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Chem 40S Unit 3 Notes
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Chem 40S Unit 3 Notes

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  • i) Temperature change over time (°C/min) As the reaction proceeds, the system will increase in temperature. The faster heat is produced, the higher the rate ii) Pressure change over time (kPa/s or mmHg/s) As the reaction proceeds, more gaseous C is produced. This increases the pressure of the system. The faster the pressure increases, the greater the rate. iii) Mass change over time (g of C/min) As the reaction proceeds, reactants are used up and converted to gas. This results in a loss of mass of reactants, or an increase in mass of product, C. iv) Other possibilities: Colour change, pH change, change in conductivity, etc. over a period of time.
  • Where:  [A] is the change in the concentration of A, etc.   t is the change in time
  • Covalent bonds involve the sharing of one or more electron pairs. These are usually much stronger than ionic bonds, which involve electrostatic interactions caused by electron transfer. Since covalent bonds are very strong they take more time to break. Ionic reactions are often instantaneous.
  • Changes in temperature affect the shape of the Kinetic Distribution Curve: If activation energy and number of particles remain constant, increasing temperature from T 2 to T 3 , shifts the curve to the right. Since the activation energy does not change, there are more particles with enough energy, activation energy, to produce an effective collision. Conversely, a decrease in the temperature to T 1 , shifts the curve to the left, decreasing the number of particles with activation energy. This decreases the reaction rate. In the example above, the reaction does not proceed at T 1 as there are little or no particles with activation energy.
  • We can think of this in terms of collisions with cars as well. During rush hour, there are more accidents because there are more vehicles on the road at one time. The frequency of collisions increases because the "concentration" of vehicles is larger and the space between each vehicle is less, increasing chances of collision.
  • If you decrease the volume of a container without changing the number of particles in the container, as in figure c in the diagram below, the concentration of the reactants increases. The spaces between the particles decreases, increasing the chances of a collision.  If the concentration of the reactants increases, the reaction rate increases. If you increase the volume of a container without changing the number of particles, as in figure a in the diagram above, the concentration of the reactants decreases. The spaces between the particles increases, decreasing the chances of a collision.  Decreasing the concentration of reactants decreases the reaction rate.
  • Here is another quick little experiment you may want to try (and this is one you can safely do): Find a friend, and each of you crumple up a piece of paper. Sitting facing one another, with knees almost touching, toss your pieces of paper together, trying to get them to hit while still in the air. Do this 10 times - how many successful "collisions" did you record? Now, grab another friend with their own wad of paper. This time, a successful collision will only occur when all three pieces of paper hit together simultaneously - just two out of three won't do. Out of 10 tosses, how many are successful? Add one more friend and repeat the exercise. Any successful collisions, where all four pieces of paper collide at the same instant? Likely not.

Chem 40S Unit 3 Notes Chem 40S Unit 3 Notes Presentation Transcript

  • Chemistry 40S Unit 3 - Kinetics
  • About KineticsChemistry focuses largely on chemicalreactions. Some reactions, such as rusting,occur very slowly, while others, likeexplosions, occur quite fast. Kinetics is thebranch of chemistry that studies the speedor rate with which chemical reactions occur.Some reactions do not occur in one simplestep. Some occur in several more complexsteps. Part of Kinetics is studying the steps,or mechanism, of a chemical reaction.
  • Defining Rate• Recall that a chemical reaction is one in which a new substance is formed. They are usually expressed in the shorthand form, known as a chemical equation. REACTANTS → PRODUCTS• The rate of a reaction refers to the speed with which a chemical reaction takes place. That is, how fast reactants are used up or products are formed.
