REACTION KINETICS Kenneth E Schnobrich WHAT IS REACTION
REACTION KINETICS Kenneth E. Schnobrich
WHAT IS REACTION KINETICS n n It is the study of the factors that can influence the RATE or SPEED of a reaction. Rate describes how fast a reactant is used or how fast a product is formed in a reaction under a specified set of conditions (temperature, pressure, etc).
Factors that Effect Rate n n n Nature of the Reactants Concentration Temperature Surface Area (Particle Size) Catalyst Pressure (Gases only)
COLLISION THEORY n n Before a reaction can take place there must be COLLISIONS between the reacting species (atoms, molecules, ions) Successful collisions are those that result in a new product being formed (not all collisions are successful)
Successful Collisions n For a collision to be successful it must have: n n A sufficient amount of energy when the collision takes place The reacting particles must be in the proper orientation so that bonds can be broken and new bonds made.
Collision Theory H 2 and I 2 molecules moving randomly H 2 I 2 intermediate as a result of collision (unstable) HI product, the result of “successful” collisions – also some unreacted H 2 and I 2
NATURE OF THE REACTANTS n Since bonds must be broken and formed in most cases - the greater the number of bonds broken and formed the slower the rate of reaction. n Reactions involving ions in aqueous solutions are almost instantaneous - the ions are already present so bonds are formed easily.
Nature of Reactants (cont. ) n n In covalent compounds the bonds are more difficult to break and form so the reactions vary in rate. The more complex the reaction generally the slower the rate of reaction. n n CH 4 + 2 O 2 -> CO 2 + 2 H 2 O(fast) 2 C 8 H 18 + 25 O 2 -> 16 CO 2 + 18 H 2 O(slow)
Nature of Reactants Double bond + Double bonds 2 4 bonds + 4 bonds Made Broken Double bond 2 2 + 25 50 bonds Broken Double bonds 16 + 18 64 bonds 36 bonds Made
CONCENTRATION n Since reactions require collisions - the greater the number of reacting particles the greater the chances for successful collisions.
CONCENTRATION (cont. ) n n n In conclusion - the greater the concentration the faster the rate of reaction (more collisions -> more successful collisions) Analogy - students in the hallways Analogy – cars on the highway
TEMPERATURE n n n Temperature is a measure of average Kinetic Energy - the greater the temperature, the faster the particles are moving (have more energy when they collide) - this assures more successful collisions (bonds can be broken) Analogy - students in the hallways running to class Analogy – a posted minimum speed limit of 95 on the highway
Temperature All reactions require a certain minimum activation energy to get them started. Activated Complex Potential Energy n Activation Energy (forward) Activation Energy (reverse) Heat of Reactants Heat of Rx (DH) = Heat of Products Reaction Pathway
#molecules Temperature At T 1+ 25° there are molecules with sufficient activation energy T 1 + 25° Activation Energy
SURFACE AREA n The greater the size of the sample the smaller the surface area. The smaller we make the particle size the greater the surface area n n . Hydrogen ions hit the outer layer of atoms Mg + 2 HCl -> Mg. Cl 2 + H 2 Increased possibility for collisions …but not those in the Center of the lump With the smaller particle size, hydrogen Ions can reach more Mg atoms
More Particle interaction 2 inches 2 inches 1 in 4 inches 4 in ch es 1 in Surface Area Total Surface (6 sides) 4 in x 4 in = 16 in 2 x 6 = 96 in 2 Total Surface (48 sides) 2 in x 2 in = 4 in 2 x 6 = 24 in 2 x 8 = 192 in 2 Total Surface (384 sides) 1 in x 1 in =1 in 2 x 6 = 6 in 2 x 64 = 384 in 2
Catalyst n n Catalysts are chemical agents that speed up a reaction without changing either the initial reactants or final products and is not consumed by the reaction. n Heterogeneous catalysts – exist in a different phase than the reaction mixture (like the catalyst in your car’s exhaust system). n Homogeneous catalysts – exist in the same phase as the reaction mixture and frequently are part of an active intermediate to speed up the reaction. Catalysts lower the activation energy required for a reaction and thus increase reaction rate.
Catalysts n Catalysts work in several ways to lower the activation energy – n n n They can provide a surface for the reaction (putting H 2 O 2 on a cut, the blood cells provide a surface for the decomposition of the H 2 O 2) They can provide proper orientation for a successful collision (enzymes frequently act this way) They can stretch bonds and make them easier to break (Au used to weaken bonds between N & O in N 2 O) They can form an unstable intermediate (decomposition of H 2 O 2 by Na. I(aq)) In each case mentioned above they lower the activation energy needed for the reaction to take place.
Potential Energy Catalysts Activation Energy with a catalyst Activation Energy (forward) Activation Energy (reverse) Heat of Reactants Heat of Rx (DH) = Heat of Products Reaction Pathway
Equilibrium n n n In most systems there are two reactions occurring, the forward reaction and the reverse reaction. When the rate of the forward reaction and the reverse reaction are equal we say the system is in equilibrium (at a specified set of conditions). At equilibrium the concentrations of both reactants and products stay constant.
