REACTION RATES RATES Cont Conten ents ts for for this this pag page e
Rela Re late ted d top topic ics s
Reactants and products Definition of the reaction rate Initial rates Catalysis Factors affecting Chemical reaction rates equilibrium The collision theory Measurement of reaction rates Additional questions
Data Glossary
Learning Outcomes After studying this section, you will (a)understand what is meant by the "rate of a chemical reaction", (b) and you will be familiar with the factors which influence it, and (c) have an understanding of how reaction rates are measured.
Reactants and products: A chemical reaction involves one or more reactants and products:
In the example above, zinc in a solid form, as indicated by the symbol (s) , reacts with sulphuric acid in an aqueous solution, as indicated by the symbol (aq). These are the REACTANTS . The reactants react to form PRODUCTS written to the right of the arrow. In the above example, the products are an aqueous solution of the salt zinc sulphate and hydrogen in gaseous
form, which is indicated by the symbol (g). If a reactant is present in amounts that are less than will completely use up the other, it is called the LIMITING REACTANT .
Definition of the reaction rate: The rate of a chemical reaction is defined as the change in the concentration of one of the reactants or products in unit time
For example, in the reaction above, we could express the rate of the reaction as the change in the concentration of H 2SO4 (Δ[H2SO4]) in a certain time interval :
Let's consider a general reaction: A → B: If one plots the concentration of A, [A] against the time, t, one typically obtains a curve such as the one shown on the right: The fact that the graph is not a straight line tells us the rate of the reaction changes with time.
form, which is indicated by the symbol (g). If a reactant is present in amounts that are less than will completely use up the other, it is called the LIMITING REACTANT .
Definition of the reaction rate: The rate of a chemical reaction is defined as the change in the concentration of one of the reactants or products in unit time
For example, in the reaction above, we could express the rate of the reaction as the change in the concentration of H 2SO4 (Δ[H2SO4]) in a certain time interval :
Let's consider a general reaction: A → B: If one plots the concentration of A, [A] against the time, t, one typically obtains a curve such as the one shown on the right: The fact that the graph is not a straight line tells us the rate of the reaction changes with time.
Now look at the graph on the left. In the time interval Δt covering t = 0 to t = t 1, there is a comparatively greater change in [A] than in the same time interval Δt = t1 to t = t2. The rate of a reaction is therefore always linked to the specific time interval in which the concentration measurements were made!
Initial rates: THE INITIAL RATE of a reaction is the rate right at the start of the reaction, when t = 0. It is calculated by determining the gradient of the tangent to the curve at time t = 0:
Looking at the curve again, we see that its steepest part occurs right at the start of the reaction. The gradient at t = 0 is the steepest, and that will be when the rate will have its highest value.
The collision theory: The COLLISION THEORY explain the factors that influence reaction rates. Basically, the theory postulates that the rates of reactions depend on how often and how energetically the reacting molecules collide with each other, the faster will the overall reaction proceed. We will see below how this theory is applied.
Factors affecting reaction rates: In order for two reactants to react, they must come into close contact, in other words, they must COLLIDE. However, it is not sufficient for two molecules to collide in order that they might react. The reaction will only take place if the collisions are "fruitful", and for this to happen, • •
the molecules have sufficient energy, and, they collide with the proper orientation (this is of particular importance when organic molecules react).
There are a number of factors that increase the likelyhood of fruitful collisions will generally tend to increase the rates of reactions. These are discussed below.
Concentration:
For reactions that take place in solution, the greater the concentration of reactants, the greater the rate of the reaction, as the likelyhood of collisions increases the more there are reactants in a given volume.
Pressure: For reactions involving gases, an increase in pressure increases the reaction rates, since, in a container having a fixed volume, the concentration of gases are directly proportional to the pressure inside the container.
Temperature: The higher the temperature, the higher the reaction rate. This is because the kinetic energy of the reactant molecules increases with temperature, and thus, the collisions are more energetic. The graphs shown here on the left are "MaxwellBoltzmann" DISTRIBUTION CURVES . Energy values are plotted on the horizontal axis, while the number of molecules with corresponding energies are plotted on the vertical axis. Two such curves are shown, one (blue) at a temperature T 1, and one, in red, at a higher temperature T2. Ea, the so-called ENERGY OF ACTIVATION , is the minimum value that a molecule must achieve in order for
the reaction to occur. At the lower temperature, T 1, only a relatively small proportion of molecules (area below the blue curve to the right of Ea will have sufficient energy to react. At the higher temperature, T 2, a much higher proportion of the molecules have the required energy. Thus, since the rate of the reaction depends on the number of molecules that react in a given time interval, the rate will be higher at T2 than at T 1.
Bond types: Reactions between ionic compounds are usually very much faster than those involving compounds where covalent bonds have to be made or broken.
State of subdivision: This is very important when solids are involved. The more finely divided the solid is, the faster the reaction will take place. In fact, for solids, it is the surface area, and not the concentration, that affects the rate. Of course, for a reaction involving a solid reactant and another that is in solution, the rate will be affected not only by the surface of the solid reactant, but also by the concentration of the one in solution. In terms of the collision theory, the grater the surface are of a solid, the more collisions will take place at the surface, and hence the greater the reaction rate.
Catalysts: These are substances which speed up reaction rates without undergoing a permanent change in the process. The process is called CATALYSIS, and it may be HOMOGENEOUS , when both the reactants and the catalyst are in the same phase (for example, when both are in solution), or HETEROGENEOUS , when the ractants and catalyst are in different phases (for example, gaseous reactants and a solid catalyst).
Measurement of reaction rates: We have seen that the rate of a chemical reaction is defined as the change in the concentration of one of the reactants or products in unit time. For a reaction A → B, since one molecule of B is produced for every molecule of A that is consumed, it does not matter whether we measure a decrease in the concentration of A or an increase in the concentration of B the absolute value of the rate will be the same. We can then write:
From an experimental point of view, if one wished to measure the rate of such a reaction, it would not matter whether one measured the decrease in concentration of A, or the increase in concentration of B. We would get the same results. In practice, it may well be that one of the two alternatives is much easier to perform in the laboratory, and that would then be the method of choice.
The rate law: Experimentally, the rate of a reaction is found to depend (amongst other things!) on the concentration of the reactants. For a generalised reaction
we find that
where k is the RATE CONSTANT for the reaction.
Experimental techniques: There are many ways of determining reaction rates. The following will give the basic principles involved in some of them: Changes in volumes:
This is useful for reactions that produce gases. The reaction flask is connected to a syringe, as shown in the diagram on the right, and the volumes indicated on the syringe scale read off at regular time intervals.
Changes in mass:
Again, if a gas is evolved during a reaction, for example, the reaction between calcium carbonate and dilute acid, the rate at which the apparatus loses mass will be proportional to the reaction rate. Calcium carbonate powder may be suspended in water in an Erlenmeyer flask and placed on the pan of an
electronic top-loading balance. The reaction is started by the addition of acid, and the mass recorded at various time intervals starting immediately after addition of the acid.