  • Measuring Rate • Rate can be measured using several different methods. • The reaction below shows that A and B combine to form C and heat is given off: A(s) + B(l) → C(g) + heat • In this reaction, rate can be measured in terms of several different properties: – i) Temperature change over time (°C/min) – ii) Pressure change over time (kPa/s or mmHg/s) – iii) Mass change over time (g of C/min) – iv) Other possibilities:
  • Measuring RateTherefore generally,Rate is usually described in terms of change in concentrationof reactant or product over time.Or, in the case of our example: A(s) + B(l) → C(g) + heat
  • Average Rate
  • Average Rate• The numerical value for the rate of a reaction can be determined by examining the change in the amount of a substance at a particular time, or over a period of time. The graphical representation of this may look like the figure below:
  • Average Rate
  • Average Rate• The average rate for a reaction is given by the following equation:• The instantaneous rate of a reaction is the rate at a specific time. This is determined by calculating the slope of the line tangent to the point on the concentration vs time curve.
  • Average Rate• Example 1 – According to the reaction A → B, the following data was collected: a) What is the average rate over the entire 50 seconds? b) What is the average rate for the interval 20 s to 40 s?
  • Average Rate• Example 2 – The decomposition of nitrogen dioxide produces nitrogen monoxide and oxygen according to the reaction: 2 NO2(g) → 2 NO(g) + O2(g) – the following data has been collected: Calculate: a) The average rate of decomposition of NO2 over 400 s. b) The average rate of production of NO over 400 s. c) The average rate of production of O over 400 s.
  • Rate and Stoichiometry
  • Graphical Determination ofAverage Rate
  • Rate and Stoichiometry• Rate of creation of products or disappearance of reactants can be predicted from reaction stoichiometry.• From the example on the previous page, for the reaction: 2 NO2(g) → 2 NO(g) + O2(g)
  • Rate and Stoichiometry• From the graph and the calculations, it can be seen: – the rate of decomposition of NO2 is equal to the production of NO. The ratio of the molar coefficients is 1:1. This means as one molecule of NO2 is decomposed, one molecule of NO is created. Their rates should be equal. – the rate of production of oxygen is half that of the NO. The ratio of the molar coefficients is 2:1, as well. The rate of production of NO should be double that of the oxygen.• Therefore,
  • Rate and Stoichiometry• Example 1 – For the decomposition of nitrogen dioxide, if the rate of decomposition is determined to be 0.50 mol/Ls at a certain temperature, predict the rate of creation of both products. 2 NO2(g) → 2 NO(g) + O2(g)
  • Rate and Stoichiometry• Example 2 – For the reaction 2 A + B → 3 C, what is the rate of production of C and the rate of disappearance of B if A is used up at a rate of 0.60 mol/Ls?
  • Collision TheoryA Thought Experiment• A group of students are taken into the school hall and blindfolded. They are asked to move around the hall. If two students crash into each other and both students fall over, then they stay lying on the ground - they have "reacted".• How can we increase the rate at which students fall down? There are two obvious choices: – 1) Put more students into the hall. This will lead to more collisions between students. – 2) Ask the students in the hall to run more quickly. This will also lead to more collisions and to a greater chance of a collision leading to the students falling over.• We will use this "thought experiment" to try and explain rates of reaction.
  • Collision Theory• The collision theory attempts to explain known information concerning the rates of chemical reactions. It uses kinetic molecular theory, and other factors, such as molecular shape and structure, in its explanations.• The basic assumption of the collision theory is that for a reaction to occur the reacting species must collide. The collision theory applies to three factors effecting reaction rate (effect of concentration, effect of temperature, effect of surface area).
  • Collision Theory• The effect of concentration• The effect of temperature• The effect of particle size
  • Collision Theory• The rate of chemical reaction, specifically depends upon the following two factors: – Number of collisions that occur per unit of time (i.e. frequency of collisions). – The fraction of the collisions that is effective in causing a “reaction”.
  • Collision Theory
  • Collision Theory• Fast Reactions:F reactants that are in a liquid (aq) or gaseous (g) state have greater mobility of the molecules, thus a good chance of contacting each otherco when reactants are ionic, the positive- negative attraction works rapidly
  • Collision Theory• Slow Reactions:Sl when reactants are both solids, the particles have very little mobilitypa often slow reactants are neutral molecules, this means that electron transfer and bond arrangement takes some time to occurso complex structures take longer to rearrange than do small ones (LEGO analogy)
  • Collision Theory• The collision theory suggest an additional reason as to why it may take “a while” for certain species to react:• Collisions will only successful in producing a reaction if the molecules have a particular orientation, so that particular chemical bonds may be either broken or formed.
