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UNIT 3: Energy Changes and Rates of Reaction Chapter 5: Energy Changes Chapter 6: Rates of Reaction.

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Presentation on theme: "UNIT 3: Energy Changes and Rates of Reaction Chapter 5: Energy Changes Chapter 6: Rates of Reaction."— Presentation transcript:

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2 UNIT 3: Energy Changes and Rates of Reaction Chapter 5: Energy Changes Chapter 6: Rates of Reaction

3 UNIT 3 An important part of studying chemical reactions is to monitor the speed at which they occur. Chemists look at how quickly, or slowly, reactions take place and how these rates of reaction are affected by different factors. TO PREVIOUS SLIDEPREVIOUS The light produced by a firefly depends on the speed of a particular chemical reaction that occurs in its abdomen. Chapter 6: Rates of Reaction

4 UNIT 3 Section Chemical Reaction Rates TO PREVIOUS SLIDEPREVIOUS Chemical kinetics is the study of the rate at which chemical reactions occur. Chapter 6: Rates of Reaction The term reaction rate, or rate of reaction refers to: the speed that a chemical reaction occurs at, or the change in amount of reactants consumed or products formed over a specific time interval

5 UNIT 3 Section 6.1 Determining Reaction Rates TO PREVIOUS SLIDEPREVIOUS The reaction rate is often given in terms of the change in concentration of a reactant or product per unit of time. The change in concentration of reactant A was monitored over time. Chapter 6: Rates of Reaction

6 UNIT 3 Section 6.1 Determining Reaction Rates TO PREVIOUS SLIDEPREVIOUS The change in concentration of reactant or product over time is often graphed. For the reaction A → B, over time, the concentration of A decreases, and the concentration of B increases. Chapter 6: Rates of Reaction

7 UNIT 3 Section 6.1 Average and Instantaneous Reaction Rates TO PREVIOUS SLIDEPREVIOUS Average rate of reaction: change in [reactant] or [product] over a given time period (slope between two points) Instantaneous rate of reaction: the rate of a reaction at a particular point in time (slope of the tangent line) Average rate of reaction and instantaneous rate of reaction can be determined from a graph of concentration vs. time. Chapter 6: Rates of Reaction

8 UNIT 3 Section 6.1 Expressing Reaction Rates in Terms of Reactants or Products TO PREVIOUS SLIDEPREVIOUS A known change in concentration of one reactant or product and coefficients of a chemical equation allows determination of changes in concentration of other reactants or products. Chapter 6: Rates of Reaction

9 Express the rate of formation of ammonia relative to hydrazine, for the reaction on the previous slide. UNIT 3 Section 6.1 Answer on the next slide TO PREVIOUS SLIDEPREVIOUS Chapter 6: Rates of Reaction L EARNING C HECK

10 The mole ratio of ammonia to hydrazine is 4:3 Section 6.1 UNIT 3 TO PREVIOUS SLIDEPREVIOUS Chapter 6: Rates of Reaction L EARNING C HECK

11 UNIT 3 Section 6.1 Methods for Measuring Rates of Reaction TO PREVIOUS SLIDEPREVIOUS Chapter 6: Rates of Reaction

12 UNIT 3 Section 6.1 Calculating Reaction Rates from Experimental Data TO PREVIOUS SLIDEPREVIOUS Chapter 6: Rates of Reaction The following data were collected in order to calculate the rate of a reaction. Calculations on the next two slides show how to use volume data and mass data to determine the average rate of a reaction.

13 UNIT 3 Section 6.1 TO PREVIOUS SLIDEPREVIOUS Chapter 6: Rates of Reaction Calculating Reaction Rates from Experimental Data

14 UNIT 3 Section 6.1 Chapter 6: Rates of Reaction Calculating Reaction Rates from Experimental Data TO PREVIOUS SLIDEPREVIOUS

15 Section 6.1 Review UNIT 3 Section 6.1 TO PREVIOUS SLIDEPREVIOUS Chapter 6: Rates of Reaction

16 UNIT 3 Section Collision Theory and Factors Affecting Rates of Reaction TO PREVIOUS SLIDEPREVIOUS According to collision theory, a chemical reaction occurs when the reacting particles collide with one another. Only a fraction of collisions between particles result in a chemical reaction because certain criteria must be met. Chapter 6: Rates of Reaction