Changes in colour: If a reaction produces a coloured substance, that is, one that has an absorption maximum at some wavelength in the visible spectrum, one can measure the increase in the absorbance of the solution in which the reaction takes place as a function of time. Since the absorbance of a solute is directly proportional to its molar concentration, the rate of the reaction is easy to determine. One will of course need a fairly expensive piece of equipment called a SPECTROPHOTOMETER . A similar method, known as TURBIDIMETRY makes use of the scattering of light from aqueous suspensions. If a precipitate gradually forms, the solution will become more opaque with time as more finely divided particles precipitate out. The rate at which the opacity of the liquid increases can be used as a measure of the reaction rate. In fact, any measurable change that takes place during a reaction may be used to determine reaction rates. One must always bear in mind the various factors that influence the rates. (See above) <>
Additional questions
ORDERS OF REACTION AND RATE EQUATIONS Changing the concentration of substances taking part in a reaction usually changes the rate of the reaction. A rate equation shows this effect mathematically. Orders of reaction are a part of the rate equation. This page introduces and explains the various terms you will need to know about. Note: If you aren't sure about why changing concentration affects rates of reaction you might like to follow this link and come back to this page afterwards - either via the rates of reaction menu or by using the BACK
button on your browser.
Rate equations Measuring a rate of reaction There are several simple ways of measuring a reaction rate. For example, if a gas was being given off during a reaction, you could take some measurements and work out the volume being given off per second at any particular time during the reaction. A rate of 2 cm 3 s-1 is obviously twice as fast as one of 1 cm 3 s-1. Note: Read cm3 s-1 as "cubic centimetres per second".
However, for this more formal and mathematical look at rates of reaction, the rate is usually measured by looking at how fast the concentration of one of the reactants is falling at any one time. For example, suppose you had a reaction between two substances A and B. Assume that at least one of them is in a form where it is sensible to measure its concentration - for example, in solution or as a gas.
For this reaction you could measure the rate of the reaction by finding out how fast the concentration of, say, A was falling per second. You might, for example, find that at the beginning of the reaction, its concentration was falling at a rate of 0.0040 mol dm -3 s-1. Note: Read mol dm-3 s-1 as "moles per cubic decimetre (or litre) per second".
This means that every second the concentration of A was falling by 0.0040 moles per cubic decimetre. This rate will decrease during the reaction as A gets used up. Summary For the purposes of rate equations and orders of reaction, the rate of a reaction is measured in terms of how fast the concentration of
one of the reactants is falling. Its units are mol dm -3 s-1.
Orders of reaction I'm not going to define what order of reaction means straight away I'm going to sneak up on it! Orders of reaction are always found by doing experiments. You can't deduce anything about the order of a reaction just by looking at the equation for the reaction. So let's suppose that you have done some experiments to find out what happens to the rate of a reaction as the concentration of one of the reactants, A, changes. Some of the simple things that you might find are: One possibility: The rate of reaction is proportional to the concentration of A That means that if you double the concentration of A, the rate doubles as well. If you increase the concentration of A by a factor of 4, the rate goes up 4 times as well. You can express this using symbols as:
Writing a formula in square brackets is a standard way of showing a concentration measured in moles per cubic decimetre (litre). You can also write this by getting rid of the proportionality sign and introducing a constant, k.
Another possibility: The rate of reaction is proportional to the square of the concentration of A This means that if you doubled the concentration of A, the rate would go up 4 times (2 2). If you tripled the concentration of A, the rate would increase 9 times (3 2). In symbol terms:
Generalising this By doing experiments involving a reaction between A and B, you would find that the rate of the reaction was related to the concentrations of A and B in this way:
This is called the rate equation for the reaction. The concentrations of A and B have to be raised to some power to show how they affect the rate of the reaction. These powers are called the orders of reaction with respect to A and B. For UK A' level purposes, the orders of reaction you are likely to
meet will be 0, 1 or 2. But other values are possible including fractional ones like 1.53, for example. If the order of reaction with respect to A is 0 (zero), this means that the concentration of A doesn't affect the rate of reaction. Mathematically, any number raised to the power of zero (x 0) is equal to 1. That means that that particular term disappears from the rate equation. The overall order of the reaction is found by adding up the individual orders. For example, if the reaction is first order with respect to both A and B (a = 1 and b = 1), the overall order is 2. We call this an overall second order reaction. Some examples Each of these examples involves a reaction between A and B, and each rate equation comes from doing some experiments to find out how the concentrations of A and B affect the rate of reaction. Example 1:
In this case, the order of reaction with respect to both A and B is 1. The overall order of reaction is 2 - found by adding up the individual orders. Note: Where the order is 1 with respect to one of the reactants, the "1" isn't written into the equation. [A] means [A] 1.
Example 2:
This reaction is zero order with respect to A because the concentration of A doesn't affect the rate of the reaction. The order with respect to B is 2 - it's a second order reaction with respect to B. The reaction is also second order overall (because 0 + 2 = 2). Example 3:
This reaction is first order with respect to A and zero order with respect to B, because the concentration of B doesn't affect the rate of the reaction. The reaction is first order overall (because 1 + 0 = 1). What if you have some other number of reactants? It doesn't matter how many reactants there are. The concentration of each reactant will occur in the rate equation, raised to some power. Those powers are the individual orders of reaction. The overall order of the reaction is found by adding them all up.
The rate constant Surprisingly, the rate constant isn't actually a true constant! It varies, for example, if you change the temperature of the reaction, add a catalyst, or change the catalyst. The rate constant is constant for a given reaction only if all you are changing is the concentration of the reactants. You will find more about the effect of temperature and catalysts on the rate constant on another page. Note: If you want to follow up this further look at rate constants you might like to follow this link. Alternatively, you could visit it later via the rates of reaction menu.
Calculations involving orders of reaction You will almost certainly have to be able to calculate orders of reaction and rate constants from given data or from your own experiments. There are all sorts of ways of doing these sums, and it is important that you practice the methods that your syllabus wants. Check your syllabus and past exam papers to see what sort of examples you need to be able to work out. Note: For UK A'level students, if you haven't got copies of your syllabus and past papers follow this link to find out how to get hold of them.
Many text books make these sums look really difficult. In fact for A' level purposes, the calculations are usually fairly trivial. You will find them explained in detail in my chemistry calculations book. Note: There are several reasons why there are very few calculations on this site. It is much easier to learn to do sums from a carefully organised book than from a website; I would be in breach of my contract with my publishers if I included material similar to what is in the book; and I need to sell a few books to generate some income! If you are interested in my chemistry calculations book you might like to follow this link.
Where would you like to go now? To a simple look at how orders of reaction are related to reaction mechanisms . . To the rates of reaction menu . . . To the Physical Chemistry menu . . . To Main Menu . . .