  • Activation Energy• Particles must collide with enough energy in order to cause a reaction.• Chemical reactions involve the making and breaking of bonds, which requires energy. If colliding particles do not have sufficient velocity (Kinetic Energy) they will not produce a reaction.• The minimum amount of energy required for colliding particles to produce a chemical reaction is called the Activation Energy (EA) of that reaction.
  • Collision Theory& Activation Energy
  • Activation Energy• The greater the activation energy, the slower the reaction rate, the longer the reaction takes.• The rate of a reaction is determined by the height of the barrier that the particles must cross in order that they are converted into products.• The activation energy is the barrier colliding particles must overcome.
  • Activation Energy• As the particles collide, they form an unstable intermediate particle called the activated complex or transition state. The energy required to produce the activated complex is the activation energy. Enthalpy (H) is the heat content or total energy possessed by the particles in a system.
  • Activation Energy• The energy released or absorbed by a reaction is called the change in enthalpy, ∆H, or heat of reaction. ∆H = HPRODUCTS - HREACTANTS• where ∆H is the change in enthalpy or energy that occurs during a reaction in Joules (J).• If ∆H is negative, heat flows out of the system. This type of reaction gives off heat and the reaction vessel feels warmer. This is called an exothermic reaction.• If ∆H is positive, heat is absorbed or flows into the system. The reaction vessel feels cooler as energy is absorbed from the surroundings. This type of reaction is called an endothermic reaction.
  • Reaction Coordinate Diagrams
  • Reaction Coordinate Diagrams
  • Reaction Coordinate Diagrams•The activated complex in this reaction is CH3CH2(OH)Br¯.•The activation energy is 88.9 kJ per mole of CH3CH2Br.•The enthalpy change is -77.2 kJ, indicating an exothermicreaction.•Consequently, this reaction does not take place unless 88.9 kJper mole of CH CH Br is added to the system.
  • Factors Affecting Reaction Rate• The rate of a reaction can be affected by a number of different variables. Listed below some factors that may affect the rate of reaction. – Nature of Reactants – Concentration of Reactants – Temperature – Catalysts – Surface Area – Pressure (Volume)
  • Factors AffectingReaction Rate• Nature of Reactants Some chemical reactions involve the rearrangement of atoms as a result of bond breaking and new bond formation. Other reactions are a result of electron transfer. The nature of the reactants involved in the reaction will affect the rate of the reaction. i) The fewer the number of bonds broken, the faster the rate. The equations give us an indication as to why: 2 NO(g) + O2(g) 2 NO2(g) 2 C8H18(l) + 25 O2(g) 16 CO2(g) + 18 H2O(g)
  • Factors AffectingReaction Rate• Nature of Reactants ii) In general, the weaker the bonds in the reactants, the faster the reaction. Reactions with molecular reactants (covalent bonds) are usually slower than ionic reactions. iii)The state or phase of the reactants will also affect the rate of a reaction. Reactions with gases as reactants are usually faster than those with liquids, which are both faster than those with solids. 
  • Factors AffectingReaction RateTemperature• According to the collision theory, reaction rate is determined by the frequency of collisions between molecules and increases as the frequency of collisions increases. According to the Kinetic Molecular Theory, the speed (Kinetic Energy) of particles increases as the temperature increases and the more likely the reactant particles will have enough energy to overcome the energy barrier to form products. A general rule is that for every 10 ºC rise in temperature the reaction rate doubles.
  • Factors AffectingReaction RateTemperature
  • Factors AffectingReaction RateConcentration of Reactants• Recall that concentration refers to the amount of reactant per unit volume. The units for concentration are mol/L.• Increasing the concentration of reactants increases the total number of particles in a container. If the number of particles in a container increases, so do the number of particles with activation energy. More particles with activation energy increases the frequency of effective collisions.
