Reaction Rates CHM 1: Chapter 18 CHM Hon: Chapter 17 & 18

PART 1 RATES OF REACTION

Collision Theory Reactions occur at different rates Speed = Distance in a given amount of time –Shelia runs 10 meters in 4 seconds 10 m / 4s 5 m / 2s 2.5 m/s

Collision Theory Rate = a measure of the speed of any change that occurs within an interval of time. –Rate of a Chemical Change or reaction = amount of reactant changing per unit time

Collision Theory Collision Theory: atoms, ions and molecules can react to form products when they collide with one another, provided that the colliding particles have enough kinetic energy –Kinetic energy = the energy possessed by a body because of its motion

Collision Theory Particles lacking the necessary kinetic energy to react bounce apart unchanged with they collide

Collision Theory If enough energy is applied to a molecule, the bonds holding the molecule together can break! Substances supplied with enough energy decompose to simpler substance or reorganize themselves into new substances.

Collision Theory Activation Energy = the minimum energy that colliding particles must have in order to react

Collision Theory When 2 reactant particles with enough activation energy collide, an activated complex (new entity) is formed

Collision Theory

Activated Complex = an unstable arrangement of atoms that forms momentarily at the peak of the activation- energy barrier (Brief existence!) –Forms only if: Colliding particles have sufficient energy Atoms are oriented properly –Also called… Transition state Reformation of the reactants or formation of products. Both outcomes equally likely!

Pop Quiz!

Collision Theory Particle size –Surface area of reactant affects the reaction rate Smaller particle size  larger the surface area High surface area  –High amount of reactant exposed –High collision frequency –High reaction rate

Collision Theory Increasing surface area –Solid: dissolve into solution, grind into powder

Collision Theory

Catalysts = a substance that increases the rate of reaction without being used up during the reaction –Permit reactions to proceed along a lower energy path –More reactants have the energy to form products within a given time

Reaction with catalyst vs. no catalyst

Collision Theory Catalysts –Important to the human body –Enzymes = catalysts –Without enzymes, digestion would take years!!

Collision Theory Inhibitor = a substance that interferes with the action of a catalyst –Reduces the amount of functional catalyst available –Reactions slow or stop

Collision Theory Heterogeneous catalyst = the catalyst exists in a different physical state that the reaction it catalyzes –Ex. Catalytic converter Homogeneous catalyst = the catalyst exists in the same physical state as the reaction it catalyzes –Ex. Both the enzyme and reaction are in aqueous solution

PART 2 REVERSIBLE REACTIONS & EQUILIBRIUM

Reversible reactions Reversible reaction = a reaction in which the conversion of reactants to products and the conversion of products to reactants occur simultaneously

Reversible reactions Chemical Equilibrium = a state of balance; when rates of the forward and reverse reactions are equal

Reversible reactions Chemical Equilibrium –No net change occurs in the amounts of the components of the system –Dynamic state: both forward and reverse reactions continue

Reversible reactions Equilibrium position = the relative concentrations of the reactants and products at equilibrium –Indicates whether the reactants or products are favored If A reacts to give B, eq. mixture contains more B, then…

Reversible reactions

Catalysts –Speeds up both forward and reverse reactions equally

Le Chatelier’s Principle When the equilibrium of a system is disturbed, the system makes ajustments to restore equilibrium –Shift in Equilibrium position

Le Chatelier’s Principle Le Chatlier’s Principle If a stress is applied to a system in dynamic equilibrium, the system changes in a way that relieves the stress.  Stresses that upset the equilibrium  changes in the concentration of reactants or products  changes in temperature  changes in pressure

Le Chatelier’s Principle Concentration –System adjusts to minimize the effects of the change –Adding product to a reaction pushes in the direction of reactants and vice a versa

Le Chatelier’s Principle Temperature –Increasing the temperature causes the equilibrium position of a reaction to shift in the direction that absorbs heat Heat is considered a product!!

Le Chatelier’s Principle Pressure –Change in pressure only affects gaseous equilibria that have an unequal number of moles of reactants and products

Le Chatelier’s Principle

Equilibrium Constants Equilibrium constant = (K eq ) the ratio of product to reactant concentrations at equilibrium –Each concentration raised to a power equal to the # of moles of that substance in the balanced chemical equation

Equilibrium Constants

Brackets indicate concentration (mol/L) Value of K eq depends on the temperature The size of Keq shows whether products or reactants are favored at equilibrium

Equilibrium Constants

K eq =12

Equilibrium Constants

PART 3 Solubility Equilibrium

Solubility Product Constant Ionic compounds have different solubilities Most salts are somewhat soluble –Most salts of the alkali metals are soluble in water (slightly / sparingly soluble) –Compounds that contain phosphate, sulfide, sulfite or carbonate ions are generally insoluble

Solubility Product Constant When salts completely dissolve it is a one way reaction… When salts do no completely dissolve it becomes a reversible reaction and an equilibrium is established…

Solubility Product Constant Solubility product Constant K sp = an equilibrium constant for the dissolving of a sparingly soluble ionic compound in water  The product of the concentrations of the ions each raised to the power of the coefficient of the ion in the dissociation equation

Solubility Product Constant

The Common Ion Effect Common Ion = an ion that is found in both salts in a solution Common Ion effect = the lowering of the solubility of an ionic compound as a result of the addition of a common ion

The Common Ion Effect

The K sp can be used to predict whether a precipitate will form when solutions are mixed. If the product of the concentrations of two ions in the mixture is greater than K sp of the compound formed from the ions, a precipitate will form.