© Jim Clark 2002
http://www.chemguide.co.uk/physical/basicrates/orders.html#top
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The rate of a reaction can be measured by the rate at which a reactant is used up, or the rate at which a product is formed. The temperature, concentration, pressure of reacting gases, surface area of reacting solids, and the use of catalysts, are all factors which affect the rate of a reaction. Chemical reactions can only happen if reactant particles collide with enough energy. The more frequently particles collide, and the greater the proportion of collisions with enough energy, the greater the rate of reaction.
Measuring rates Different reactions can happen at different rates. Reactions that happen slowly have a low rate of reaction. Reactions that happen quickly have a high rate of reaction. For example, the chemical weathering of rocks is a very slow reaction: it has a low rate of reaction. Explosions are very fast reactions: they have a high rate of reaction.
Reactants and products There are two ways to measure the rate of a reaction:
1. measure the rate at which a reactant is used up 2. measure the rate at which a product is formed The method chosen depends on the reaction being studied. Sometimes it is easier to measure the change in the amount of a reactant that has been used up; sometimes it is easier to measure the change in the amount of product that has been produced.
Things to measure The measurement itself depends on the nature of the reactant or product: •
the mass of a substance - solid, liquid or gas - is measured with a balance
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the volume of a gas is usually measured with a gas syringe, or sometimes an upside down measuring cylinder or burette
It is usual to record the mass or total volume at regular intervals and plot a graph. The readings go on the vertical axis, and the time goes on the horizontal axis.
For example, if 24cm3 of hydrogen gas is produced in two minutes, the mean rate of reaction = 24 ÷ 2 = 12cm 3 hydrogen / min.
http://www.bbc.co.uk/schools/gcsebitesize/science/add_aqa/chemreac/ratesrev1.shtml
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Reaction Rates
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Foundation We will assume an understanding of the postulates of the Kinetic Molecular Theory and of the energetics of chemical reactions. We will also assume an understanding of phase equilibrium and reaction equilibrium, including the temperature dependence of equilibrium constants.
Goals We have carefully examined the observation that chemical reactions come to equilibrium. Depending on the reaction, the equilibrium conditions can be such that there is a mixture of reactants and products, or virtually all products, or virtually all reactants. We have not considered the time scale for the reaction to achieve these conditions, however. In many cases, the speed of the reaction might be of more interest than the final equilibrium conditions of the reaction. Some reactions proceed so slowly towards equilibrium as to appear not to occur at all. For example, metallic iron will eventually oxidize in the presence of aqueous salt solutions, but the time is sufficiently long for this process that we can reasonably expect to build a boat out of iron. On the other hand, some reactions may be so rapid as to pose a hazard. For example, hydrogen gas will react with oxygen gas so rapidly as to cause an explosion. In addition, the time scale for a reaction can depend very strongly on the amounts of reactants and their temperature. In this concept development study, we seek an understanding of the rates of chemical reactions. We will define and measure reaction rates and develop a quantitative analysis of the dependence of the reaction rates on the conditions of the reaction, including concentration of reactants and temperature. This quantitative analysis will provide us insight into the process of a chemical reaction and thus lead us to develop a model to provide an understanding of the significance of reactant concentration and temperature. We will find that many reactions proceed quite simply, with reactant molecules colliding and exchanging atoms. In other cases, we will find that the process of reaction can be quite complicated, involving many molecular collisions and rearrangements leading from reactant molecules to product molecules. The rate of the chemical reaction is determined by these steps.
Observation 1: Reaction Rates
Doc Brown's Chemistry KS4 science GCSE/IGCSE Revision Notes Factors affecting the Speed-Rates of Chemical Reactions This page describes the factors controlling the speeds of chemical reactions and the collision theory behind it discussed. The factors affecting the speed of reaction are also presented using particle models to give a theoretical basis to the rules on the effects of concentration, pressure, temperature, solid reactant particle size (surface area), stirring, catalysts and light. Methods of how to collect data are also described and graphical treatment of the observations and how to draw conclusions. How do we graphically interpret tables of results and what graphs are useful and why. How to we calculate the speed of a reaction? Revision KS4 Science GCSE/IGCSE/O level Chemistry Information Study Notes for revising for AQA GCSE Science, Edexcel 360Science/IGCSE Chemistry & OCR 21stC Science, OCR Gateway Science WJEC gcse science chemistry CCEA/CEA gcse science chemistry O Level Chemistry (revise courses equal to US grade 8, grade 9 grade 10) SECTIONS on this page: 1. What do mean by rate/speed of reaction and its measurement? * 2. Collision theory of reaction * 3. Factors: 3a concentration, 3b pressure, 3c stirring, 3d particle size/surface area, 3e temperature, 3f
catalyst,, 3g light * 4. Examples of graphs catalyst KEY WORDS-PHRASES in alphabetical order for this rates web page: hydrochloric/sulphuric hydrochloric/sulphu ric acid-metal e.g. Mg/carbon Mg/carbonate ate reaction * hydrochloric acid-sodium thiosulphate reaction * Activation energy * Catalysts * Concentrat Concentration ion effect * Graphs-gas collection * Graphs-examples * hydrogen peroxide decomposition * How reactions happen * Interpreting results * Light (catalyst) effect * Methods of measuring rate * Pressure effect * Rate of reaction * Reaction profiles * Stirring effect * Surface area/size of solid particle reactant effect * Temperature effect
See also
Advanced Level Chemistry Theory pages on "CHEMICAL KINETICS"
1. What do we mean by Rate and how is it measured? •
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WHAT DO WE MEAN BY SPEED OR RATE IN THE CONTEXT OF A CHEMICAL REACTION? IS IT TO FAST OR TO SLOW TO MEASURE THE SPEED? WHAT SORT OF WAYS CAN WE MEASURE THE SPEED OF A CHEMICAL REACTION? The phrase ‘rate ‘rate of reaction’ reaction’ means ‘how fast is the reaction’ or 'the speed of the reaction'. It can be measured as the 'rate of formation of product' product ' (e.g. collecting gaseous product in a syringe) or the 'rate ' rate of removal of reactant'. reactant'. The speeds of reactions are very varied. o Rusting is a ‘slow’ reaction, reaction, you hardly see any change looking at it! The weathering of rocks is an extremely very slow reaction. reaction . o o The fermentation of sugar to alcohol is quite slow but you can see the carbon dioxide bubbles forming in the 'froth' in a laboratory experiment or beer making in industry! o A faster reaction example is magnesium reacting with hydrochloric acid to form magnesium chloride and hydrogen or the even faster reaction between sodium and water to form sodium hydroxide. Combustion reactions e.g. when a fuel burns in air or oxygen, o is a very fast reaction. A 'use of words' revision note : Reacting and/or dissolving? Chemical or physical change? If you take the solids magnesium chloride or sodium hydroxide and mix them with water they dissolve to form a solution, but no chemical reaction to form new substances takes place i.e. dissolving on its own is basically a physical change. change . However, the two substances mentioned above are formed in a chemical reaction change, change,