  • Factors Affecting Reaction RateConcentration of Reactants• Increasing the number of particles in a container will also increase the chances of a collision. The diagram below illustrates the effect of doubling the concentration of just one of the reactants: – In (a), with two particles of each, there exists four possible collisions which could produce a reaction. – In (b), by doubling the number of red particles, the number of possible collisions increases to eight! This will increase collision frequency.
  • Factors AffectingReaction RateCatalysts• A catalyst is a substance that speeds up or initiates a reaction without itself being permanently changed.• Catalysts reduce the activation energy needed for a reaction to occur and so increase the number of effective collisions. Catalysts speed up reaction rate but do not enter into the reaction and are therefore not changed or used up. This makes them very effective molecules for creating change. Our body has many enzymes (organic catalysts) which speed up reaction rates allowing the body to control its metabolism.• An inhibitor also affects reaction rates by stopping or slowing down reactions. It is the opposite of a catalyst.
  • Factors AffectingReaction RateCatalysts
  • Factors AffectingReaction RatePressure• Changing the pressure on a system usually only affects the reaction rates of gaseous reactions.• Pressure is the force of particles upon the walls of their container. If the number of particles in a container increases without changing volume, the pressure increases.• There are 3 ways to change pressure: – Add more product and/or reactant particles to the container. – Increase or decrease the volume of the container – Add an inert or unreactive gas.
  • Factors AffectingReaction RatePressure• If pressure is increased by adding more reactant particles, the concentration increases causing an increased rate. If the pressure is reduced by removing reactant, the rate decreases due to decreased concentration. If you decrease the volume of a container without changing the number of particles in the container, as in figure c in the diagram below, the concentration of the reactants increases. If you increase the volume of a container without changing the number of particles, as in figure a in the diagram above, the concentration of the reactants decreases.
  • Factors Affecting Reaction RateSurface Area• The size of the reactants, or surface area in contact, is usually only a factor in heterogeneous reactions. Increasing the surface area of the reactants by crushing, grinding or other means, increases the number of particles of reactants in contact. The rate of reactions increases when surface area in contact also increases.• Increasing surface area increases the frequency of collisions, increasing reaction rate.
  • Collision Theory &Reaction Mechanisms• The rate of a reaction cannot always be determined from the chemical equation of the reaction. Consider the reaction: 2 NO (g) + O2 (g) → 2 NO2 (g)• According to the collision theory, for this reaction to occur in one step, 3 particles must collide: two NO molecules and one O2. These particles must collide at exactly the same time with the correct orientation and enough energy.
  • Collision Theory &Reaction Mechanisms• Three particle collisions are quite rare. We can think of a three particle collision in terms of a pool shot. Striking the cue ball to create a collision which puts one ball in a pocket is difficult. Making a "combination" shot, striking two balls with enough energy and the correct angles to sink both, is much more difficult, and rarely attempted.
  • Collision Theory &Reaction Mechanisms• Chemical reactions tend to take place in steps, with each step involving a collision between two particles, or bimolecular. The chances of a two particle collision being successful are much better than a collision between three or more particles.
  • Collision Theory &Reaction MechanismsReaction Intermediates• Reactions which take place in one elementary step are called, appropriately, simple reactions.• Reactions which take place in more than one step are called complex reactions.• The complex reaction: 2 NO (g) + O2 (g) → 2 NO2 (g) does not take place in one elementary step, but actually takes place in two steps:
  • Collision Theory &Reaction MechanismsReaction Intermediates• The complex reaction: 2 NO (g) + O2 (g) → 2 NO2 (g)• does not take place in one elementary step, but actually takes place in two steps:
  • Collision Theory &Reaction MechanismsReaction Intermediates• The complex reaction: 2 NO (g) + O2 (g) → 2 NO2 (g)• Compounds such as N2O2, products of one reaction that immediately become reactants in another reaction, are called reaction intermediates. All complex reactions contain at least one reaction
  • Collision Theory &Reaction MechanismsNet Equations• The steps in which a reaction occurs is called that reactions mechanism. The sum of the steps of a mechanism must equal the total or net equation.• For the reaction, NO2 + CO → NO + CO2 the mechanism is as follows:
  • Collision Theory &Reaction MechanismsNet Equations• This mechanism agrees with the initial equation for this reaction, as the sum of the steps equals the original reaction. The NO3 is the reaction intermediate, so it does not appear in the net reaction. Since NO2 is found twice on the left and once on the right, we can cancel one of the NO2s just as we would adding equations in math.
