Doc Brown's Chemistry Clinic

KS4 Science GCSE-IGCSE Chemistry revision information-notes on

Factors affecting the Speed-Rates of Chemical Reactions

  see also the brainstorm of GCSE rates coursework-projects investigation ideas and two ADVANCED theory pages on 'Kinetics' for GCE-AS-A2-IB/US gr 11-12) and adventurous GCSE students! * EMAIL query?comment

SECTIONS on this page: 1. What do mean by rate/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

KEY WORDS-PHRASES in alphabetical order for this rates web page: hydrochloric/sulphuric acid-metal e.g. Mg/carbonate reaction * hydrochloric acid-sodium thiosulphate reaction * Activation energy * Catalysts * Concentration 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  

• AND on other associated web pages:

  • Enzymes and biotechnology (rates factors, temperature and pH effects, lock-key mechanism)

  • Ammonia synthesis (applying rates and equilibrium factors)

  • Uses of catalysts in Transition Metals and Industrial Chemistry and Pollution control

  • Energy changes in chemistry

  • RATES QUIZ QUESTIONS:

    • MULTIPLE CHOICE QUIZ * SHORT ANSWER QUIZ * MULTI-WORD FILL QUIZ

    • and MATCHING PAIRS QUIZ Q1 and Q2, drop down menu or drag/drop and printout

      • CROSSWORD PUZZLE (printout  version) * JUMBLED SENTENCES

1. What do we mean by Rate and how is it measured?

• The phrase ‘rate of reaction’ means ‘how fast is the reaction’ or 'the speed of the reaction'. It can be measured as the 'rate of formation of product' (e.g. collecting gaseous product in a syringe) or the 'rate of removal of reactant'. The speeds of reactions are very varied.

  • Rusting is a ‘slow’ reaction, you hardly see any change looking at it!

  • The weathering of rocks is a very slow reaction.

  • The fermentation of sugar to alcohol is quite slow but you can see the carbon dioxide bubbles forming in the 'froth'!

  • A fast 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.

    • 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.

      • However, the two substances mentioned above are formed in a chemical reaction change, 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.

  • Explosions and burning/combustion reactions would be described as ‘very fast’!

• The importance of "Rates of Reaction knowledge":

  • Time is money in 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.

  • Health and Safety 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.
• 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 below.
• 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).
  • The initial gradient of the graph e.g. in cm³/min (speed/rate) gives an accurate measure of how fast a gaseous product is being formed in  meta/carbonate - acid reaction (forming H₂/CO₂ respectively).
  • The most accurate measurements are made early on in the reaction when the gas volume versus time is almost linear. You can take a series of measurements and draw the graph (origin 0,0) to get the rate from the gradient (e.g. cm³/min) or measure the time to make a fixed volume of gas.
  • 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 goes slower and slower as the reactant amounts/concentrations are decreasing.
  • The reciprocal of the reaction 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 meta/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.
• Reactions involving:
  • (i) metals dissolving in acid ==> hydrogen gas, (test is lit splint => pop!),
    • e.g. magnesium + sulphuric acid ==> magnesium sulphate + hydrogen
      • Mg_(s) + H₂SO_4(aq) ==> MgSO_4(aq) + H_2(g)
  • (ii) carbonates dissolving in acids => carbon dioxide gas, (test is limewater => cloudy),
    • calcium carbonate (marble chips)  + hydrochloric acid ==> calcium chloride + water + carbon dioxide
      • CaCO_3(s) + 2HCl_(aq) ==> CaCl_2(aq) + H₂O_(l) + CO_2(g)
  • and (iii) the manganese(IV) oxide catalysed decomposition of hydrogen peroxide (è oxygen gas, test is glowing splint => relights)
    • hydrogen peroxide ==> water + oxygen
      • 2H₂O_2(aq) ==> 2H₂O_(l) + O_2(g)
        • can all be followed with the gas syringe method.
  • You can do all sorts of investigations to look at the effects of 
    • (a) the solution concentration,
    • (b) the temperature of the reactants,
    • (c) the size of the solid particles (surface area effect),
    • (d) the effectiveness of a catalyst on hydrogen peroxide decomposition.

• The shape of the graph is quite characteristic (see diagram above and notes below).
  • The reaction is fastest at the start when the reactants are at a maximum (steepest gradient in cm³/min).
  • The gradient becomes progressively less as reactants are used up and the reaction slows down.
  • Finally the graph levels out when one of the reactants is used up and the reaction stops.
  • 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 cm³ 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 cm³ of gas formed, so the rate of reaction is 20/2 = 10 cm³/min.
    • 
    • 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 above 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.

