Gases have important properties that distinguish them from solids and liquids:

Transcription

1 Kinetic molecular theory Gases have important properties that distinguish them from solids and liquids: Gases diffuse to occupy available space. For example, the molecules responsible for the scent of an apple diffuse throughout a home when an apple pie is baked in an oven. Gases exert pressure. Propane gas is contained in rigid steel cylinders because the force exerted by a compressed gas is large enough to rupture thin walled containers. Another example is the expansion of a rubber balloon when it is filled with air. The pressure of the gas forces the wall of the balloon to expand. When you push on the walls of a balloon, the gas inside "pushes" back. Gas molecules have little or no volume compared to the space they occupy. Furthermore, they have little or no attractions for other gas molecules. Gas molecules move in straight lines from one location to another at very high velocities until they collide with something and change course. The collisions are perfectly elastic, that is, the kinetic energy (energy of motion) possessed by the colliding particles is the same before and after the collision. The collision theory states that reacting particles must collide with one another in order for a reaction to occur. However, in the same way that not all personal encounters result in friendships, not all collisions between particles result in chemical changes. The theory of chemical reactions places conditions on the nature of the particle to particle collisions that cause old bonds to break and/or new bonds to form. The particles must collide with: correct orientation (collide so that bonds can be broken). and sufficient intensity (meets the activation energy).

2 Chemical Reactions and Potential Energy Chemical reactions involve the the breaking and forming of chemical bonds. Bond breaking is an endothermic process or one which requires the input of energy. Bond forming is an exothermic process or one that results in release of energy. For any chemical reaction, there is always a net change in chemical potential energy - it is either lost or gained as a result of bond breaking and bond forming. The potential energy changes that occur as reactants become products during a chemical change can be illustrated in a potential energy diagram (also known as a energy profile diagram). Drawing and interpreting these diagrams is the focus of this lesson; however, before you begin to work with these diagrams, it would be helpful to review some key concepts. You will start with a review of kinetic energy and temperature. Temperature and Kinetic Energy Kinetic energy is energy of motion. It is a function of the the mass and velocity (speed) of an object like a molecule, a piece of lint, or a truck. In chemistry, we are interested in the motion of atoms, ions and molecules and the relationship between kinetic energy and potential energy. Temperature is an indicator of kinetic energy. For example, the air temperature in your room is probably close to 20 C right now. This means that molecules in the air are moving faster than they do at 0 C, but slower than at 40 C. Moving particles possess kinetic energy. As two molecules approach each other, gravitational forces accelerate them towards each other (see the little bumps on the kinetic energy graph), but just before the particles collide, repulsive forces begin to slow the molecules down (see the steep declines on the kinetic energy graph). At the instant of collision kinetic energy is zero - the molecules come to a complete momentary stop. The potential energy of a particle is a function of its position relative to another particle. As two molecules approach, weak attractions cause a decrease in potential energy (see the little declines in the potential energy graph), but as the molecules collide the bonds are stretched to the breaking point (see the sharp increases in potential energy). The harder two molecules collide, the higher the spike on potential energy graph. If collisions are hard enough, the potential energy gain may be substantial enough to cause old bonds to break and new bonds to form. In many circumstances, only some of the particles in a sample have enough kinetic energy to collide hard enough to cause bonds to stretch to the breaking point. This critical kinetic energy level is called the threshold energy.

3 The high energy molecules are the ones that possess enough kinetic energy to result in the breaking of old bonds when they collide. The other molecules will just bounce off each other when they collide. Energy Profiles - Potential Energy Diagrams The kinetic energy of colliding particles is converted to potential energy. A potential energy diagram shows the potential energy changes that occur as reactants become products. It has five distinct regions: 1. the potential energy of the reactants 2. the potential energy gain that must take place in order for old bonds to be stretched to the breaking point 3. the potential energy of the transition state 4. the potential energy released as new bonds form during a chemical change 5. the potential energy of the products. Details: 1. The flat region labelled "Reactants" shows the potential energy of the reacting particles relative to the products. The actual potential energy of the reactants is an unknown. 2. Moving particles possess kinetic energy. When they collide, their kinetic energy is converted to potential energy. The rising part of the graph represents the increase in