  • Collision Theory &Reaction MechanismsRate Determining Step• Not all steps in a mechanism have the same rate. The step with the slowest rate is called the rate determining step (RDS), since that step affects the rate of the reaction more than the others.• The slowest step in the reaction mechanism determines the rate, and therefore only reactants in the slowest step affect rate. Increasing the concentrations of these reactants will speeding up the slow step will increase the overall reaction rate. Increasing the concentration of reactant in fast steps will have little effect on rate. To understand how this works consider the following analogy.
  • Collision Theory & Reaction MechanismsRate Determining Step• A student has to get to school in the morning, after having a good nights sleep in his own bed. The student might perform the following tasks (steps) in order to get to school (overall reaction)?• Sleeping in bed (reactant) --------> Arriving at school (product) – wake up – uses the washroom to clean up and have a shower. – have breakfast – walk, ride with family or friends or take the bus to school• The slowest step in this sequence will determine how fast the student can get to school (product). Waking up may be the slow step for some while taking a shower may be the slow step for others.• Speeding up the slow step will speed up the overall reaction. If for example taking a shower was the slow step the student might shower the night before and simply wash their face in the morning. If taking the bus is the slow step they might get a ride with a friend that might be quicker and therefore speed up the mechanism.
  • Collision Theory & Reaction MechanismsRate Determining Step• Assignment :Analogies of Reaction Rates• List 4 jobs that you do either at home or at work, that involve more that one task or step to accomplish.(For example cleaning your room, or cooking chicken at a local restaurant)• List what might be considered the starting point (reactant).• List what might be considered the finished product.• List the steps required and the amount of time needed to complete each step.• Describe which step is rate determinining.• Describe how this step might be sped up.
  • Collision Theory &Reaction MechanismsRate Determining Step• According to our last mechanism: step 1 is the RDS since it is the slowest step.
  • Collision Theory &Reaction MechanismsEnergy Changes• According to our last mechanism:• if the reaction is exothermic, we can predict its shape:
  • Collision Theory & Reaction MechanismsEnergy Changes• Without knowing the actual energy values, we can predict the shape of the curve: – Recall that the larger the activation energy, the slower the reaction. Step 1 is the rate determining step, so it should have the largest activation energy. – Step 2 is fast so it should have the lowest activation energy. Notice the activation energy for Step 2 begins where Step 1 ends. – The enthalpy of reactants should be higher than the enthalpy of products
  • Collision Theory & Reaction MechanismsMechanisms Problem• Since the RDS affects the rate of the entire reaction the most, changes to the reactants in the other steps will have very little effect on the rate of the reaction.• Example 1• Given the following mechanism: P+Q → X+T (slow) X+P → Y+R (fast) Y+S → T (moderate) a) What is the net reaction? b) What are the reaction intermediates? c) Which is the rate determining step? d) What would be the effect of increasing the concentration of P? e) What would be the effect of decreasing the concentration of Q? f) What would be the effect of increasing the concentration of S?
  • Rate Law• Rate Law is an expression which relates the rate of a reaction to the concentration of the reactants. Rate law is a tool which helps us calculate the rate of a reaction with given concentrations of reactants.• The rate law shows the quantitative effect of concentration on reaction rate.• For the reaction: A + B → products• The rate of consumption of A and B is directly proportional to their concentration. That is, the greater the concentration of A and B, the faster the rate.