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where the word 'dissolving' on its own is inadequate. The phrases reaction with ... or reaction between ... are much more appropriate, but there is no denying that the magnesium/sodium dissolve in acid/water, BUT only because they have formed a water soluble compound. compound . o Explosive reactions would be described as ‘very fast’ e.g. the pop of a hydrogen-air mixture on applying a lit splint or the production of a gas to inflate the air bags safety feature of many cars. The importance of "Rates of Reaction knowledge": knowledge" : o Time is money in industry, industry, the faster the reaction can be done, the more economic it is. You need to know how long reactions are likely to take. Hence the great importance of catalysts e.g. transition metals or enzymes which reduce time and save money. o Health and Safety Issues: Issues : Mixtures of flammable gases in air present an explosion hazard (gas reactions like this are amongst the fastest reactions known). e.g. Methane gas in mines, petrol vapour etc. are all potentially dangerous situations so knowledge of 'explosion/ignition threshold concentrations', ignition temperatures and activation energies are all important knowledge to help design systems of operation to minimise risks. Flammable fine dust powders can be easily ignited e.g. coal dust in mines, flour in mills. Fine powders have a large surface area which greatly increases the reaction rate causing an explosion. Any spark from friction is enough to initiate the reaction! A reaction will continue until one of the reactants is used up. up. To measure the ‘speed’ or ‘rate’ of a reaction depends on what the reaction is, and can what is formed be measured as the reaction proceeds? Two examples are outlined be low. When a gas is formed from a solid reacting with a solution, it can be collected in a gas syringe (see diagram below and the graph graph). ). The initial gradient of the graph e.g. in cm3 /min (speed or rate) o gives an accurate measure of how fast a gaseous product is being formed in metal/carbonate - acid acid reaction (forming (forming H2/CO2 respectively). You can measure the gas formed every e.g. 30 seconds and plot the graph and measure the initial gradient in e.g. cm3/min or cm3/sec. The most accurate measurements are made early on in the o reaction when the gas volume versus time is almost linear. You can take a series of measurements and draw the g raph (origin 0,0) 0,0) 3 to get the rate from the gradient (e.g. cm /min) or measure the time to make a fixed volume of gas (* see below). o If the reaction is allowed to go on, you can measure the final maximum volume of gas and the time at which the reaction stops, though this a very poor measure of rate, because the reaction just
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goes slower and slower as the reactant amounts/concentrations are decreasing - so don't use this as a method of measuring reaction speed. (*) The reciprocal of the reaction time, 1/time, 1/time, can also be used as a measure of the speed of a reaction. The time can represent how long it takes to form a fixed amount of gas first few minutes of a metal/carbonate - acid reaction, or the time it takes for so much sulphur to form to obscure the X in the sodium thiosulphate hydrochloric acid reaction. The time can be in minutes or seconds, as long as you stick to the same unit for a set of results e.g. a set of experiments varying the concentration of one of the reactants.
For more details see Advanced Level Chemistry Theory pages on "CHEMICAL KINETICS" Examples of reactions involving gas formation (i) metals dissolving in acid ==> hydrogen gas, (test is lit splint => o pop!), e.g. magnesium + sulphuric acid ==> magnesium sulphate + hydrogen Mg(s) + H2SO4(aq) ==> MgSO4(aq) + H2(g) o (ii) carbonates dissolving in acids => carbon dioxide gas, (test is limewater => cloudy), calcium carbonate (marble chips) + hydrochloric acid ==> calcium chloride + water + carbon dioxide CaCO3(s) + 2HCl(aq) ==> CaCl2(aq) + H2O(l) + CO2(g) o and (iii) the manganese(IV) oxide catalysed decomposition of hydrogen peroxide (oxygen gas, test is glowing splint => relights) hydrogen peroxide ==> water + oxygen 2H2O2(aq) ==> 2H2O(l) + O2(g) can all be followed with the gas syringe method. You can do all sorts of investigations to look o at the effects of (a) the solution concentration, concentration, (b) the temperature of the reactants, (c) the size of the solid particles (surface ( surface area effect), o
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(d) the effectiveness of a catalyst on hydrogen peroxide decomposition.
The shape of the graph is quite characteristic (see diagram above and notes below). o The reaction is fastest at the start when the reactants are at a maximum (steepest gradient in cm3/min). o The gradient becomes progressively less as reactants are used up and the reaction slows down. o Finally the graph levels out when one of the reactants is used up and the reaction stops.
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The amount of product depends on the amount of reactants used . The initial rate of reaction is obtained by measuring the gradient at the start of the reaction. A tangent line is drawn through the first part of the graph, which is usually reasonably linear from the x,y origin 0,0. This gives you an initial rate of reaction in cm 3 gas/minute, Typical results from a gas producing reaction are shown below, for different amounts or concentrations of reactants. How to calculate the reaction rate is explained below. e.g. for run q [ ], after 2 mins, 20 cm 3 of gas formed, so the rate of reaction is 20/2 = 10 cm3/min. From the graph of results you can measure the relative rate of reaction from (i) the initial gradient in cm3/min (see on diagram above), (ii) you can estimate from the graph the volume of gas formed after a particular time e.g. 3 minutes or (iii) you can estimate the time it takes to form a particular volume of gas. (i) is the best method i.e. the best straight line covering several results at the start of the reaction.
Keeping the temperature constant is really important for a 'fair test' if you are investigating speed of reaction/rate of reaction factors such as concentration of a soluble reactant or the particle size/surface area of a solid reactant. On the advanced gas calculations page, temperature sources of error and their correction are discussed in calculation example Q4b.3, although the calculation is a bove GCSE level, the ideas on sources of errors are legitimate for GCSE level. Note that if the temperature of a rates experiment was too low compared to all the other experiments, the 'double error' would occur again, but this time the measured gas volume and the calculated speed/rate of reaction would be lower than expected.
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The rate of a reaction that produces a gas can also be measured by following the mass loss as the gas is formed and escapes from the reaction flask. The method is ok for reactions producing carbon dioxide or oxygen, o o but not very accurate for reactions giving hydrogen (too low a mass loss for accuracy). o The reaction rate is expressed as the rate of loss in mass from the flask in e.g. g/min based on the initial gradient (see graph below).
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When sodium thiosulphate reacts with an acid, a yellow precipitate of sulphur is formed and forms the basis of a good project for assessment.
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To follow this reaction in your investigation you can measure how long it takes for a certain amount of sulphur to form . You do this by observing the reaction down through a conical flask, viewing a black cross on white paper (see diagram below). The X is eventually obscured by the sulphur precipitate and the time noted. sodium thiosulfate + hydrochloric acid ==> sodium chloride + sulfur dioxide + water + sulfur Na2S2O3(aq) + 2HCl(aq) ==> 2NaCl(aq) + SO2(aq) + H2O(l) + S(s) Note: You do not see gas bubbles because the very nasty sulphur dioxide gas is very soluble in water so take care you do not inhale any of the air near the flask when you are doing the experiment or washing out the apparatus afterwards.