The Common Ion Effect

Part 4 Entropy & Free Energy

Free Energy and spontaneous reactions Many chemical and physical processes release energy that can be used to bring about other changes Free energy = energy that is available to do work –Not necessarily used efficiently! Example: internal-combustion engine (cars) only 30 % efficient No process 100% efficient

Free Energy and spontaneous reactions Spontaneous reaction = occurs naturally and favors the formation of products at the specified conditions. Nonspontaneous reaction = a reaction that does not favor the formation of products at the specified conditions

Free Energy and spontaneous reactions Spontaneous reactions –Produce substantial amounts of products at equilibrium and release free energy Nonspontaneous reactions –Do not give substantial amounts of products at equilibrium

Free Energy and spontaneous reactions Spontaneous or nonspontaneous? SPONTANEOUS

Free Energy and spontaneous reactions Note: –Spontaneous and nonspontaneous do not refer to rate of reaction! May take 100’s of years, but still my be considered spontaneous if products are favored! May speed up reaction by introducing energy as heat (increase temperature)

Free Energy and spontaneous reactions Note: –Some nonspontaneous reactions may be made to occur if it is coupled to a spontaneous reaction (one that releases free energy) Common in Biological processes taken place in living organisms

Entropy Enthalpy = the heat content of a system at constant pressure –Heat changes accompany most chemical and physical processes Exothermic = release of heat Endothermic = absorption of heat

Entropy The Combustion of carbon is exothermic and spontaneous –Heat is released during the reaction  kj / mole of carbon burned

Entropy Sometimes a reaction may be spontaneous but absorb heat –As this reaction turns from solid to liquid (melts) 1 mol of ice at 25 C absorbs 6.0 kj/mol of heat from its surroundings

Entropy Entropy = a measure of the disorder of a system Law of Disorder: (entropy change) the natural tendency is for systems to move in the direction of maximum disorder or randomness –An increase in entropy favors the spontaneous chemical reaction; a decrease favors the nonspontaneous reaction

Entropy

Enthalpy, Entropy & Free Energy In every chemical reaction, heat is either released or absorbed and entropy or randomness either increases or decreases The size and direction of enthalpy changes and entropy changes together determine whether a reaction is spontaneous (whether it favors products and releases free energy)

Enthalpy, Entropy & Free Energy ~ Entropy and enthalpy changes affect the spontaneity of chemical reactions! ~ Either of the two (not both) can be unfavorable for a spontaneous process

Gibbs Free-Energy Gibbs free-energy change = the maximum amount of energy that can be coupled to another process to do useful work. Temperature in K Change in entropy Change in enthalpy

Gibbs Free-Energy All spontaneous processes release free energy The numerical value of is negative in spontaneous processes because the system loses free energy!

Gibbs Free-Energy Solid calcium carbonate decomposes to give calcium oxide and carbon dioxide –Entropy increases: one of the products formed from the (s) reactant is a (g) –Endothermic reaction –Nonspontaneous reaction

PART 5 Progress of Chemical Reactions

Rate Laws Rate of reaction depends partly on the concentrations of the reactants. The rate at which A yields B is expressed as the change in A over time ( ).

Rate Laws A is the reactant –Reactant is decreasing –The concentration of A is smaller at a later time than initially, therefore will always be negative

Rate Laws Rate law = an expression for the rate of a reacion in terms of the concentration of reactants Specific rate constant = (k) a proportionality constant relating the concentrations of reactants to the rate of the reaction

Rate Laws The value of the specific rate constant, k, is large if the products form quickly –Less time! The value is small if the products form slowly –More time!

Rate Laws Order of a reaction = the power to which the concentration of a reactant must be raised to give the experimentally observed relationship between concentration and rate First-order reaction = one in which the reaction rate is directly proportional to the concentration of only one reactant. –Reaction rate is directly proportional to the concentration of only one reactant

Rate Laws First-order reaction –Example: Conversion of A to B in a one step reaction

Rate Laws When two substances react to give products: –Example, double-replacement reaction

Rate Laws Note: –When each of the exponents a and b in the rate law equals 1, the reaction is said to be… First-order in A First-order in B Second-order over-all

Rate Laws Example problem: Finding the order of a reaction from experimental data The rate law for the one-step reaction aA  B is of the form: Rate=k[A] a. From the data in the following table, find the order of the reaction with respect to A and the overall order of the reaction.

Rate Laws

Practice problem: –Show that the unit of k for a first-order reaction is a reciprocal unit of time, such as a reciprocal second (s -1 ) –Answer:

Reaction Mechanism Elementary reaction = a reaction in which reactants are converted to products in a single step –Has only one activation-energy peak between products and reactants and therefore only one activated complex –Most chemical reactions consist of a number of elementary reactions!

Reaction Mechanism Reaction mechanism = The series of elementary reactions or steps that take place during the course of a complex reaction

Reaction Mechanism The peaks correspond to the energies of the activated complexes. Each valley corresponds to the energy of an intermediate Reaction progress curve

Reaction Mechanism Intermediates do not appear in the overall chemical equation for a reaction. Example: The reaction mechanism for the decomposition of nitrous oxide

Reaction Mechanism IMPORTANT: The overall chemical equation for a complex reaction fives no information about the reaction mechanism. –Reaction mechanisms must be determined experimentally.