4 potential energy that occurs when reactants collide. The minimum gain in potential energy that results in the stretching of reactant bonds to the breaking point is called the activation energy (E a ). It can be determined by experiment. 3. The top of the curve represents the point at which the bonds of the colliding particles are stretched to the breaking point. The unstable group of atoms formed at this point are neither reactants nor products but something in between - a transitional structure called the activated complex. The potential energy of this structure is very high because the bonds are stretched as far as possible. This structure exists for the shortest amount of time imaginable. In an instant, the particles either form new bonds to give new products or reform old bonds to give the original reactants. 4. The falling part of the curve represents the energy released when new bonds form between particles to make one or more products. The potential energy difference between the reactants and the products is called the heat of reaction (DH). It represents the net energy change of the reaction. If the potential energy of the products is greater than that of the reactants, then the reaction is classified as endothermic. If the potential energy of the products is less than that of the reactants, then the reaction is classified as exothermic. 5. The second flat region represents the potential energy of the products. The actual potential energy of the the products is also an unknown. The rate of a chemical reaction is a function of the number of successful collisions per unit of time where successful means that particles have collided with the correct orientation and sufficient intensity to form an activated complex. The potential energy gain resulting from the collision is the activation energy. It corresponds to the threshold kinetic energy of the colliding particles. An activated complex is a species formed at the transition point of a collision. The species is highly unstable - it can become reactants again or turn into products. Potential energy values for reactants, activated complexes and products are relative. We can never know the actual potential energy values for chemical species, but we can determine the energy changes they undergo during physical and chemical changes. Heat of reaction (DH) is negative for exothermic changes and positive for endothermic changes. Potential energy is measured in Joules (J) or kilojoules (kj). The formation of water is exothermic, 286 kj of heat is released per mole of water formed. By convention, the sign of this value is negative, -286 kj. The decomposition of water is endothermic, 286 kj of heat is absorbed per mole of water broken down into hydrogen and oxygen gas. By convention, the sign of this value is positive, +286 kj. If the sign of the forward reaction is negative, then the sign of the reverse reaction is positive (and vice versa). Given two of the three labelled values in the diagram above, you can calculate the third.

5 Potential Energy Calculations Given a potential energy diagram or information about the potential energy of species in a chemical reaction, you may be required to perform simple calculations to determine a missing quantity. Heat of reaction (DH) is the potential energy of the products minus the potential energy of the reactants. Heat of reaction (DH) is also activation energy of forward reaction (E a forward ) minus activation energy of reverse reaction (E a reverse ). Activation energy of forward reaction (E a forward ) is activation energy of reverse reaction (E areverse ) plus heat of reaction of forward reaction (?H forward ). Activation energy of reverse reaction is activation energy of forward reaction (E a forward ) plus heat of reaction of reverse reaction (?H reverse ). Note: ALL activation energy values are positive! Heat of reaction values may be positive or negative. The factors that affect reaction rates are: surface area of a solid reactant (larger surface area faster reaction) concentration or pressure of a reactant (increase = faster reaction) temperature (increase= faster reaction) nature of the reactants presence/absence of a catalyst (present = faster reaction) A change in one or more of these factors may alter the rate of a reaction. In this lesson, you will define these factors, and describe and predict their effects on reaction rates. Chemical changes involve bond breaking and bond forming. Bonds tend to break as a result of collisions between particles. New bonds form to make products. When you look at a chemical equation, you are looking at a summary of all the processes that have occurred in a reaction. For example, indicates that four C-H bonds and two O=O bonds have been broken followed by the forming of two C=O bonds and four H-O bonds. The likelihood that all of these bond breaking and forming events occur at exactly the same time are just about nil. Chemists try to determine the individual events and the order in which they occur. They attempt to determine which events have the highest activation energy. By knowing this, they may come up with alternate reaction events that have lower activation energies thereby making the overall reaction faster. In many cases this is achieved using a catalyst.

