Chemical Kinetics Unit 12 Part 2

Kinetics –Kinetics – the study of the speeds of chemical reactions and the mechanisms by which reactions occur. Reaction RatesAlso referred to as “Reaction Rates” Rates of chemical change usually are expressed as the amount of reactant forming products per unit time. ∆[reactants] ∆ time Kinetics

How do reactions really occur?How do reactions really occur? –Reactant particles MUST collide! Rates of chemical reactions are described in a model called collision theory.Rates of chemical reactions are described in a model called collision theory. –Studies show most molecular collisions do NOT result in a reaction…why?? Atoms, ions, and molecules react to form products only when they collide with the proper orientation and sufficient energy.Atoms, ions, and molecules react to form products only when they collide with the proper orientation and sufficient energy. Collision Theory

Effective collisions are defined by two conditions:Effective collisions are defined by two conditions: –exactly the right orientation to react (head on collision is best) –enough energy to break bonds and form new ones The minimum amount of energy required is called the reaction’s activation energy.The minimum amount of energy required is called the reaction’s activation energy. –The activation energy is a barrier that reactants must get over to react –The higher the barrier the larger the amount of energy needed for the reaction to proceed Collision Theory

During a reaction, a particle that is neither reactant nor product forms momentarily, called an activated complexDuring a reaction, a particle that is neither reactant nor product forms momentarily, called an activated complex –if there is sufficient energy –and if the atoms are oriented properly An activated complex is a kind of transition molecule which has similarities to reactants & productsAn activated complex is a kind of transition molecule which has similarities to reactants & products –An activated complex is the arrangement of atoms at the peak of the activation-energy barrier. Collision Theory

Collision theory explains why some naturally occurring reactions are immeasurably slow at room temp.Collision theory explains why some naturally occurring reactions are immeasurably slow at room temp. –Carbon and Oxygen react when charcoal burns, but this reaction has a high activation energy –At room temp, the collisions of oxygen and carbon molecules aren’t energetic enough to react –But the reaction can be helped along a number of ways Collision Theory

It is possible to vary the conditions of the reaction, the rate of almost any reaction can be modifiedIt is possible to vary the conditions of the reaction, the rate of almost any reaction can be modified ocollision theory can help explain why the rates can be modified Several strategies can be used to speed up reactions:Several strategies can be used to speed up reactions: oIncrease the temperature oIncrease the concentration oDecrease the particle size oEmploy a catalyst Reaction Rates

Increasing the temperature speeds up the reaction, while lowering the temperature slows down the reactionIncreasing the temperature speeds up the reaction, while lowering the temperature slows down the reaction For every 10 o C increase in temperature, the reaction rate usually DOUBLES!For every 10 o C increase in temperature, the reaction rate usually DOUBLES! Recall, temperature is directly proportional to average kinetic energy:Recall, temperature is directly proportional to average kinetic energy: –Average KE = ½ x mass x velocity 2 At a higher temperature, there are more effective collisions – more molecules have the velocity / energy needed to reactAt a higher temperature, there are more effective collisions – more molecules have the velocity / energy needed to react Temperature

Just sitting out, charcoal does not react at a measurable rateJust sitting out, charcoal does not react at a measurable rate –However, when a starter flame touches the charcoal, the temperature is increased so atoms of reactants collide with higher energy and greater frequency –Some of the collisions are high enough in energy that the product CO 2 is formed o The energy released by the exothermic reaction then supplies enough energy to get more C and O 2 over the activation- energy barrier Remove the starter flame: the reaction will continue on its own. Remove the starter flame: the reaction will continue on its own.

The more reacting particles you have in a given volume, the higher the rate of reaction.The more reacting particles you have in a given volume, the higher the rate of reaction. Cramming more particles into a fixed volume increases the concentration of reactants:Cramming more particles into a fixed volume increases the concentration of reactants: Increasing the concentration, increases the frequency of the collisions, and therefore increasing the reaction rate.Increasing the concentration, increases the frequency of the collisions, and therefore increasing the reaction rate. ***For gases, increasing pressure and/or decreasing volume increases concentration.***For gases, increasing pressure and/or decreasing volume increases concentration. Concentration

