rendered":"Rate of reaction"},"content":{"rendered":"

Consider a chemical reaction<\/p>\n

aA + bB\u00a0\u2192 cC + dD<\/p>\n

rate with respect to the change of concentration of the reactant A = – \u0394[A]\/\u0394t<\/p>\n

rate with respect to the change of concentration of the product D = \u0394[D]\/\u0394t<\/p>\n

\"Average<\/p>\n

\u2022 In a given reaction, the rates of removal of each reactant and rates of formation of each products are not equal.<\/p>\n

\u2022 The rate of removal of a reactant or formation of a product depends on the stoichiometric coefficients of the respective substances.<\/p>\n

rate of reaction = – (1\/a)\u0394[A]\/\u0394t \u00a0 \u00a0 \u00a0= \u00a0-(1\/b)\u0394[B]\/\u0394t \u00a0 \u00a0 = \u00a0(1\/c)\u0394[C]\/\u0394t \u00a0 = (1\/d)\u0394[D]\/\u0394t<\/p>\n

 <\/p>\n

\u2022 Rate law for the above generalized equation = k [A] x [B] y, where x and y are
\norder of reaction with respect to reactants A and B respectively.
\n\u2022 (x + y) of the above expression is referred as the overall order of the reaction.<\/p>\n

Example\u00a0<\/span><\/p>\n

Express the rate of the reaction given, in terms of reactants and product concentrations.<\/span><\/p>\n

<\/span><\/p>\n

The rate of the reaction is expressed as<\/span><\/p>\n

<\/span><\/p>\n

<\/h4>\n

Zero order reaction<\/span><\/h4>\n

\u00a0<\/span>Ammonia (NH<\/span>3<\/span><\/sub>) gas decomposes over platinum catalyst to nitrogen gas (N<\/span>2<\/span><\/sub>) and hydrogen gas (H<\/span>2<\/span><\/sub>). The chemical reaction is as follows:<\/span><\/p>\n

\"example<\/span><\/h5>\n

The reaction follows zero order kinetics. Therefore the rate law is rate r = k[NH<\/span>3<\/span><\/sub>]<\/span>o<\/span><\/sup><\/p>\n

For a zero order reaction the concentration versus time profile is linear and the rate of reaction versus time has the profile.<\/span><\/p>\n

Concentration versus time profile for a zero order reaction<\/strong><\/h5>\n
\"Concentration<\/span><\/h5>\n

 <\/p>\n

\u00a0Rate of reaction versus time for a zero order reaction<\/strong><\/p>\n

\"Rate<\/span><\/h5>\n

<\/h4>\n

First order reaction<\/h4>\n

Cyclopropane (C3<\/span><\/sub>H<\/span>6<\/span><\/sub>) at room temperature has a ring structure. When it is heated, the ring opens up and cyclopropane isomerizes to propylene.<\/span><\/p>\n

\"example<\/span><\/h5>\n

The reaction takes place in a first order manner. Therefore, the rate law for the reaction is, rate (r) = k [C<\/span>3<\/span><\/sub>H<\/span>6<\/span><\/sub>]<\/span>1<\/span><\/sup><\/p>\n

First Order Reaction: Concentration versus time profile<\/strong><\/h5>\n
\"example<\/span><\/h5>\n

The above figure is an example of the first order reaction. The rate of the reaction depends on the concentration of the reactant. Initially the rate is fast and then it slows down as the concentration of the reactant falls.<\/p>\n

 <\/p>\n

Second order reaction<\/h4>\n

Nitrogen dioxide (NO2<\/span><\/sub>) gas reacts with fluorine gas (F<\/span>2<\/span><\/sub>) to give nitrosyl fluoride. The chemical reaction is given by the equation.<\/span><\/p>\n

<\/span><\/p>\n

The reaction has the rate law as:<\/span><\/p>\n

rate (r) = k [NO<\/span>2<\/span><\/sub>] [F<\/span>2<\/span><\/sub>]<\/span><\/p>\n

The rate has order one with respect to nitrogen dioxide concentration and fluorine concentration and the overall order is (1 + 1) which is two.<\/span><\/p>\n

This depicts the change of concentration of second order reaction with only one reactant, that is, a reaction of the type A P where the rate law is rate r = k [A]<\/span>2<\/span><\/sup>.<\/span><\/p>\n

 <\/p>\n

Experiment I : Determination of the effect of concentration of the acid on the reaction between magnesium and acids.<\/h4>\n

Requirements : 250 cm3<\/sup> of 0.1 mol dm-3<\/sup> H2<\/sub>SO4<\/sub> , Six pieces of 2.0 cm cleaned magnesium ribbon.<\/p>\n

 <\/p>\n

Method<\/strong><\/p>\n

\"\"<\/a><\/p>\n

Mg(s) + H2<\/sub>SO4<\/sub>(aq)\u00a0\u2192 MgSO4<\/sub>(aq) + H2<\/sub>(g)<\/p>\n