6 What is a Reaction Mechanism? A reaction mechanism is a detailed description of the pathway followed as reactants become products during a chemical change. Here's an example. Dinitrogen monoxide decomposes into nitrogen and oxygen gas: On the surface it might appear that a single collision between two N 2 O molecules results in the formation of three product molecules. However, laboratory analysis of this reaction reveals that it actually takes place in two stages: First, a collision between two N 2 O molecules causes one N-O bond to break: Then a collision between the O atom and another N 2 O molecule produces O 2 and N 2 molecules: These two equations can be added together to give the overall equation: In this example, the first two equations represent elementary processes - single events in which the bonding in one or more particles is changed to produce new bonding arrangements. Elementary processes are often called steps. The O atom is an example of a reaction intermediate, or a species that is generated in one step and consumed in a subsequent step. It is highlighted in red above. Do not confuse a reaction intermediate with an activated complex. An activated complex exists only at the point of collision - it is not detectable. A reaction intermediate exists after bonds are rearranged - its presence in a chemical system can be detected. In fact, detection of intermediates is crucial to the acceptance of a proposed reaction mechanism. Detection could involve measuring a short-lived ph change or observing a specific colour change. Rate Determining Step

7 The overall rate of a reaction is a function of the rate of the slowest elementary process (step). The slow step of a reaction is conveniently known as the rate determining step. In the example above, the first elementary process is the rate determining step. Analogy. When you measure the rate of a reaction (as you will in Lab 1), you are usually measuring the rate of the rate determining step. Sample Exercise 1 Use the equations provided to identify the reaction intermediates and the rate determining step for the reaction between four moles of hydrogen bromide and one mole of oxygen gas. Plan a strategy. 1. The reaction intermediates are species that are formed during elementary processes, so look for species that appear on the right side of one equation and the left side of following equations. 2. The rate determining step is the slowest step in a reaction mechanism. Look for the step labelled slow. Communicate the answer. The reaction intermediates are HOOBr and HOBr because these species are formed in early elementary processes and consumed in subsequent elementary processes. The slowest step is the reaction between one mole of HBr and one mole of O 2, so it is the rate determining step. The sample exercise illustrates a very important point about reaction mechanisms.

8 Acceptable elementary processes in reaction mechanisms generally show collisions between two particles or the breakdown of single particles. Simultaneous collisions between more than two particles are rare. Consider the overall equation from Sample Exercise 1: If this reaction proceeded in one step, it would require the simultaneous collision of five molecules (4 HBr and 1 O 2 ). The likelihood of more than two particles colliding at once is very, very, very small. A reaction mechanism that shows four or five molecules colliding at once is basically unacceptable. Catalysts In the previous lesson, a catalyst was defined as something that speeds up a chemical change by providing a reaction pathway with lower activation energy than that of the uncatalyzed reaction. Use of a catalyst often results in the replacement of a single high activation energy step with two or more lower activation energy steps. Here's an example: The thermal, uncatalyzed decomposition of methanoic acid: has high activation energy. Adding an acid catalyst results in a reaction mechanism involving three elementary processes.

9 In step one, the catalyst (H + ) combines with methanoic acid to form the reaction intermediate. The intermediate decomposes in step two to form a second reaction intermediate, HCO +. In step 3, the catalyst, H +, is regenerated. The second step is slow and determines the rate of the catalyzed reaction, but notice that it is not as slow as the single step in the uncatalyzed reaction. How can you tell? Activation energy is the barrier to a reaction. The greater the activation energy, the slower the reaction. In the diagram above, look for the step with the highest potential energy relative to the other steps to find the rate determining (i.e. slowest) step. Some Key Points About Reaction Mechanisms 1. Each elementary process has its a unique activation energy value. The step with the highest potential energy is the slowest or the rate determining step. 2. A catalyst is a type of reaction intermediate that appears on the reactants side of one elementary process and on the products side of a subsequent step. A hydrogen ion is the catalyst in the example above. It first appears as a "reactant" in step 1 and eventually reappears as a "product" in step 3. It is not part of the overall equation.

10 3. Reaction intermediates first appear as products side and are consumed in subsequent steps. and HCO + fit this description in the above example. These species are produced in steps 1 and 2 and consumed in steps 2 and 3. They are reaction intermediates. 4. Finally, the heat of reaction DH is not affected by the addition of a catalyst. It takes 72.8 kj of heat to drive the uncatalyzed reaction. The same amount of heat is needed to drive the catalyzed reaction.