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By using the same flask and paper X you can obtain a relative measure of the speed of the reaction in forming the same amount of sulphur . The speed or rate of reaction can expressed as 'x amount of sulphur'/time , so the rate is proportional to 1/time for a particular run of the experiment. In other words since you don't know the absolute mass of sulphur formed, the reciprocal of the time is taken as a measure of the relative rate of reaction. o You can investigate the effects of (a) the hydrochloric acid or sodium thiosulphate concentration (b) the temperature of the reactants. to show the effects of changing one of the variables you can plot graphs of e.g. o reaction time versus temperature or concentration, or rate of reaction (1/reaction time) versus temperature or concentration. You can also measure the speed of this reaction by using a light gate to detect the precipitate formation. The system consists of a light beam emitter and sensor connected to computer and the reaction vessel is placed between the emitter and sensor. The light reading falls as the sulphur precipitate forms. Further examples of graphs that may be obtained from the different methods. For more details see Advanced Level Chemistry Theory pages on "CHEMICAL KINETICS"
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2. The theory of how reactions happen • • •
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WHAT CAUSES A CHEMICAL REACTION? WHAT MUST HAPPEN FOR A CHEMICAL REACTION TO TAKE PLACE? CAN WE MAKE PREDICTIONS ABOUT HOW THE SPEED OF A REACTION MAY CHANGE IF THE REACTION CONDITIONS ARE CHANGED? COLLISION THEORY: Reactions can only happen when the reactant particles collide, but most collisions are NOT successful in forming product molecules despite the high rate of collisions. about 109 per second!) The reason is that particles have a wide range of kinetic energy BUT only a small fraction of particles have enough kinetic energy to break bonds and bring about chemical change. The minimum kinetic energy required for reaction is known as the activation energy. (see also AS-A2 Advanced Theory) The minority high kinetic energy collisions between particles which do produce a chemical change are called 'fruitful collisions'. Here the reactant molecules collide with enough kinetic energy to break the original bonds and form new bonds in the product molecules. Nearly all the rate-controlling factors described below are to do with the collision frequency (chance of collision) OR the energy of reactant particle collision (>= activation energy) which can be summed up as the 'chance of a fruitful collision' leading to product formation. In the case of temperature, the energy of the collision is even more important than the frequency effect (see later ). The particle theory of gases and liquids and the particle diagrams and the explanations below, will all help you understand or describe in your coursework what is going on. For more details see Advanced Level Chemistry Theory pages on "CHEMICAL KINETICS"
3. The Factors affecting the Rate of Chemical Reactions
3a The effect of Concentration (see also graphs 4.6, 4.7 and 4.8) • • • •
WHAT IS THE EFFECT OF CHANGING THE CONCENTRATION OF A REACTANT? AND WHY IS THE REACTION SPEED CHANGED? Why does increase in concentration speed up a reaction? If the concentration of any reactant in a solution is increased, the rate of reaction is increased Increasing the concentration, increases the probability of a collision between o reactant particles because there are more of them in the same volume and so increases the chance of a fruitful collision forming products. e.g. Increasing the concentration of acid molecules increases the frequency or o chance at which they hit the surface of marble chips to dissolve them (slower => faster, illustrated below)
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In general, increasing the concentration of reactant A or B will increase the chance or frequency of a successful collision between them and increase the speed of product formation (slower => faster, illustrated below). •
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Increasing the concentration of reactant A or B will increase the chance or frequency of collision between them and increase the speed of product formation (slower => faster). See also graphs 4.6, 4.7 and 4 .8 for a numerical-quantitative data interpretation. For more details on concentration see Advanced Level Chemistry Theory pages on "CHEMICAL KINETICS" •
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3b The effect of Pressure •
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WHAT IS THE EFFECT OF CHANGING PRESSURE ON THE SPEED OF A REACTION? DOES INCREASING THE PRESSURE ALWAYS HAVE AN EFFECT? Why does an increase in pressure speed up a reaction with a gaseous reactant? If one or more of the reactants is a gas then increasing pressure will effectively increase the concentration of the reactant molecules and speed up the reaction. The particles are, on average, closer together and collisions between the particles will occur more frequently. The A and B particle diagrams above could represent lower/higher pressure, resulting in lesser or greater concentration and so slower or faster reaction all because of the increased chance of a 'fruitful' collision. Solid reactants and solutions are NOT affected by change in pressure, there concentration is unchanged.
3c The effect of Stirring • •
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CAN STIRRING AFFECT THE RATE OF A REACTION? DOES STIRRING AFFECT THE SPEED OF THE REACTION BETWEEN A SOLID AND A SOLUTION? Why does stirring speed up a reaction between a solid and a solution? In doing rate experiments with a solid and solution reactant e.g. marble chips-acid solution or a solid catalyst like manganese(IV) oxide catalysing the decomposition of hydrogen peroxide solution, it is sometimes forgotten that stirring the mixture is an important rate factor . If the reacting mixture is n ot stirred ‘evenly’, the reactant concentration in solution becomes much less near the solid, which tends to settle out at the bottom of the flask. Therefore, at the bottom of the flask the reaction prematurely slows down distorting the overall rate measurement and making the results uneven and therefore inaccurate. The 'unevenness' of the results is even more evident by giving the reaction mixture the 'odd stir'! You get jumps in the graph!!! Stirring cannot affect a completely mixed up solution at the particle level i.e. two solutions of soluble substance that react together are unaffected by stirring.
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3d The effect of Surface Area - particle size of a solid reactant •
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WHAT HAPPENS TO THE SPEED OF A REACTION IF WE CHANGE THE PARTICLE SIZE OF A REACTING SOLID? WHAT DOES BREAKING UP A SOLID REACTANT INTO FINER PIECES DO TO IT IN TERMS OF HOW IT REACTS? If a solid reactant or a solid catalyst is broken down into smaller pieces the rate of reaction increases. The speed increase happens because smaller pieces of the same mass of solid have a greater surface area compared to larger pieces of the solid. Therefore, there is more chance that a reactant particle will hit the s olid surface and react. The diagrams below illustrate the acid–marble chip reaction (slower => faster, but they could also represent a solid catalyst mixed with a solution of reactants. See also graphs 4.1 and 4.8(iii) for a numerical-quantitative data interpretation.
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3e The effect of Temperature (see also graphs 4.3, 4.4 and 4 .8) • • • •
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DOES TEMPERATURE AFFECT THE SPEED OF A CHEMICAL REACTION? IF SO, HOW AND WHY? Why does a reaction go faster at a higher temperature? When gases or liquids are heated the particles gain kinetic energy and move faster (see diagrams below). The increased speed increases the c hance (frequency) of collision between reactant molecules and the rate increases. BUT this is NOT the main reason for the increased reaction speed, so be careful in you r theory explanations if investigating the effect of temperature, so read on after the pictures!