• 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,
  • but not very accurate for reactions giving hydrogen (too low mass loss for accuracy).
  • The reaction rate is expressed as the rate of loss in mass from the flask in e.g. g/min based on the intitial gradient (see graph below).

• When sodium thiosulphate reacts with an acid, a yellow precipitate of sulphur is formed and forms the basis of a good project for assessment.
  • 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
  • Na₂S₂O_3(aq) + 2HCl_(aq) ==> 2NaCl_(aq) + SO_2(aq) + H₂O_(l) + S_(s)
    • Note: You do not see gas bubbles because the very nasty sulphur dioxide gas is very soluble in water.

mixongoingwatch stopped

• 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.
  • 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.
    • 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.

• 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 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 chance at which they hit the surface of marble chips to dissolve them (slower => faster, illustrated below)

• 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).

• 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.

• 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 A and B particle diagrams above could represent lower/higher pressure , resulting in lesser è greater concentration and so slower è faster reaction all because of the increased chance of a 'fruitful' collision.

• 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 not 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!!!

• 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 solid 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.

• When gases or liquids are heated the particles gain kinetic energy and move faster (see diagrams below).

• The increased speed increases the chance (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 your theory explanations if investigating the effect of temperature, so read on after the pictures!

• 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 collision.

    • The purple arrow up represents this minimum energy needed to break bonds to initiate the reaction, that is the activation energy.

  • 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.

• 

  • 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 due to the more 'energetic' situation - but this is the minor factor when considering why rate of a reaction increases with temperature.
2. 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. 
3. 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.
4. This is usually only fully discussed 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.

3f The effect of a Catalyst (see also light effect and graph 4.8)

• 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:

  • a 'positive catalyst' meaning speeding up the reaction (plenty of examples in most chemistry courses)

  • OR a 'negative catalyst' slowing down a reaction (rarely mentioned at GCSE, sometimes at AS-A2 level, e.g. adding a chemical that 'mops up' free radicals or 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 hydrogen peroxide.
  • hydrogen peroxide ==> water + oxygen
    • 2H₂O_2(aq) ==> 2H₂O_(l) + O_2(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.

  • 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.

  • 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 disintegration 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 ..

  • nickel catalyses the hydrogenation of unsaturated fats to margarine

  • iron catalyses the combination of unreactive nitrogen and hydrogen to form ammonia

  • enzymes in yeast convert sugar into alcohol

  • zeolites catalyse the cracking of big hydrocarbon molecules into smaller ones

  • most polymer making reactions require a catalyst surface or additive in contact with or mixed with the monomer molecules.

• Enzymes are biochemical catalysts are dealt with on another page - enzymes and biotechnology.

  • They have the advantage of bringing about reactions at normal temperatures and pressures which would otherwise need more expensive and energy-demanding equipment.

---

3g The Effect of Light

• 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 energy 'bullets' called photons and they have sufficient 'impact' 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 molecules are likely to gain the energy react, so the reaction speed increases.

• Examples:

  • 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 + 2e^-

      • so 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) + 6CO_2(g) ==> C₆H₁₂O_6(aq) + 6O_2(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.

---

4. More examples of interpreting graphical results ('graphing'!)

Note (i) rate of reaction = speed, (ii) see two other graphs and notes

Graphs 4.1 to 4.7 just show the shape of graph for an individual run of results, but graph 4.8 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.

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 cm3/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.