  • Rate LawThe Molarity Relationship• M = n, V• the molar concentration α number of moles of substance – therefore R α [A] and R α [B] – AND: R α [A]x [B]y AND: R = k [A]x [B]y• R = rate of reaction (Moles/1/s or M/s)• [A] = molar concentration of A• [B] = molar concentration of B• k = specific rate constant• x, y = is the power, called the order of the reaction
  • Rate LawThe Molarity Relationship• R = k [A]x [B]y• The constant k, is known as the specific rate constant for the reaction. The rate constant and the order can only be determined experimentally.• The rate constant is specific for each reaction at a specific temperature, since it’s value depends upon the size, speed and types of molecules in the reaction.• Changing temperature would change the speed of the reactant particles and hence change the rate constant.• Temperature is the only factor, which affects the rate constant.
  • Rate LawThe Law of Mass Action• The rate of a reaction is proportional to the product of the concentrations of the reactants, each raised to a power equal to the coefficient in the original equation.
  • Rate LawReaction Order• The order of a reaction indicates how concentration of reactants affects the rate of a reaction.• Example, A → products,• If the order of a reaction is a first order reaction, x = 1, this means the reaction rate was directly proportional to changes in reactant concentration. If the concentration of A were doubled, the rate would double. If the concentration were tripled, the rate would triple, etc.
  • Rate LawReaction Order• If the reaction is a second order reaction, x = 2, doubling the concentration would increase the rate by a factor of 2x = 22 = 4. That is, the rate would increase four times. Tripling the concentration of A would cause the rate to increase nine times (3x=32=9).• If the rate of the reaction did not depend on the concentration of A, it would be a zero order reaction, x = 0. This means a change in the concentration of A does NOT change the rate of the reaction.
  • Rate LawReaction Order• For a reaction with more than one reactant, such as A + B → products• The rate law would be: Rate = k[A]x[B]y• The rate depends on both A and B concentrations. Each reactant can affect the rate differently.• The total order of the reaction is the sum of the order with respect to A and the order with respect to B, x + y.
  • Rate LawCalculating Rate Law• The rate law can only be determined experimentally. Even though the rate of formation of the products and the rate of consumption of reactants is related to their reaction stoichiometry, the rate law cannot usually be determined from the molar coefficients.• One method of determining rate law is by measuring the effect of changes in concentration of one reactant on the initial reaction rate, while keeping the other reactant concentrations constant.
  • Rate LawCalculating Rate Law• Example 1:• What is the rate law for the following reaction, given the experimental data below? H2O2 + 2 HI → 2 H2O + I2
  • Rate LawCalculating Rate Law• Example 2:• For the reaction 3 A(g) + B(g) + 2 C(g) → 2 D(g) + 3 E(g)a. Write the rate law for this reaction.b. Calculate the value of the rate constant.c. Calculate the rate for Trial #5.d. Calculate the concentrationof A in Trial #6.
  • Rate LawRate Law & Stoichiometry• For most reactions, the rate law, the specific rate constant, k, and the mechanism of a reaction can only be determined experimentally, not from the reaction stoichiometry.• However, for reactions that occur in a single step (elementary reactions) the order of each reactant in the rate law is equal to the coefficient in the reactions balanced equation.• For the elementary reaction: aA + bB → cC + dD• the rate law is rate = k[A]a[B]b• where a and b are the molar coefficients for the elementary reaction.
  • Rate LawRate Law & Reaction Mechanism• We have seen that rate law for reactions which occur in more than one step do NOT correspond to the coefficients in its balanced net equation.• The rate law for these equations corresponds to the stoichiometry of the rate determining step. Those reactants which are zero order, do not appear in the RDS and, hence do not affect reaction rate significantly.
  • Rate LawRate Law & Reaction Mechanism• Example• The mechanism for the reaction 3M + N → P + 2Q• is below: 2M → P + X (slow) X+M→ Q + Y (moderate) N+Y → Q (fast) a) What is the rate law for this reaction? b) What would be the effect of tripling the [M]? c) What would be the effect of doubling the [N]?