Kinetics and Equilibrium

Kinetics and Equilibrium

Kinetics Kinetics - concerned with the rates of chemical reactions.
Basic concept; in order for a reaction to occur, reactant particles must collide Collision Theory – collisions between particles can produce a reaction if both the spatial orientation and energy of the colliding particles are conducive to a reaction.

Collision Theory An effective collision of oxygen and hydrogen molecules produces water molecules.

Collision Theory An ineffective collision of oxygen and hydrogen molecules produces no reaction; the reactants bounce apart unchanged.

Collision Theory Particles must collide in a specific arrangement for the collision to be effective, allowing products to form.

Collision Theory Properly positioned particles during a collision allows the activated complex to form Activated complex – temporary, intermediate product. Allows for the rearrangement of atoms to form new products OR the reformation of reactants

Collision Theory The minimum energy that colliding particles must have in order to react is called the activation energy. Think of it as a barrier that reactants must cross before products can form. When two reactant particles collide, they may form an activated complex. An activated complex is an unstable arrangement of atoms that forms for a moment at the peak of the activation-energy barrier.

Collision Theory When two reactant particles collide, they may form an activated complex. The activated complex forms only if the colliding particles have enough energy and if the atoms are oriented properly. The lifetime of an activated complex is typically about seconds. Its brief existence ends with the reformation of the reactants or with the formation of products. Thus, the activated complex is sometimes called the transition state.

Potential Energy Diagrams
Rates of Reaction Potential Energy Diagrams The activation-energy barrier must be crossed before reactants are converted to products. Note: Energy is measured in kJ (kilojoules)

Potential Energy Diagram
Recall: chemical bonds contain energy  specifically that energy is potential energy. These diagrams illustrate the potential energy change that occurs during a chemical reaction. Y-axis: change in potential energy X-axis: reaction coordinate – the progress of the reaction In order for a reaction to occur: reactants must have sufficient potential energy to collide effectively

The figure below illustrates the progress of a typical reaction
The figure below illustrates the progress of a typical reaction. Over time, the amount of reactant decreases and the amount of product increases.

Collision / Reaction Process
1. Energy is applied to reactants. Reactant particles absorb energy;  kinetic energy. Remember that reactants are more unstable then their products. Effective collisions allow kinetic energy to be converted to potential 2. 3. 1.

Potential Energy (P.E.) Diagram
P.E. of reactants Activation Energy (A.E.) Heat of Rx./ Difference in P.E. b/w reactants and products P.E. of products A.E. of forward Rx. /Difference b/w the reactants and the activated complex A.E. of reverse Rx./ Difference b/w the products and the activated complex E F

P.E. of activated complex
P.E. of Products P.E. of Reactants

Potential Energy Diagrams
Rates of Reaction Potential Energy Diagrams Remember: An endothermic reaction absorbs heat, and an exothermic reaction releases heat. Endothermic or Exothermic

Exothermic Products have less potential energy then reactants
Energy is released, and the temperature of the surrounding increases Heat of the Rx is always negative

Exothermic Reactant has more P.E. then products. Net loss of energy.
More energy was released to form the product than was gained to form the activated complex. ΔH = negative

Endothermic Products have more potential energy then the reactants
Energy is absorbed, temperature of the surrounding decreases Heat of reaction is always positive

Endothermic Product has more P.E. then reactants. Net gain of energy.
More energy was absorbed to form the activated complex than was released to form the product. ΔH = positive

Topic 9 - Kinetics & Equilibrium Note Packet
Complete Practice Problems #9-14 on page 161. Complete Practice Problems #23-25 on page 163. Complete Practice Problems #26-32 on page 165.

Factors Affecting Rates of Rx
Nature of the reactants Concentration Surface area/ Particle size Pressure – gasses only Presence of a catalyst Temperature

Nature of the Reactants Reactions involve the breaking of existing bonds and the formation of new ones. Ionic bonds – faster to react low bond dissociation energy/ easy to break Covalent bonds – slower to react high bond dissociation energy/ hard to break Greater number of bonds must be broken before the reaction can occur. Breaking more bonds requires that the particles must have more energy when they collide

Concentration If one or more of the reactants is increased, the reaction will proceed faster. Cramming more particles into a fixed volume increases the concentration of reactants, and, thus, the frequency of collision. Increased collision frequency leads to a higher reaction rate.

Surface Area/ Particle Size More surface area exposed, the greater chance for reactant particles to collide. Only atoms at the surface of the metal are available for reaction. Dividing the metal into smaller pieces increases the surface area and the number of collisions.