The smaller the particle size, the larger the surface area for a given mass of particles. Effectively is increasing the concentration.The smaller the particle size, the larger the surface area for a given mass of particles. Effectively is increasing the concentration. An increase in surface area increases the amount of the reactant exposed for collision to take place…An increase in surface area increases the amount of the reactant exposed for collision to take place… –Which increases the collision frequency and the reaction rate. Methods:Methods: –Grinding the reactants into a powder –Dissolving in a solvent. Particle Size

A catalyst is often the best way to speed up an reaction. In fact, some reactions simply will not go forward measurably without one. A catalyst is a substance that increases the rate of a reaction without being changed or used up during the reaction The key is that they permit reactions to proceed at lower activation energy than is normally required Catalyst

Catalysts lower the required activation energy by providing an alternative path or arrangement for the molecules to react. Not always understood why certain catalyst behave the way the do…they just do! By lowering the Ea threshold, more molecules in the system will have the required energy to surmount the barrier. A negative catalyst has the opposite effect. It interferes with effective collisions and increases the Ea required. Catalyst

The type of reactant substances also plays a role in reaction rates. Weak bonds = easier to break, reactants with weak bonds will react faster. –Essentially results in a low activation energy. Strong bonds = harder to break, reactants with strong bonds will react slower (high Ea) Electronegativity and ionization energy also plays a role. Nature of Reactants

Reaction Rate Laws relate the speed of a reaction to the concentrations of the reactants.Reaction Rate Laws relate the speed of a reaction to the concentrations of the reactants. –Determined from experiment The General Rate LawThe General Rate Law –For reaction: xA + yB  Products –Rate = k[A] m [B] n Example: N 2 + 2H 2  2NH3Example: N 2 + 2H 2  2NH3 –Rate Law: Rate = k[N 2 ] 1 [H 2 ] 2 Reaction Rate Laws

Reaction Mechanism Most chemical reactions consist of complex multi-step sequences of two or more simpler (elementary) reactions. For example: Ozone in the upper atmosphere can decompose into Oxygen: 2O 3  3O 2 Happens after UV light liberates Cl atoms from chlorofluorocarbons and breaks down ozone:Happens after UV light liberates Cl atoms from chlorofluorocarbons and breaks down ozone: Elementary Step 1: Cl + O 3  ClO + O 2 Elementary Step 2: O 3  O + O 2 Elementary Step 3: ClO + O  Cl + O 2 Overall Net Reaction: Overall Net Reaction: 2O 3  3O 2 UV

Reaction Mechanism Since ClO molecules and O atoms were formed in one elementary step and then used up in another, they are called “intermediates”. Like catalysts, they do not appear in the net chemical equation. Intermediates live much longer than activated complexes Many intermediates exist only for a fraction of a second: often measured in femtoseconds (1 x seconds). Caltech scientists have used laser flashes to record the formation of intermediates.

For a complex reaction, its Energy Diagram has a series of hills & valleysFor a complex reaction, its Energy Diagram has a series of hills & valleys –The peaks correspond to the energies of the activated complexes –Each valley represents an intermediate product which becomes a reactant in the next stage of the reaction For the complex reaction:For the complex reaction: The elementary steps are:The elementary steps are: Reaction Mechanism H 2 (g) + 2ICl (g)  I 2 (g) + 2HCl(g) 1) H 2 + 2ICl(g)  ICl(g) + HCl(g) + HI(g) 2) ICl(g) + HCl(g) + HI(g)  I 2 (g) +2HCl(g)

Energy Diagram with Intermediates

Reaction Mechanism Rate Determining Step – every reaction has one step that is the slowest. A reaction can not go any faster than the slowest step, so it “determines” the rate. Complex reaction: 2NO + 2H 2  N 2 + 2H 2 O Step 1: 2NO  N 2 O 2 Step 2: N 2 O 2 + H 2  N 2 O + H 2 O Step 3: N 2 O + H 2  N 2 + H 2 O Thus, the second step is the rate determining step. Fast! Slow

Reaction Mechanism What would be the rate determining step for the reaction:What would be the rate determining step for the reaction: 4A + 2B + 2C  2D + 3E Step 1: 2A + B  F + D FAST Step 2: 2C + B  D + 2F SLOW Step 3: 2A + 3F  3E FAST Which is the rate determining step? (Step 2) Which reactant concentrations have an effect on rate? (C and B only)