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Most molecular collisions do not result in chemical change . Before any change takes place on collision, the colliding molecules must have a minimum kinetic energy called the Activation Energy shown on the energy level diagrams below (sometimes called reaction profile/progress diagrams - shown below). Going up and to the top 'hump' represents bond breaking on reacting particle o collision. The purple arrow up represents this minimum energy needed to break bonds to initiate the reaction, that is the activation energy. o Going down the other side represents the new bonds formed in the reaction products. The red arrow down represents the energy released - exothermic reaction. It does not matter whether the reaction is an exothermic or an endothermic in terms of energy change, its the activation energy which is the most important factor in terms of temperature and reaction speed. Now heated molecules have a greater average kinetic energy, and so at higher temperatures, a greater proportion of them have the required activation energy to react. This means that the increased chance of 'fruitful' higher energy collision greatly increases the speed of the reaction, depending on the fraction of molecules with enough energy to react. For this reason, generally speaking, and in the absence of catalysts or extra energy input, a low activation energy reaction is likely to be fast and a high activation energy reaction much slower , reflecting the trend that the lower the energy barrier to a reaction, the more molecules are likely to have sufficient energy to react on collision.
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See also graphs 4.3, 4.4 and 4.8 for a numerical-quantitative data interpretation. Trying to resolve an apparent confusion for GCSE students!
1. With increase in temperature, there is an increased frequency (or chance) of collision 2.
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due to the more 'energetic' situation - but this is the minor factor when considering why rate of a reaction increases with temperature. The minimum energy needed for reaction, the activation energy (to break bonds on collision), stays the same on increasing temperature. However, the average increase in particle kinetic energy caused by the absorbed heat means that a much greater proportion of the reactant molecules now has the minimum or activation energy to react. It is this increased chance of a 'successful' or 'fruitful' higher energy collision leading to product formation, that is the major factor, and this effect increases more than the increased frequency of particle collision, for a similar rise in temperature. This is usually only fully d iscussed at AS-A2 level, but it may impress the teacher for GCSE coursework if you look up the Maxwell-Boltzmann distribution of kinetic energies, its quite difficult to get over some of these ideas without considering graphs of probability versus particle KE, but that's up to you! There is also the Arrhenius Equation relating rate of reaction and temperature - but this involves advanced level mathematics. For more details see Advanced Level Chemistry Theory pages on "CHEMICAL KINETICS"
3f The effect of a Catalyst (see also light effect and graph 4.8)
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WHAT IS A CATALYST? HOW DOES IT AFFECT THE SPEED OF A CHEMICAL REACTION? HOW DOES A CATALYST WORK? Why does a catalyst speed up a reaction? I was once asked "what is the opposite of a catalyst? There is no real opposite to a catalyst, other than the uncatalysed reaction! The word catalyst means an added substance, in contact with the reactants, that changes the rate of a reaction without itself being chemically changed in the end. There are the two phrases you may come across: o a 'positive catalyst' meaning speeding up the reaction (plenty of examples in most chemistry courses) o OR a 'negative catalyst' slowing do wn a reaction (rarely mentioned at GCSE, sometimes at AS-A2 level, e.g. adding a chemical that 'mops up' free radicals or other reactive species). Catalysts increase the rate of a reaction by helping break chemical bonds in reactant molecules and provide a 'different pathway' for the reaction. This effectively means the Activation Energy is reduced, irrespective of whether its an exothermic or endothermic reaction (see diagrams below).
Therefore at the same temperature, more reactant molecules have enough kinetic energy to react compared to the uncatalysed situation. The catalyst does NOT increase the energy of the reactant molecules! Although a true catalyst does take part in the reaction and may change chemically temporarily, but it does not get used up and can be reused/regenerated with more reactants. It does not change chemically or get used up in the end . Black manganese(IV) oxide (manganese dioxide) catalyses the decomposition of o hydrogen peroxide. hydrogen peroxide ==> water + oxygen o 2H2O2(aq) ==> 2H2O(l) + O2(g)
The manganese is chemically the same at the end of the reaction but it may change a little physically if its a solid e.g. o In the hydrogen peroxide solution decomposition by the solid black catalyst manganese dioxide, the solid can be filtered off when reaction stops 'fizzing' i.e. all of the hydrogen peroxide has reacted-decomposed. o After washing with water, the catalyst can be collected and added to fresh colourless hydrogen peroxide solution and the oxygen production 'fizzing' is instantaneous! In other words the catalyst hasn't changed chemically and is as effective as it was fresh from the bottle! Note: At the end of the experiment the solution is sometimes stained brown from minute manganese dioxide particles. The reaction is exothermic and the heat has probably caused some d isintegration of the catalyst into much finer particles which appear to be (but not) dissolved. In other words the catalyst has changed physically BUT NOT chemically. Different reactions need different catalysts and they are extremely important in industry: examples .. o nickel catalyses the hydrogenation of unsaturated fats to margarine iron catalyses the combination of unreactive nitrogen and hydrogen to form o ammonia o enzymes in yeast convert sugar into alcohol zeolites catalyse the cracking of big hydrocarbon molecules into smaller ones o o most polymer making reactions require a catalyst surface or additive in contact with or mixed with the monomer molecules. o
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Enzymes are biochemical catalysts are dealt with on another page enzymes and biotechnology They have the advantage of bringing about reactions at normal o temperatures and pressures which would otherwise need more expensive and energy-demanding equipment. For more details on catalysis see "CHEMICAL KINETICS"
Advanced Level Chemistry Theory pages on
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CAN LIGHT AFFECT THE SPEED OF ANY REACTIONS? IF IT DOES, HOW DOES CHANGE THE SPEED OF A CHEMICAL REACTION? Why does increasing light intensity sometimes increase the speed of a reaction? Light energy (uv or visible radiation) can initiate or catalyse particular chemical reactions. As well as acting as an electromagnetic wave, light can be considered as an o energy 'bullets' called photons and they have sufficient 'impact energy' to break chemical bonds, that is, enough energy to overcome the activation energy. The greater the intensity of light (visible or ultra-violet) the more reactant o molecules are likely to gain the energy react, so the reaction speed increases. Examples: o Silver salts are converted to silver in the chemistry of photographic exposure of the film. Silver chloride (AgCl), silver bromide (AgBr) and silver iodide (AgI) are all sensitive to light ('photosensitive'), and all three are used in the
production of various types of photographic film to detect visible light and beta and gamma radiation from radioactive materials. Each silver halide salt has a different sensitivity to light. When radiation hits the film the silver ions in the salt are reduced by electron gain to silver Ag+ + e- ==> Ag (X = halogen atom, Cl, Br or I) and the halide ion is oxidised to the halogen molecule by electron loss 2X- ==> X2 + 2eso overall the change via light energy is: 2AgX ==> 2Ag + X2 AgI is the least sensitive and used in X-ray radiography, AgCl is the most sensitive and used in 'fast' film for cameras. Photosynthesis in green plants: The conversion of water + carbon dioxide ==> glucose + oxygen 6H2O(l) + 6CO2(g) ==> C6H12O6(aq) + 6O2(g) requires the input of sunlight energy and the green chlorophyll molecules absorb the photon energy packets of light and initiate the chemical changes summarised above. Photochemical Smog: This is very complex chemistry involving hydrocarbons, carbon monoxide, ozone, nitrogen oxides etc. Many of the reactions to produce harmful chemicals are catalysed or promoted by light energy.