17 UNIT 3 Section 6.2 Effective Collision Criteria 1: The Correct Orientation of Reactants TO PREVIOUS SLIDEPREVIOUS For a chemical reaction to occur, reactant molecules must collide with the correct orientation relative to each other (collision geometry). Five of many possible ways that NO(g) can collide with NO 3 (g) are shown. Only one has the correct collision geometry for reaction to occur. Chapter 6: Rates of Reaction

18 UNIT 3 Section 6.2 Effective Collision Criteria 2: Sufficient Activation Energy TO PREVIOUS SLIDEPREVIOUS The shaded part of the Maxwell- Boltzmann distribution curve represents the fraction of particles that have enough collision energy for a reaction (ie the energy is ≥ E a ). Chapter 6: Rates of Reaction For a chemical reaction, reactant molecules must also collide with sufficient energy. Activation energy, E a, is the minimum amount of collision energy required to initiate a chemical reaction. Collision energy depends on the kinetic energy of the colliding particles.

19 UNIT 3 Section 6.2 Representing the Progress of a Chemical Reaction TO PREVIOUS SLIDEPREVIOUS Chapter 6: Rates of Reaction From left to right on a potential energy curve for a reaction: potential energy increases as reactants become closer when collision energy is ≥ maximum potential energy, reactants will transform to a transition state products then form (or reactants re-form if ineffective) ExothermicEndothermic

20 UNIT 3 Section 6.2 Activation Energy and Enthalpy TO PREVIOUS SLIDEPREVIOUS The E a for a reaction cannot be predicted from ∆H. ∆H is determined only by the difference in potential energy between reactants and products. E a is determined by analyzing rates of reaction at differing temperatures. Reactions with low E a occur quickly. Reactions with high E a occur slowly. Potential energy diagram for the combustion of octane. Chapter 6: Rates of Reaction

21 UNIT 3 Section 6.2 Activation Energy for Reversible Reactions TO PREVIOUS SLIDEPREVIOUS Potential energy diagrams can represent both forward and reverse reactions. follow left to right for the forward reaction follow right to left for the reverse reaction Chapter 6: Rates of Reaction

22 UNIT 3 Section 6.2 Analyzing Reactions Using Potential Energy Diagrams TO PREVIOUS SLIDEPREVIOUS The BrCH 3 molecule and OH - collide with the correct orientation and sufficient energy and an activated complex forms. When chemical bonds reform, potential energy decreases and kinetic energy increases as the particles move apart. Chapter 6: Rates of Reaction

23 Describe the relative values of E a(fwd) and E a(rev) for an exothermic reaction UNIT 3 Section 6.2 Answer on the next slide TO PREVIOUS SLIDEPREVIOUS Chapter 6: Rates of Reaction L EARNING C HECK

24 E a(rev) is greater than E a(fwd) Section 6.2 UNIT 3 TO PREVIOUS SLIDEPREVIOUS Chapter 6: Rates of Reaction L EARNING C HECK

25 UNIT 3 Section 6.2 Factors Affecting Reaction Rate TO PREVIOUS SLIDEPREVIOUS 1. Nature of reactants reactions of ions tend to be faster than those of molecules 2. Concentration a greater number of effective collisions are more likely with a higher concentration of reactant particles 3. Temperature with an increase in temperature, there are more particles with sufficient energy needed for a reaction (energy is ≥ E a ) Chapter 6: Rates of Reaction

26 UNIT 3 Section 6.2 Factors Affecting Reaction Rate TO PREVIOUS SLIDEPREVIOUS Chapter 6: Rates of Reaction 4. Pressure for gaseous reactants, the number of collisions in a certain time interval increases with increased pressure 5. Surface area a greater exposed surface area of solid reactant means a greater chance of effective collisions 6. Presence of a catalyst a catalyst is a substance that increases a reaction rate without being consumed by the reaction

27 UNIT 3 Section 6.2 A Catalyst Influences the Reaction Rate TO PREVIOUS SLIDEPREVIOUS A catalyst lowers the E a of a reaction. this increases the fraction of reactants that have enough kinetic energy to overcome the activation energy barrier a catalyzed reaction has the same reactants, products, and enthalpy change as the uncatalyzed reaction A catalyst decreases both E a(fwd) and E a(rev). Chapter 6: Rates of Reaction

28 UNIT 3 Section 6.2 Catalysts in Industry TO PREVIOUS SLIDEPREVIOUS A metal catalyst is used for industrial-scale production of ammonia from nitrogen and hydrogen. Hydrogen and nitrogen molecules break apart when in contact with the catalyst. These highly reactive species then recombine to form ammonia. Chapter 6: Rates of Reaction A catalyst (V 2 O 5 ) is used for industrial-scale production of sulfuric acid from sulfur.

