Lecture 11 Notes on Chemistry 2025 PDF
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Uploaded by FervidDune
ETH Zurich
2025
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These notes cover lecture 11 on chemistry, in particular, chemical kinetics, and cover various topics and questions. The document details announcements, problem sets, office hours, and a final exam schedule.
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Lecture #11, p. 1 Lecture 11: Announcements Today: Brown Ch. 14 Chemical Kinetics II 14.5 Temperature and Rate 14.6 Reaction Mechanisms 14.7 Catalysis P...
Lecture #11, p. 1 Lecture 11: Announcements Today: Brown Ch. 14 Chemical Kinetics II 14.5 Temperature and Rate 14.6 Reaction Mechanisms 14.7 Catalysis Problem Set 10: Due before Exercise #11 tomorrow; upload on Moodle link Problem Set 11: Posted on Moodle; due before Exercise #12 next week Study Center: Wednesdays 18:00–20:00 in ETA F 5 Office Hours: Prof. Norris and Brisby, Thursdays 17:00–18:00 in LEE P 210 Final Exam: Monday, February 3, 2025, at 8:30–10:30 Chemistry Lecture #11, p. 2 Lecture 12 Next Week: Brown Ch. 15 Chemical Equilibrium 15.1 The Concept of Equilibrium 15.2 The Equilibrium Constant 15.3 Understanding and Working with Equilibrium Constants 15.4 Calculating Equilibrium Constants 15.5 Le Châtelier’s Principle 19.7 Free Energy and the Equilibrium Constant Chemistry Lecture #11, p. 3 Red Thread Where are we going? Acid-Base Catalysis Properties Christmas! Kinetics Batteries Equilibrium Chemistry Lecture #11, p. 4 Review In Lecture 10, we discussed kinetics Chemical kinetics tells us about the speed of reactions/processes Speed quantified by reaction rate Rate depends on reactant concentrations according to rate law: Rate = ![A]! [B]" [C]# ⋯ ( + * + + ⋯ = reaction order Rate law and reaction order determined by experiment! Rate law can give concentrations versus time Result depends on reaction order (zero-, first-, second-, etc.) Half-life, -$/& , is when half of the initial reactant is gone Chemistry Lecture #11, p. 5 Important Open Questions Given: Rate = ![A]! [B]" [C]# ⋯ , what do ), *, + mean? Given two spontaneous reactions, why can one be fast and one slow? Why do we often heat our reactions? What role does temperature play in kinetics? Chemistry Lecture #11, p. 6 Speed of Ants versus Temperature Ants are “cold blooded” Body temperature changes with ambient Cool experiment a century ago! If ∆" = +10°C, ant speed doubles! Related to chemistry? Can we explain/understand? Today: we will see! H. Shapley, PNAS 6 , 204 (1920). Chemistry Lecture #11, p. 7 Let's start with Why Heat Chemical Reactions? Collision Model Because reaction rates increase with T ! Iiii reasons torehact.tn For two molecules to react they must first collide Reaction rate ∝ collisions per second Collision frequency depends on concentrations Agrees with rate law! But also depends on temperature ⇒ At higher T, molecules move faster more often At higher T, molecules collide more forcefully according to the kinetic-molecular theory (Lecture #2) To react, molecules must collide with proper kinetic energy and orientation Thus, only small percentage of collisions leads to reaction ! Chemistry Lecture #11, p. 8 Why doorientation and kinetic energy of molecules matter Why Does Orientation Matter? Generic reaction Correct orientation Wrong orientation: “A” hits wrong side of “BC” How can it form product? Reactants must collide correctly to form product ! Chemistry Lecture #11, p. 9 Why Does Kinetic Energy Matter? Svante Arrhenius (1888) Proposed that molecules need minimum energy to react Activation Energy, !a Very important concept! Svante Arrhenius,1859–1927 (wikipedia.org) Why does this make sense? Side comment: First to predict greenhouse effect Reactants stable due to human-caused CO2 (1896) Products stable What happens in between? Chemistry iiiiiiEiii Lecture #11, p. 10 Activation Energy, !a In between, chemical bonds are: Half broken and half formed TajistEption Less stable with higher energy !a ≡ energy needed to get through this less stable state iii At peak * reactants are in: transition state or activated complex Moving from reactants to products Mechanical analogy: tipping over bottle ∗ ≡ “tipping point” Chemistry Lecture #11, p. 11 Activation Energy, !a Kinetic “barrier” Implications Larger !a ⇒ bigger energetic “hill” to overcome Molecules carry energy into reaction via their motion Molecules have distribution of kinetic energies !k according to the kinetic-molecular theory (Lecture #2) To react, molecules need !