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4. More examples of interpreting graphical results ('graphing'!) PLOTTING GRAPHS - PLOTS OF GRAPHS OF DATA AND HOW TO INTERPRET THEM PLEASE Note (i) rate of reaction = speed, (ii) see two other graphs and notes (ii) Graphs 4.1, 4.2 and 4.5 just show the theoretical shape of a graph for a single particular experiment. Graphs 4.3 and 4.4 (temperature), 4.6 and 4.7 (concentration) and 4.8 (several factors illustrated) shows the effect of changing a variable on the rate of the reaction and hence the relative change in the curve-shape of the graph line. (iii) The rate of reaction may be expressed as the reciprocal of the reaction time (1/time) e.g. for the time for sulphur formation (to obscure the X) in the sodium thiosulfate - hydrochloric acid reaction or where a fixed volume of gas is formed, though in this can also be expressed as gas volume/time too as cm3 /s or cm 3 /min even though the gas volume is the same for a given set of results of changing one variable whether it be concentration or temperature.
If you have detailed data e.g. multiple gas volume readings versus time, the best method for rate analysis is the initial rate method described on and below the diagram of the gas syringe gas collection system. (iv) for detailed observations of gas versus time
Graph 4.1 shows the decrease in the amount of a solid reactant with time. The graph is curved, becoming less steep as the gradient decreases because the reactants are being used up, so the speed decreases. Here the gradient is a measure of the rate of the reaction. In the first few minutes the graph will (i) decline less steeply for larger 'lumps' and (ii) decline more steeply with a fine powder i.e. (i) less surface area gives slower reaction and (ii) more surface area a faster reaction. Graph 4.2 shows the increase in the amount of a solid product with time. The graph tends towards a maximum amount possible when all the solid reactant is used up and the graph becomes horizontal. This means the speed has become zero as the reaction has stopped. Here the gradient is a measure of the rate of the reaction. Graph 4.3 shows the decrease in reaction time with increase in temperature as the reaction speeds up. The reaction time can represent how long it takes to form a fixed amount of gas in e.g. in the first few minutes of a metal/carbonate - acid reaction, or the time it takes for so much sulphur to form in the sodium thiosulphate hydrochloric acid reaction. The time can be in minutes or seconds, as long as you stick to the same unit for a set of results e.g. a set of experiments varying the concentration of one of the reactants. Theory of temperature effect Graph 4.4 shows the increase in speed of a reaction with increase in temperature as the particles have more and more kinetic energy. The rate of reaction is proportional to 1/t where t = the reaction time. See the notes on rate in the Graph 4.7 paragraph below and the theory of temperature effect. Graph 4.5 shows the increase in the amount of a gas formed in a reaction with time. Here the gradient is a measure of the rate of the reaction. Again, the graph becomes horizontal as the reaction stops when one of the reactants is all used up! More on this type of graph.
Graph 4.6 shows the effect of increasing concentration, which decreases the reaction time, as the speed increases because the greater the concentration the greater the chance of fruitful collision. See the notes on rate in the Graph 4.3 paragraph above and the theory of concentration effect Graph 4.7 shows the rate/speed of reaction is often proportional to the concentration of one particular reactant. This is due to the chance of a fruitful collision forming products being proportional to the concentration. The initial gradient of the product-time graph e.g. for gas in cm3 /min (or s for timing the speed/rate) gives an accurate measure of how fast the gaseous product is being formed. The reciprocal of the reaction time, 1/time, can also be used as a measure of the speed of a reaction. The time can e.g. represent how long it takes to make a fixed amount of gas, or the time it takes for so much sulphur to form in the sodium thiosulphate - hydrochloric acid reaction. The
time can be in minutes or seconds, as long as you stick to the same unit for a set of results for a set of experiments varying the concentration or mass of one of the reactants. Theory of concentration effect Graph 4.8 A set of results for the same reaction (i) The graph lines W, X, original, Y and Z on the left diagram are typical of when a gaseous product is being collected. The middle graph might represent the original experiment 'recipe' and temperature. Then the experiment repeated with variations e.g. (ii) X could be the same recipe as the original experiment but a catalyst added, forming the same amount of product, but faster. (iii) Initially, the increasing order of rate of reaction represented on the graph by curves Z to W i.e. W > X > original > Y > Z might represent progressively increasing concentrations of reactant or progressively higher temperature of reaction or progressively smaller lumps-particle/increasing surface area of a solid reactant. All three trends in changing a reactant/reaction condition variable produce a progressively faster reaction shown by the increasing gradient in cm 3 /min which represents the rate/speed of the reaction. (iv) Z could represent taking half the amount of reactants or half a concentration. The reaction is slower and only half as much gas is formed. (v) W might represent taking double the quantity of reactants , forming twice as much gas e.g. same volume of reactant solution but doubling the concentration, so producing twice as much gas, initially at double the speed (gradient twice as steep).
See also graphs for enzymes showing effects of pH, temperature and concentration.
For more details on concentration results analysis see on "CHEMICAL KINETICS"
Advanced Level Chemistry Theory pages
This page should help with ideas about factors controlling the rates of chemical reactions for coursework projects, assignments, investigations etc.
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AND on other associated web pages: o
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Enzymes and biotechnology (rates factors, temperature and pH effects, lock-key mechanism) Ammonia synthesis (applying rates and equilibrium factors) Uses of catalysts in ...
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Transition Metals, control
Industrial Chemistry and
Pollution
Energy changes in chemistry RATES QUIZ QUESTIONS and WORKSHEETS printout or online:
GCSE/IGCSE MULTIPLE CHOICE QUIZ on RATES of reaction
MULTI-WORD GAP-FILL QUIZ on Rates of Reaction MATCHING PAIRS QUIZ on rates of reaction Q1 and Q2
CROSSWORD PUZZLE on rates of reaction and ANSWERS! 10 JUMBLED SENTENCES to sort out! See also the brainstorm of GCSE rates coursework-projects investigation ideas and two
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Advanced Level Chemistry Theory pages on "CHEMICAL KINETICS" ADVANCED LEVEL theory pages on 'Kinetics' for GCE-AS-A2-IB/US gr 1112) and adventurous GCSE students!
http://www.docbrown.info/page03/3_31rates.htm
What factors influence the rate of a chemical reaction?