Equilibrium n n To indicate a system has reached equilibrium, double arrows are frequently used or an sign could be used, the arrows are preferred. There are two types of systems to be studied – n n Phase Equilibrium Chemical Equilibrium
Equilibrium n There also factors that affect a system in equilibrium – n n Concentration Temperature Pressure/Volume (gases) The underlying rule is referred to as Le. Chatlier’s Principle
Equilibrium and Rate of reaction Rate of forward rx decreases over time Rates are equal (Equilibrium) Rate of reverse rx increases over time Time
Equilibrium and Concentration reactants products Time Equilibrium Concentration [ ] Equilibrium Time Case 1 Time Case 2 At equilibrium the concentrations of reactants and products are constant and almost the same At equilibrium the concentrations of reactants are greater than the concentration of the products Case 3 At equilibrium the concentrations of reactants are less than the concentration of the products ** In each case, at equilibrium, the concentrations remain constant
Phase Equilibrium n n Phase equilibrium refers to a case where there is an equilibrium between two different phases of the same substance. H 2 O(l) = H 2 O(g) in a sealed container The rate of vaporization is equal to the rate of condensation Polar Water Molecule
Phase Equilibrium n In a saturated solution at a specific temperature you have another example of phase equilibrium - Rate of solution = the rate of crystallization Chloride ion (Cl-) Sodium ion (Na+) Polar Water Molecule
Chemical Equilibrium n N 2 H 2 Rate of reaction In a chemical reaction when the rate of the forward reaction is equal to the rate of the reverse reaction at a specified set of conditions the system is in equilibrium, N 2(g) + 3 H 2(g) = 2 NH 3(g) n Rates are equal (Equilibrium) NH 3 Time
Le. Chatlier’s Principle n n Le. Chatlier’s Principle states that a system will always try to reach equilibrium. If a stress, such as concentration change, temperature change, or pressure change (or volume change in systems involving gases) is placed on the system, the system will temporarily favor the forward or the reverse reaction until equilibrium is re-established. To indicate the reaction that is favored as a result of a stress we will use arrows of unequal length. In this case we are indicating that the forward reaction is favored (or there is a shift to the right)
D In Concentration n An increase in concentration of either the reactant(s) or product(s) will favor the reaction that uses the substance added. In the reaction for the Haber Process, let’s say we increase the concentration of N 2(g) n N 2(g) + 3 H 2(g) 2 NH 3(g) n Note that the forward reaction is favored (there is a shift to the right) n Eventually, over time, equilibrium will be re-established at a new equilibrium point Suppose H 2(g) is removed (the concentration is decreased) n N 2(g) + 3 H 2(g) 2 NH 3(g) n Note the reverse reaction is now favored to produce more H 2(g) (there is a shift to the left) n Eventually, over time, equilibrium will be re-established at a new equilibrium point
D In Concentration n Given the reaction, CH 4(g) + 2 O 2(g) = CO 2(g) + 2 H 2 O(g) fill in the following table Stress Inc [CH 4] Dec [O 2] Inc [H 2 O] Inc [O 2] Dec [CO 2] Shift R/L [CH 4] [O 2] [CO 2] [H 2 O]
D In Temperature n n An increase in temperature will always favor the endothermic reaction. A decrease in temperature favors the exothermic reaction. 2 H 2(g) + O 2(g) = 2 H 2 O(g) + 483. 6 k. J …. . exothermic reaction n n A temperature increase would favor the reverse reaction 2 H 2(g) + O 2(g) 2 H 2 O(g) There would be a shift to the left The [H 2] and [O 2] would increase and the [H 2 O] would decrease
D In Temperature Given the reaction, N 2(g) + 3 H 2(g) = 2 NH 3(g) + heat……exothermic Stress Inc T Dec T Shift R/L [N 2] [H 2] [NH 3]
D In Pressure/Volume (Gases) n n A change in pressure is comparable to a change in volume. If the volume decreases the pressure increases. If the volume increases the pressure decreases. An increase in pressure always favors the reaction that produces fewer mols of gas in the system. A decrease in pressure favors the reaction that produces more mols of gas in the system. After the pressure is changed equilibrium will be reestablished at a new equilibrium point.
D In Pressure/Volume (Gases) Given the reaction – N 2(g) + 3 H 2(g) 2 NH 3(g) 4 mols of reactants 1 mol of products Pressure Increase Reaction shifts to the right Fewer molecules N 2 and H 2 , more NH 3 , fewer molecules overall
D In Pressure/Volume (Gases) n n In cases where the number of mols of gas are the same for both reactants and products a pressure/volume change has no effect. Example - H 2(g) + I 2(g) 2 HI(g) 2 mols n 2 mols Complete the following table for the reaction: n N 2(g) + 3 H 2(g) Stress Inc. P Inc. V Shift (R/L) 2 NH 3(g) [N 2] [H 2] [NH 3]
Equilibrium Constants n n The equilibrium constant is known by several names but in general it is labeled, Keq. The various versions of equilibrium constant describe the make up of the system at various temperatures. Given a general reaction (at equilibrium) – n n n w. A + x. B = y. C + z. D (w, x, y, z would all represent numerical coefficients) By definition Keq = [C]y[D]z/[A]w[B]x - in other words products over reactants. If the Keq = 1. 8 x 10 -5 that would tell us that there are very few products in the system at equilibrium. The smaller the value for Keq the fewer the products.
Equilibrium Constants n n n For saturated solutions of salts the equilibrium constant is frequently called the Ksp (solubility product constant). The Ksp gives us an indication of how soluble a salt is in solution at equilibrium. Ag. Cl has a Ksp of 1. 8 x 10 -10, that should indicate to you that in the saturated solution of Ag. Cl there are very few Ag+(aq) and Cl Ag+(aq) + Cl-(aq) ions. Ag. Cl(s) For acids and bases they frequently talk about Ka (for acids) and Kb (for bases). The Ka for CH 3 COOH is 1. 85 x 10 -5. That should tell us there are few H+(aq) and CH 3 COO-(aq) ions in solution (it is a weak acid). CH 3 COOH(aq) H+(aq) + CH 3 COO-(aq).
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