Pressure An increase in pressure on a gas, increases the concentration of gaseous particles, increasing the rate of reaction for gases Pressure has little or no effect on solids and liquids

Presence of a Catalyst a catalyst is a substance that increases the rate of a reaction without being used up during the reaction. Catalysts permit reactions to proceed along a lower energy path. A catalyst is not consumed during a reaction. Therefore, it does not appear as a reactant in the chemical equation. Instead, the catalyst is often written above the yield arrow, as in the equation above. 2H2(g) + O2(g) H2O(l) Pt

Advantage of a catalyst
Heat of the reaction – difference between the P.E. of the reactants and the P.E. of the products ΔH does not change with the presence of a catalyst. Activation Energy of the forward reaction or the reverse reaction are both lowered.

Catalysts At normal body temperature (37C), reactions in the body would be too slow without catalysts. The catalysts that increase the rates of biological reactions are called enzymes. When you eat a meal containing protein, enzymes in your digestive tract help break down the protein molecules in a few hours.

Catalysts An inhibitor is a substance that interferes with the action of a catalyst. Some inhibitors work by reaction with, or “poisoning,” the catalyst itself. Thus, the inhibitor reduces the amount of catalyst available for a reaction. Reactions slow or even stop when a catalyst is poisoned.

Temperature Usually, raising the temperature speeds up a reaction. Lowering the temperature usually slows down a reaction. At higher temperatures, particles move faster. The frequency of collisions increases. Reacting particles will have more energy, allowing the percentage of particles that have enough kinetic energy to slip over the activation-energy barrier. Thus, an increase in temperature causes products to form faster. At a higher temperature, not only will there be more collisions, but the reacting particles will have more energy, maing it more likely that the collisions will be effective

Surprise Question Which of the following factors could be increased in order to decrease a reaction rate? A. Catalyst concentration B. Concentration C. Temperature D. Particle size

Topic 9 - Kinetics & Equilibrium Note Packet
Complete Practice Problems #1-8 on page 158.

Progress of Chemical Rx
N2O4(g) → 2NO2(g) A B The rate of the rx depends in part on the concentration of the reactants. The rate at which A, (N2O4) forms B, (2NO2) can be expressed as a change in concentration A, (∆A) with time (∆t) The initial concentration of A at t1 verses the final concentration of A at t2. Rate=− ∆A ∆t =− concentration A 2 −concentration A 1 t 2 − t 1

Rate Laws ΔA [A] Δt ΔA rate = = k × [A] Δt
The rate of disappearance of A is proportional to the concentration of A. The proportionality can be expressed as the concentration of A, [A], multiplied by a constant, k. Rate law, an expression for the rate of a reaction in terms of the concentration of the reactants. Specific rate constant (k) for a reaction is a proportionality constant relating the concentrations of reactants to the rate of the reaction. ΔA Δt [A] rate = = k × [A] ΔA Δt

Rate Laws The value of the specific rate constant, k, in a rate law is large if the products form quickly; the value is small if the products form slowly. rate = = k × [A] ΔA Δt

First Order Reaction N2O4(g) → 2NO2(g) A B
Conversion of A  B in one step reaction Reaction rate is directly proportional to the concentration of only one reactant. As the reaction progresses, the rate of the rx decreases

Higher Order Reactions
𝑎A+𝑏B →𝑐C+𝑑D For the reaction of A with B, the rate of reaction is dependent on the concentrations of both A and B. When each exponent in the rate law equals 1 (that is, x = y = 1) the reaction is said to be first order in A and first order in B. The overall order of a reaction is the sum of the exponents for the individual reactants. i.e. second order reaction overall. rate = k[A]a[B]b

Equilibrium Reaction going “forward” Current understanding: Potential energy diagrams depict rx in 1 direction; L  R in a “forward” direction. Reactions can occur in both directions; Dynamic Equilibrium R P Reaction going “backward” R P The reactants collide to form the activated complex and then they form products. The opposite is also true. Products can collide to form the activated complex and become reactants. Reaction in dynamic equilibrium R P

Equilibrium 2SO2(g) + O2(g) 2SO3(g)
Equilibrium – when both the forward and reverse reactions occur at the same rate Rxs are shown proceeding in both directions with a double arrow 2 opposite processes are occurring at the same time = equal rates. Must occur in a closed system Reactants and products can not leave the system 2SO2(g) + O2(g) 2SO3(g)

Equilibrium Quantities of reactants and products ARE NOT equal when at equilibrium. Rates of reaction of reactants and products ARE equal when at equilibrium. Constant flux between reactants and products. If one side is increased or decrease, the other side will compensate for the change.