\u2022 Into a boiling tube, 10.0 cm3 of water is added and marked with a rubber band.<\/p>\n

\u2022 Into the same boiling tube, 40.0 cm3<\/sup> of 0.1 mol dm-3<\/sup> H2<\/sub>SO4<\/sub> is added and filled with water brimfully.<\/p>\n

\u2022 The cleaned magnesium ribbon is fixed as shown in the diagram to the bung, the tube is closed with the bung and turned upside down while switching on the stop watch simultaneously.<\/p>\n

\u2022 Time taken by the gas to fill up to the level of the rubber band is measured.<\/p>\n

\u2022 Measurement of time is repeated for different solutions tabulated below, every time
\nfixing a fresh Mg ribbon.<\/p>\n

\"\"<\/a><\/p>\n

\u2022Calculate the concentration of each acid solution and plot a graph between [H+<\/sup>] and 1\/t.<\/p>\n

R\u00a0\u221d [H+<\/sup>]n<\/sup><\/p>\n

Initial rate =\u00a0Average rate for a small change from the starting point<\/p>\n

= Volume of gas produced\/Time \u00a0 \u00a0= Constant\/Time \u00a0= k\/t<\/p>\n

R\u00a0\u221d 1\/t<\/p>\n

[H+<\/sup>]n<\/sup>\u00a0\u221d 1\/t<\/p>\n

<\/h4>\n

Experiment II : Determination of the effect of concentration on the rate of the reaction between Na2<\/sub>S2<\/sub>O3<\/sub> and HCl.<\/h4>\n

Requirements : 0.10 mol dm-3 Na2<\/sub>S2<\/sub>O3<\/sub> solution. 2.0 mol dm-3<\/sup> HCl solution, stop
\nwatch, boiling tubes, measuring cylinders, 50 cm3<\/sup> beaker.<\/p>\n

S2<\/sub>O3<\/sub>2-<\/sup>(aq) + 2H+<\/sup>(aq) \u00a0\u2192 \u00a0S(s) + SO2<\/sub>(g) + H2<\/sub>O(l)<\/p>\n

Method<\/strong>
\n\u2022 Given volume of Na2<\/sub>S2<\/sub>O3<\/sub> solution as shown in the table is added to a 50 cm3<\/sup> beaker.<\/p>\n

\u2022 The beaker containing the Na2<\/sub>S2<\/sub>O3<\/sub> solution is kept on a white paper marked with a cross. Relevant volume of the HCl solution is added and the time to disappear the cross is measured keeping the eye at a constant height from the beaker.<\/p>\n

\u2022 The beaker is cleaned and the experiment is repeated mixing the solutions given in the table below.<\/p>\n

(i) Determining the relationship between the reaction rate and the thiosulphate ion concentration.
\nThe experiment is done using thiosulphate solutions of different concentrations as given in the table.<\/p>\n

\"\"<\/a><\/p>\n

R \u221d [S2<\/sub>O3<\/sub>2-<\/sup>(aq)]n<\/sup><\/p>\n

 <\/p>\n

(ii) Determining the relationship between the reaction rate and the concentration of hydrogen ions.
\nThe experiment is done using acid solutions of different concentrations as given in the
\ntable.<\/p>\n

\"\"<\/a><\/p>\n

R\u00a0\u221d [H+<\/sup>]n<\/sup><\/p>\n

In both the above instances, assume that the rate is constant during the period of measurement of time taken to disappear cross and it is equal to the initial rate. Then, Rate = constant \/ t
\nIn both instances examine how concentration varies with 1\/t.<\/p>\n

<\/h4>\n

Experiment III : Determination of the effect of concentration on the rate of the reaction between Fe(III) ions and KI.<\/h4>\n

2Fe3+<\/sup>(aq) + 2I–<\/sup>(aq)\u00a0\u2192 I2<\/sub>(aq) + 3Fe2+<\/sup>(aq)<\/p>\n

The amount of I2<\/sub> produced can be used to determine the rate of this reaction. The
\nminimum iodine concentration required to turn starch blue is 1.0×10-5<\/sup> mol dm-3<\/sup>. Since this is very small, measuring time is difficult. Therefore, time taken to appear blue colour should be delayed to measure the time. This can be affected through a faster reaction which converts I2<\/sub> to I–<\/sup>. For this Na2<\/sub>S2<\/sub>O3<\/sub> can be used.<\/p>\n

2S2O3<\/sub>2-<\/sup>(aq) + I2<\/sub>(aq) \u2192\u00a0S4<\/sub>O2-<\/sup>(aq) + 2I–<\/sup>(aq)<\/p>\n