ks4 science examinations gcse-igcse chemistry revision *  ks4 science examinations-gcse-igcse chemistry revision *  ks4 science examinations-gcse-igcse chemistry revision *  ks4 science examinations-gcse-igcse chemistry revision *  ks4 science examinations-gcse-igcse chemistry revision *  ks4 science examinations-gcse-igcse chemistry revision * SITE PURPOSE EDUCATION - online learning or 'self-private-tuition' using revision notes, quizzes, practice tests involving GCSE Science CHEMISTRY in the areas of REVISING only the CHEMISTRY-Earth Science-Radioactivity at Doc Brown's Chemistry Clinic via HOMEPAGE in secondary school/schools, 6th form college/colleges, academy/academies or home self-study. Hopefully it will encourage interest and understanding of Chemistry, Earth Science and Radioactivity in any country of the world, though the site is written entirely in English. The website is designed to help and unofficially support students/teachers revise-learn/teach the chemistry for modular or co-ordinated examination science courses from UK QCA based AQA, OCR (Oxford and Cambridge) Twenty First (21st) Century and Gateway Science, Edexcel 360Science , Nuffield, Salters, Cambridge International (CIE), London International, WJEC, CCEA exams etc. Also, national award assessments-examinations for GCSE-IGCSE-KS4-O level-BTEC-NVQ applied, additional and chemistry national science courses. Also covers, mainly via quizzes the UK National KS3 SATs Science-biology/chemistry/physics (SAT revision levels 3-5 or 5-7) and covers much of the revising, learning and teaching chemistry examinations for the national curriculum for secondary schools and colleges. The site does not support the content of England, Wales or Northern Ireland primary science KS1 or KS2. The notes should also provide some background theory for a coursework assignment or project. BUT please note that my on-line revision notes and quizzes are no substitute for good classroom teaching-lecturing and thorough studying of your own notes and textbooks, practicing past papers and a copy of the syllabus which are readily downloaded from the examination board sites, but I hope here and there they will lend a tutoring hand on some topic, unit, module etc. For final revision you have to be intellectually honest about what you don't know or follow, YOU have to take the stuff to pieces, analyse what you do/do not understand and reconstruct it so it all makes sense in the end. There is no other way, there are no magic secrets on how to revise and learn, its mainly down to hard work and just good old fashioned study and employing teach-yourself strategies without the need for extra tutors and tutoring lessons. I also think there is too much hit and miss revision using past papers (which I do NOT supply) and not enough systematic revision. I also hope it will help teachers in planning lessons and developing schemes of work for science-chemistry. There are no lesson plans on the site but there are plenty of quizzes to incorporate into classroom activities whether photocopied or on electronic whiteboard projector for use as self-tuition-assessment purposes and a variety of teaching and learning styles and the images may be used in Microsoft Word documents and powerpoint projections. The site seems to be used by a large number of home study tutors, particularly the revision notes. An individual tutor may print out the notes for science-chemistry learning teaching-tuition purposes and for background material for assignments and projects. I have no interest or time in producing WORD.doc or xxxx.pdf files of the notes at the moment. Neither have I time to write up many practical laboratory experiments ('lab'-'labs') at the moment, but the notes contain lots of background information of chemical reactions in terms of observations-balanced equations-reactants-products-theory etc. I also find it difficult to recommend specific exam websites or syllabus textbooks, it depends exactly on what you need, what you have time for, and there are so many of them to choose from and I do not supply past examination papers for classes. The sites resources include revision notes, quizzes and worksheets which provide support for home study or tuition for homework and coursework help e.g. science investigations for any of the key stage courses indicated, but I do not supply lesson plans.  Dr W P Brown gcse 10-11-2007 *  ks4 science examinations gcse-igcse chemistry revision *  ks4 science examinations-gcse-igcse chemistry revision *  ks4 science examinations-gcse-igcse chemistry revision *  ks4 science examinations-gcse-igcse chemistry revision *  ks4 science examinations-gcse-igcse chemistry revision *  ks4 science examinations-gcse-igcse chemistry revision

Online free help resources for Key Stages 3 SATs (S.A.T.s), 4 & 5AQA, Edexcel, OCR, CIE GCSE IGCSE BTEC Science, GCE, AS, A2 Advanced subsidiary Chemistry A levels, IB Diploma and US K12 (K-12 grades) courses and examinations and revising for the various syllabuses and specifications. Exploring the site for lessons, plans, ideas for projects and coursework, professional development. Through hard work the site has been built up over the course of many years with no need of special pc software except FrontPage and Hot Potatoes (uvic) for quizzes and worksheets. It is used in the classroom, home learning-tutoring-schooling and guidance, private tuition, school retakes revision. Whether you are a teacher/tutor teaching, a student studying, using the pages as self-study guides for your science-chemistry studies etc. etc. I hope the site supports your endeavour. 15-12-07 © Dr W P Brown

This page should help with rates of reaction coursework projects or assignments.

KS4 SCIENCE - Additional & Applied Chemistry help

AQA GCSE Science - Chemistry

CCEA GCSE Science - Chemistry

Edexcel GCSE 360Science - Chemistry

OCR GCSE 21st Century Science Suite - Chemistry

OCR GCSE Gateway Science Suite - Chemistry

OCR GCSE Applied Science - Chemistry (double award)

WJEC GCSE Science - Chemistry