Physical Equilibrium Phase Equilibrium
Ex. condensation, evaporation, and triple point Solution Equilibrium Ex. Solids in liq. or gases in liq. If additional solute (solid or gas) is added to a saturated soln. in a closed system, the solute will still dissolve, but it dissolves at the same rate that it also recrystallizes back out of the solution. This of how this physical equilibrium between phase and solution can be effect. As you increase temperature, solubility of a solid in a liquid increases, but with increase in temperature, the solubility of a gas decreases.

If a system is not closed, equilibrium cannot be reached.
Chemical Equilibrium If a system is not closed, equilibrium cannot be reached. Forward Reaction Rate slows as reactants are consumed Reverse Reaction Rate increases as products are formed Dynamic Equilibrium Both forward and reverse reactions occur at the same rate. 𝐇 𝟐 + 𝐈 𝟐 →𝟐𝐇𝐈 𝟐𝐇𝐈→𝐇 𝟐 + 𝐈 𝟐

Le Chatelier’s Principle
Used to explain how a system at equilibrium responds to relive any stress on the system. If you had more to the left side (reactants), the lever will tip to the left. In order to restore balance, the system responds by making more of what is on the right (products) and vice versa. Stressors include changes in: Temperature Concentration Pressure N 2 g + 3H 2 (g) ↔ NH 3 g +heat Reactants Products

When achieving equilibrium …
Particles move from high concentration  low concentration (diffusion) Heat travels from hot  cold (melting) Pressure changes flow from high  low

Concentration Changes Effect of Increasing the Concentration of CH4 CH 4 g + H 2 O g ↔ 3H 2 g + CO (g) + heat Stress Effect System Shift + CH 4 - H 2 O → + 3H 2 + CO + Heat Increase Decrease Away from stress Effect of Decreasing the Concentration of CH4 CH 4 g + H 2 O g ↔ 3H 2 g + CO (g) + heat Stress Effect System Shift - CH 4 + H 2 O ← - 3H 2 - CO Heat Decrease Increase Toward the stress

Temperature Changes Effect of Increasing the Heat CH 4 g + H 2 O g ↔ 3H 2 g + CO (g) + heat Effect System Shift Stress + CH 4 + H 2 O ← - 3H 2 - CO + Heat Increase Away from stress Decrease Effect of Decreasing the Heat CH 4 g + H 2 O g ↔ 3H 2 g + CO (g) + heat Effect System Shift Stress - CH 4 - H 2 O → + 3H 2 + CO Heat Decrease Toward the stress Increase

Le Chatelier’s Principle Pressure Changes
CO 2 (g)↔ CO 2 (aq) Pressure changes only have effects on gases As pressure , concentration of gases . Rx will shift to the right, causing more dissolved CO2 As pressure , concentration of gases . Rx. Shifts to the left, forming more gaseous CO2 Remember: Particles travel from high pressure to low pressure. THINK SODA POP. Pressure in the bottle is greater that that outside the bottle. What happens when you open it?

Equilibrium Constants
Think of it this way: 𝑎A+𝑏B ↔𝑐C+𝑑D a mol of reactant A and b mol of reactant B react to produce c mol of product C and d mole of product D at equilibrium. Coefficients: amounts in moles of reactants & products

Equilibrium Constant [C]c x [D]d Keq = [A]a x [B]b
The equilibrium constant (Keq) is the ratio of product concentrations (mol/L) to reactant concentrations (mol/L) at equilibrium. The value of Keq depends on the temperature of the reaction. The size of the equilibrium constant indicates whether reactants or products are more common at equilibrium. 𝐾 eq >1, products favored over equilibrium 𝐾 eq <1, reactants favored over equilibrium From the general equation, each concentration is raised to a power equal to the number of moles of that substance in the balanced chemical equation. Reminder: concentrations are Molarity

Equilibrium Constant N2O4(g) 2NO2(g) Keq = [C]c x [D]d [A]a x [B]b
The colorless gas dinitrogen tetroxide (N2O4) and the brown gas nitrogen dioxide (NO2) exist in equilibrium with each other. A liter of the gas mixture at equilibrium contains mol of N2O4 and mol of NO2 at 10oC. Write the expression for the equilibrium constant (Keq) and calculate the value of the constant for the reaction. N2O4(g) NO2(g) Keq = [C]c x [D]d [A]a x [B]b Keq = [NO2]2 [N2O2] NOTE THAT THE MOLAR CONCENRATIONS HAVE TO EXIST AT EQUILIBRIUM Keq = (0.030 mol/L)2 ( mol/L) = (0.030 mol/L x mol/L) Keq = 0.20 mol/L = 0.20 Reactants favored