Temperature Concentrations of reactants Catalysts Surface area of a solid reactant Pressure of gaseous reactants or products If you are planning an investigation, I suggest that you investigate the effects of temperature or the effects of the concentration of the reactants because these will allow you to choose a suitable range of values for the controlled or independent variable. The dependent variable will be the rate of the reaction. Keep all the other variables fixed. To make a prediction for your investigation you will have to ask yourself the question: What will happen to the rate of the reaction when I increase the temperature? or What will happen to the rate of the reaction if I increase the concentration of one of the reactants? The answer to that question is your prediction. The next thing to do is to explain your prediction. You will have to answer the question: Why will the reaction go faster if I increase the temperature? or Why will the reaction go faster if I increase the concentration of one of the reactions? The answer to this question is your explanation, and to get the highest possible marks, you will have to provide a full scientific explanation. Once you have written your hypothesis (prediction with explanation) you will decide how to do the experiments, i.e. write the proposed method.
How does temperature affect the rate of a chemical reaction?
When two chemicals react, their molecules have to collide with each other with sufficient energy for the reaction to take place. This is collision theory. The two molecules will only react if they have enough energy. By heating the mixture, you will raise the energy levels of the molecules involved in the reaction. Increasing temperature means the molecules move faster. This is kinetic theory. If your reaction is between atoms rather than molecules you just substitute "atom" for "molecule" in your explanation. How do catalysts affect the rate of a reaction?
Catalysts speed up chemical reactions. Only very minute quantities of the catalyst are required to produce a dramatic change in the rate of the reaction. This is really because the reaction proceeds by a different pathway when the catalyst is present. Adding extra catalyst will make absolutely no difference. There is a whole page on this site devoted to catalysts. How does concentration affect the rate of a reaction?
Increasing the concentration of the reactants will increase the frequency of collisions between the two reactants. So this is collision theory again. You also need to discuss kinetic theory in an experiment where you vary the concentration. Although you keep the temperature constant, kinetic theory is relevant. This is because the molecules in the reaction mixture have a range of energy levels. When collisions occur, they do not always result in a reaction. If the two colliding molecules have sufficient energy they will react. If reaction is between a substance in solution and a solid, you just vary the concentration of the solution. The experiment is straightforward. If the reaction is between two solutions, you have a slight problem. Do you vary the concentration of one of the reactants or vary the concentration of both? You might find that the rate of reaction is limited by the concentration of the weaker solution, and increasing the concentration of the other makes no difference. What you need to do is fix the concentration of one of the reactants to excess. Now you can increase the concentration of the other solution to produce an increase in the rate of the reaction. How does surface area affect a chemical reaction?
If one of the reactants is a solid, the surface area of the solid will affect how fast the reaction goes. This is because the two types of molecule can only bump into each other at the liquid solid interface, i.e. on the surface of the solid. So the larger the surface area of the solid, the faster the reaction will be. Smaller particles have a bigger surface area than larger particle for the same mass of solid. There is a simple way to visualize this. Take a loaf of bread and cut it into slices.
Each time you cut a new slice, you get an extra surface onto which you can spread butter and jam. The thinner you cut the slices, the more slices you get and so the more butter and jam you can put on them. This is "Bread and Butter Theory". You should have come across the idea in your biology lessons. By chewing your food you increase the surface area so that digestion can go faster. What affect does pressure have on the reaction between two gasses?
You should already know that the atoms or molecules in a gas are very spread out. For the two chemicals to react, there must be collisions between their molecules. By increasing the pressure, you squeeze the molecules together so you will increase the frequency of collisions between them. This is collision theory again. In a diesel engine, compressing the gaseous mixture of air and diesel also increases the temperature enough to produce combustion. Increasing pressure also results in raising the temperature. It is not enough in a petrol engine to produce combustion, so petrol engines need a spark plug. When the petrol air mixture has been compressed, a spark from the plug ignites the mixture. In both cases the reaction (combustion) is very fast. This is because once the reaction has started, heat is produced and this will make it go even faster.
Factors influencing rate of reaction The nature of the reaction: Some reactions are naturally faster than others. The number of reacting species, their physical state (the particles that form solids move much more slowly than those of gases or those in solution), the complexity of the reaction and other factors can influence greatly the rate of a reaction. Concentration: Reaction rate increases with concentration, as described by the rate law and explained by collision theory. As reactant concentration increases, the frequency of collision increases. Pressure: The rate of gaseous reactions increases with pressure, which is, in fact, equivalent to an increase in concentration of the gas. For condensed-phase reactions, the pressure dependence is weak. Order : The order of the reaction controls how the reactant concentration (or pressure) affects reaction rate. Temperature: Usually conducting a reaction at a higher temperature delivers more energy into the system and increases the reaction rate by causing more collisions between particles, as explained by collision theory. However, the main reason that •
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temperature increases the rate of reaction is that more of the colliding particles will have the necessary activation energy resulting in more successful collisions (when bonds are formed between reactants). The influence of temperature is described by the Arrhenius equation. As a rule of thumb, reaction rates for many reactions double for every 10 degrees Celsius increase in temperature,[2] though the effect of temperature may be very much larger or smaller than this. For example, coal burns in a fireplace in the presence of oxygen but it doesn't when it is stored at room temperature. The reaction is spontaneous at low and high temperatures but at room temperature its rate is so slow that it is negligible. The increase in temperature, as created by a match, allows the reaction to start and then it heats itself, because it is exothermic. That is valid for many other fuels, such as methane, butane, and hydrogen. Reaction rates can be independent of temperature (non-Arrhenius) or decrease with increasing temperature (anti-Arrhenius). Reactions without an activation barrier (e.g. some radical reactions), tend to have anti Arrhenius temperature dependence: the rate constant decreases with increasing temperature. Solvent : Many reactions take place in solution and the properties of the solvent affect the reaction rate. The ionic strength also has an effect on reaction rate. Electromagnetic radiation and intensity of light : Electromagnetic radiation is a form of energy. As such, it may speed up the rate or even make a reaction spontaneous as it provides the particles of the reactants with more energy. This energy is in one way or another stored in the reacting particles (it may break bond s, promote molecules to electronically or vibrationally excited states...) creating intermediate species that react easily. As the intensity of light increases, the particles absorb more energy and hence the rate of reaction increases. •
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For example when methane reacts with chlorine in the dark, the reaction rate is very slow. It can be sped up when the mixture is put under diffused light. In bright sunlight, the reaction is explosive. A catalyst : The presence of a catalyst increases the reaction rate (in both the forward and reverse reactions) by providing an a lternative pathway with a lower activation energy. •
For example, platinum catalyzes the combustion of hydrogen with oxygen at room temperature. Isotopes: The kinetic isotope effect consists in a different reaction rate for the same molecule if it has different isotopes, usually hydrogen isotopes, because of the mass difference between hydrogen and deuterium. Surface Area: In reactions on surfaces, which take place for example during heterogeneous catalysis, the rate of reaction increases as the surface area does. That is because more particles of the solid are exposed and can be hit by reactant molecules. •
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