29 UNIT 3 Section 6.2 Catalysts in Industry TO PREVIOUS SLIDEPREVIOUS The Ostwald process uses a platinum-rhodium catalyst for the industrial production of nitric acid. Chapter 6: Rates of Reaction Many industries use biological catalysts, called enzymes, which are most often proteins. For example: the use of enzymes decreases the amount of bleach (an environmental hazard) needed to whiten fibres used in paper production.

30 Section 6.2 Review UNIT 3 Section 6.2 TO PREVIOUS SLIDEPREVIOUS Chapter 6: Rates of Reaction

31 UNIT 3 Section Reaction Rates and Reaction Mechanisms TO PREVIOUS SLIDEPREVIOUS Initial rate is found by determining the slope of a line tangent to the curve at time zero. Chapter 6: Rates of Reaction Initial rate is the rate of a chemical reaction at time zero. products of the reaction are not present, so the reverse reaction cannot occur it is a more accurate method for studying the relationship between concentration of reactant and reaction rate

32 UNIT 3 Section 6.3 Graphing Reaction Rate in Terms of Concentration TO PREVIOUS SLIDEPREVIOUS To study the effects of concentration on reaction rate: different starting concentrations of reactant are used initial rates are calculated using the slopes of the tangent lines from concentration vs time curves initial rates are plotted against starting concentration Chapter 6: Rates of Reaction Initial rates are determined (A) and these are plotted against concentration (B).

33 UNIT 3 Section 6.3 First-order Reactions TO PREVIOUS SLIDEPREVIOUS The initial rate vs starting concentration graph on the previous slide is a straight line. Chapter 6: Rates of Reaction the equation of the line can be expressed as: rate = k[A] This represents a first-order reaction For reactions with more than one reactant (e.g. A and B): if experiments for each reactant produce straight lines, the rate is “first order with respect to reactant A and first order with respect to reactant B.”

34 UNIT 3 Section 6.3 Second-order Reactions TO PREVIOUS SLIDEPREVIOUS For chlorine dioxide in this reaction: the initial rate vs concentration curve is parabolic the reaction is proportional to the square of [ClO 2 ] it is a second order reaction with respect to this reactant Chapter 6: Rates of Reaction rate = k[A] 2

35 UNIT 3 Section 6.3 The Rate Law TO PREVIOUS SLIDEPREVIOUS The rate law shows the relationship between reaction rates and concentration of reactants for the overall reaction. Chapter 6: Rates of Reaction rate = k[A] m [B] n m: order of the reaction for reactant A n: order of the reaction for reactant B k: rate constant m + n: order of the overall reaction

36 UNIT 3 Section 6.3 Reaction Mechanisms TO PREVIOUS SLIDEPREVIOUS A reaction mechanism is the series of elementary steps that occur as reactants are converted to products. Chapter 6: Rates of Reaction For example, oxygen and nitrogen are not formed directly from the decomposition of nitrogen dioxide: It occurs in two elementary steps:

37 UNIT 3 Section 6.3 The Rate-determining Step TO PREVIOUS SLIDEPREVIOUS This reaction occurs in three elementary steps: Chapter 6: Rates of Reaction Step 2 is the rate-determining step: it is the slowest elementary step the overall rate of the reaction is dependent on this step the E a for this step is higher than E a for each of the other steps

38 UNIT 3 Section 6.3 A Proposed Reaction Mechanism TO PREVIOUS SLIDEPREVIOUS Experiments show that this reaction is zero order with respect to OH – (i.e. its rate does not depend on [OH – ]) This can be explained by a two-step mechanism Step 2 is very fast and depends on completion of Step 1, not on the concentration of OH –. Chapter 6: Rates of Reaction

39 Section 6.3 Review UNIT 3 Section 6.3 TO PREVIOUS SLIDEPREVIOUS Chapter 6: Rates of Reaction


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