Equilibrium Constant Keq = [C]c x [D]d [A]a x [B]b Keq = [NH3]2
The reversible reaction N 2 g + 3H 2 g ↔ 2NH 3 (g) produces ammonia, whih is a fertilizer. At equilibrium, a 1-L flask contains 0.15 mol H2, mol N2, and 0.10 mol NH3. Calculate Keq for the reaction. Keq = [C]c x [D]d [A]a x [B]b Keq = [NH3]2 [N2][H2]3 Keq = 12 Products favored

One is the inverse of the other.
Equilibrium Constant For the same mixture, under the same conditions described in the previous problem, calculate Keq for 2NH 3 g ↔ N 2 g + 3H 2 g . How is Keq for a forward reaction related to the Keq for a reverse reaction? Keq = [C]c x [D]d [A]a x [B]b Keq = [NH3]2 [N2][H2]3 Keq = 0.083 Reactants favored One is the inverse of the other.

2 mol HI = 1.56 mol at equilibrium
Equilibrium Constant One mole of colorless hydrogen gas and one mole of violet iodine vapor are sealed in a 1-L flask and allowed to react at 450oC. At equilibrium, mol of colorless hydrogen iodide is present, together with some of the reactant gases. Calculate Keq for the reaction. PROBLEM: The molar concentrations are NOT given in equilibrium. Initial molar concentrations are 1 mol of H2 and 1 mol of I2 react to produce 2 mol of HI 2 mol HI = 1.56 mol at equilibrium H2(g) + I2(g) ↔ 2HI(g) The molar concentrations are not given at equilibrium. They give the initial molar concentrations: 1 mole of hydrogen and 1 mole of iodine react to yield 2 moles of HI x x = 2 x + x = 1.56 mol 2x = 1.56 mol x = mol H2(g) + I2(g) ↔ 2HI(g)

mol H2 = mol I2 = (1.00 mol – 0.780 mol) = 0.22 mol
Continued… One mole of colorless hydrogen gas and one mole of violet iodine vapor are sealed in a 1-L flask and allowed to react at 450oC. At equilibrium, mol of colorless hydrogen iodide is present, together with some of the reactant gases. Calculate Keq for the reaction. You must subtract the # of moles that react from the number of moles originally present in order to determine the number moles of remaining at equilibrium.. H2(g) + I2(g) ↔ 2HI(g) X = mol mol H2 = mol I2 = (1.00 mol – mol) = 0.22 mol Recall from the problem that 1.56 mol of HI remained. The 0.78 mol of H2 and I2 was what reacted from the original 1 mol. So how many mol remain on the reactant side that did not react? Keq = [C]c x [D]d [A]a x [B]b Keq = [HI]2 [H2] x [I2] Keq = 50.28 Products favored

Equilibrium Constant Suppose the following system reaches equilibrium, N 2 g + O 2 g ↔2NO(g) Analysis of the equilibrium mixture in a 1-L flask gives the following results: N2 = 0.50mol, O2 = 0.50mol, and NO = 0.020mol. Calculate Keq for the reaction. Keq = [C]c x [D]d [A]a x [B]b Keq = 1.6 10-3 Reactants favored One is the inverse of the other.

Equilibrium Constant At 750°C the following reaction reaches equilibrium in a 1-L flask. H 2 g + CO 2 g ↔ H 2 O g +CO(g) Analysis of the equilibrium mixture gives the following results: H2 = mol, CO2 = 0.053mol, H2O = 0.047mol, and CO = 0.047mol. Calculate Keq for the reaction. Keq = [C]c x [D]d [A]a x [B]b Keq = 0.79 Reactants favored One is the inverse of the other.

Equilibrium & Reversible Rx
In nearly all reversible reactions, one reaction is favored over the other. So which is favored?? Consider the decomposition of carbonic acid in water. The forward reaction is spontaneous and releases free energy. The combination of carbon dioxide and water to form carbonic acid is a nonspontaneous reaction. H2CO3(aq) CO2(g) + H2O(l) <1% >99%

Spontaneous Reactions Nonspontaneous Reactions
C(s) + O2(g)  CO2(g) CO2(g)  C(s) + O2(g) Rx that occur naturally, on their own. Favors formation of products at equilibrium Releases free energy. Free energy – energy available to do work Rx that tend not to occur, or not efficiently. Does NOT favor formation of products at equilibrium Does NOT release free energy Important to note that the words spontaneous and nonspontaneous do not refer to how fast reactants go to products.