A known amount of Na2<\/sub>S2<\/sub>O3<\/sub> is added to the medium. The moment in which Na2<\/sub>S2<\/sub>O3<\/sub> is over, the solution turns blue. The amount of I2<\/sub> formed depends on the amount of Na2<\/sub>S2<\/sub>O3<\/sub> added
\nRequirements :
\n0.10 mol dm-3<\/sup> KI solution, 0.10 mol dm-3<\/sup> FeCl3<\/sub> or Fe(NH4<\/sub>)(SO4<\/sub>)2<\/sub> solution,
\n0.10 mol dm-3<\/sup> Na2<\/sub>S2<\/sub>O3<\/sub> solution, 0.10 mol dm-3<\/sup> H2SO4 solution, stopwatch\"\"<\/a><\/p>\n

Mix the solutions given in the above table as shown here. Measure the time taken to turn the solution blue.<\/p>\n

\"\"<\/a><\/p>\n

\"\"<\/a><\/p>\n

In both instances examine how rate (1\/t) varies with concentration.
\n\"\"<\/a><\/p>\n

According to the shape of the graph, order of the reaction (n) can be determined.<\/p>\n

 <\/p>\n

Initial rate, instantaneous rate and average rate<\/h3>\n

Determining the Average Rate from Change in Concentration over a Time Period<\/span><\/b><\/h4>\n

We calculate the average rate of a reaction over a time interval by dividing the change in concentration over that time period by the time interval. For the change in concentration of a reactant, the equation, where the brackets mean “concentration of”, is<\/span><\/p>\n

\"Average<\/span><\/center>Note: We use the minus sign before the ratio in the previous equation because a rate is a positive number. We do not need the minus sign when calculating average rates from products.<\/span><\/p>\n

<\/center>
\u00a0<\/span><\/center><\/p>\n

Determining the Instantaneous Rate from a Plot of Concentration Versus Time<\/span><\/b><\/h4>\n

An instantaneous rate is the rate at some instant in time. An instantaneous rate is a differential rate: -d[reactant]\/dt or d[product]\/dt.<\/span><\/p>\n

We determine an instantaneous rate at time t:<\/span><\/p>\n

    \n
  • by calculating the <\/span>negative<\/span> of the slope of the curve of concentration of a reactant versus time at time t.<\/span><\/li>\n
  • by calculating the slope of the curve of concentration of a product versus time at time t.<\/span><\/li>\n<\/ul>\n

    \"Graph<\/span><\/center>
    \u00a0<\/span><\/center><\/p>\n

    Determining the Initial Rate from a Plot of Concentration Versus Time<\/span><\/b><\/h4>\n

    The initial rate of a reaction is the instantaneous rate at the start of the reaction (i.e., when t = 0). The initial rate is equal to the negative of the slope of the curve of reactant concentration versus time at t = 0.<\/span><\/p>\n

     <\/p>\n

    Half life of a reaction<\/h4>\n

    Half life of a reaction is defined as the time required for reducing the concentration of a reactant to half its initial value. It is denoted as t<\/span>1<\/span><\/sub>\/<\/span><\/sub>2<\/span><\/sub>.<\/span><\/p>\n

    The t<\/span>1<\/span><\/sub>\/<\/span><\/sub>2<\/span><\/sub> of a zero order reaction is given as<\/span><\/p>\n

    \"half<\/span><\/h5>\n

    while for the first order reaction it is given as<\/span><\/p>\n

    \"half<\/span><\/h5>\n

     <\/p>\n

    Only for the first order reaction t<\/span>1<\/span><\/sub>\/<\/span><\/sub>2<\/span><\/sub> is independent of the initial concentration of the reactant. This relation can be used to determine the order of a reaction.<\/span><\/p>\n

    <\/p>\n

     <\/p>\n","protected":false},"excerpt":{"rendered":"

    Consider a chemical reaction aA + bB\u00a0\u2192 cC + dD rate with respect to the change of concentration of the reactant A = – \u0394[A]\/\u0394t rate with respect to the change of concentration of the product D = \u0394[D]\/\u0394t \u2022 In a given reaction, the rates of removal of each reactant and rates of formation […]<\/p>\n","protected":false},"author":842,"featured_media":0,"comment_status":"open","ping_status":"open","sticky":false,"template":"","format":"standard","meta":{"_acf_changed":false,"_monsterinsights_skip_tracking":false,"_monsterinsights_sitenote_active":false,"_monsterinsights_sitenote_note":"","_monsterinsights_sitenote_category":0,"footnotes":""},"categories":[14,1674],"tags":[],"class_list":["post-6035","post","type-post","status-publish","format-standard","hentry","category-chemistry","category-unit-12"],"acf":[],"yoast_head":"\nRate of reaction - Learning & Education Portal<\/title>\n<meta name=\"robots\" content=\"index, follow, max-snippet:-1, max-image-preview:large, max-video-preview:-1\" \/>\n<link rel=\"canonical\" href=\"https:\/\/astan.lk\/al_virtualclassroom\/chemical-reactions-rate\/\" \/>\n<meta property=\"og:locale\" content=\"en_US\" \/>\n<meta property=\"og:type\" content=\"article\" \/>\n<meta property=\"og:title\" content=\"Rate of reaction - 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