To be or not to be… spontaneous
Some rx that are nonspontaneous at one set of conditions may be spontaneous at another set of conditions. Changing temperature or pressure Ex. Photosynthesis: nonspontaneous rx. Needs sunlight to be driven to completion. Coupled to a spontaneous rx This occurs in complex biological processes

Entropy & Enthalpy What determines whether a spontaneous chemical or physical change will occur? Entropy – a measure of the disorder or randomness of a system Enthalpy – a measure of energy change between initial and final states.

ENTROPY Low Entropy High Entropy

Entropy Physical Changes in Entropy Chemical Changes in Entropy
Solid Liquid Gas Physical Changes in Entropy Low High Entropy Chemical Changes in Entropy Compounds represent a state of low entropy, greater order than the free elements of which they are composed of The side of the Rx. with greater # of molecules has greater amount of entropy

Entropy – What does a Rx want?
There is a tendency of nature to change from a state of order to a state of greater randomness or disorder. Rx are expected to go in the direction of greater entropy/ more randomness. Spontaneous rx favor high entropy; less order Nonspontaneous rx favor low entropy; more order Disorder is more probable then order. Big Bang Theory Ice cubes melting Organic decomposition/ rotting 2nd Law of Thermodynamics

Entropy For a given substance, the entropy of the gas is greater than the entropy of the liquid or the solid. Thus, entropy increases in reactions in which solid reactants form liquid or gaseous products.

Entropy Entropy increases when a substance is divided into parts.
For instance, entropy increases when an ionic compound dissolves in water.

Entropy Entropy tends to increase in chemical reactions in which the total number of product molecules is greater than the total number of reactant molecules.

Entropy Entropy tends to increase when the temperature increases. As the temperature rises, the molecules move faster and faster, which increases the disorder.

Which of the following is ALWAYS true of spontaneous reactions?
They produce heat and are not reversible at the stated conditions. They release free energy and favor the formation of products at the stated conditions. They are coupled with a nonspontaneous reaction and are easily reversible at the stated conditions.

Which of the following would have an increase in the entropy of the reaction system?
A. 2NH4NO3(s)  2N2(g) + 4H2O(l) +O2(g) B. 2H2(g) + O2(g)  2H2O(l) C. C3H8(g) + 5O2(g)  3CO2(g) + 4H2O(l) D. 2Fe(s) + O2(g) + 2H2O(l)  2Fe(OH)2(s)

ENTHALPY a measure of energy change between initial and final states.
Endothermic – systems move towards higher energy states Exothermic – systems move towards lower energy states

Enthalpy – What does a Rx what?
Nature favors a state of lower energy (enthalpy). Reactions aim to achieve low energy/ low enthalpy Exothermic – Heat/ energy is lost. Rx moves from high P.E. in reactants to lower P.E. in products. Endothermic – Heat/ energy is absorbed. Rx moves from low P.E. in reactants to a higher P.E. in the products. Most reactions are expected to be exothermic.

How Enthalpy Changes and Entropy Changes Affect Reaction Spontaneity
Is the reaction spontaneous? Decreases (exothermic) Increases (more disorder in products than in reactants) Yes Increases (endothermic) Increases Only if unfavorable enthalpy change is offset by favorable entropy change Decreases (less disorder in products than in reactants) Only if unfavorable entropy change is offset by favorable enthalpy change Decreases No

Would the following exothermic reaction be spontaneous
Would the following exothermic reaction be spontaneous? Explain why or why not. 2KClO3(s) KCl(s) +3O2(g) Two molecules of solid are transformed into 2 molecules of solid and 3 molecules of gas, so entropy is increased in the reaction. A reaction that is exothermic with an increase in entropy will be spontaneous.

Change in Free Energy ΔG = ΔH – TΔS Expressed as Gibbs free energy.
Josiah Gibbs, the scientist who defined this thermodynamic property. Free energy can either be released or absorbed during a physical or chemical process. ΔS is the change in entropy. ΔH is the change in enthalpy. T is the temperature in Kelvins. ΔG = ΔH – TΔS

Gibbs Free Energy When the value of ΔG is negative, the process is spontaneous. When the value is positive, the process is nonspontaneous.

ΔG = ΔH – TΔS C(s) + O2(g) CO2(g)
The entropy change for the following reaction at 298 K is 3.0 J/mol·K, and the enthalpy change is –394 kJ/mol. Calculate the Gibbs free energy change and determine whether the reaction will occur spontaneously. C(s) + O2(g) CO2(g) ΔG = –394 kJ/mol – (298 K  kJ/mol·K) ΔG = –395 kJ/mol The reaction is spontaneous.

Kinetics and Equilibrium

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