# > !a If !a is large, few molecules can react Larger !a ⇒ slower rates! Chemistry Ofmolecules Lecture #11, p. 12 Temperature Dependence of Rate Law Rate = ! [A]"[B]# [C]$ ⋯ !! Molecules Rate constant ! ⇒ depends on T that can Iii react Kinetic Energy Iiii ! " = collisions per unit time ⋅ fraction of collisions properly oriented ⋅ fraction of molecules with ! > !! = ! " # ⋅ exp −)! ⁄*" Iii # ≡ T-independent “frequency factor” Note: # is different than reactant A Chemistry 1 re EE.ie i EE i exp 1 0.36 dependent duetotheexponentialtermexp Ii It justchangesmoreslowlythanexponentialterm Lecture #11, p. 13 temperature É Eatmfier she it Yesitself then late i Arrhenius Equation Very important! ! " = #⋅exp −)%⁄*" Take natural log of both sides: ln ! ln ! = ln # 7 exp −)%⁄*" %! Slope = − & ln ! = ln exp −)%⁄*" + ln # )% 1 ln ! = − + ln # * " 0 $ = %*' + ) (line) 0 13 " Chemistry iii Lecture #11, p. 14 Backto questionabout speedy ants Ants at +10 °C ? !' "' = #⋅exp −)%⁄*"' !& "& = #⋅exp −)%⁄*"& Adobe Stock !' )% )% !' )% 1 1 = exp − + OR ln = − !& *"' *"& !& * "& "' Can calculate rate at a different temperature Also, explains “rule of thumb” for chemical reactions: For typical )%, " ⇒ Reaction rate doubles if ∆" = +10 °C Chemistry Exp 11,0T E Lecture #11, p. 15 Role of ∆"? Note: ∆) did not affect reaction rate above! But ∆) does influence reverse rate: Activation energy for reverse reaction: )%()*)(+) = )%,-(.%(/ − ∆) Elitism For ∆) < 0 : reverse reaction is slower For ∆) > 0 : reverse reaction is faster Chemistry Here forward on is exothermic eco If thenreverse ratewould befasterthanforwardrate I i.ITEEF EfFEF Lecture #11, p. 16 Sowe havenow answered 2 of our 3 Butwhataboutthe meaning of min line Inti Em wilt the rntabt.IE 1 Reaction Mechanisms Describe how a reaction proceeds The order that the bonds break and form Write reaction in terms of elementary reactions A simple step where one “thing” happens A complex multistep reaction can involve several elementary reactions Typically, one bond breaks or forms Each elementary reaction has a molecularity The number of molecules participating Number of molecules involved in elementary reaction: Unimolecular A ⟶ Products one common Bimolecular A + A ⟶ Products two common Termolecular A + A + A ⟶ Products three possible, but less likely Chemistry sainse Thereactantmole.ae Remote.ie omisnule etenialied FEE is pianets Lecture #11, p. 17 It is perhapseasier to understand with a specificexample Example Many reactions that we have discussed actually involve a sequence of elementary reactions Tromlast leiture 2 O3(g) 3 O2(g) Involves 2 elementary steps: UV-B Step 1 O3(g) O2(g) + O(g) O(g) appears as intermediate Step 2 O3(g) + O(g) 2 O2(g) ⇒ formed in step 1 ⇒ consumed in step 2 Sum 2 O3(g) 3 O2(g) ni it Elementary steps must add to give correct overall reaction Chemistry in th imi i iii ii e iai izinti in iii aiiiii.ieiiiiiin it EE i i IE i a Y ii Lecture #11, p. 18 Now we can address question about min Rate Laws for Elementary Reactions (A + *B → ,C +.D Rate = # [A]! B " Lecture 10: “0, 1 are not necessarily equal to (, *, respectively” When are they equal? 0 = ( and 1 = * for elementary reaction OR Given an elementary reaction, the reaction order depends on molecularity Maybe clearer if we make table... Chemistry Lecture #11, p. 19 Elementary Reactions and Their Rate Laws P ≡ Products Molecularity Elementary reaction Rate law Unimolecular A⟶P k [A] Bimolecular A+A⟶P k [ A ]2 Bimolecular A+B⟶P k [A][B] Termolecular A+A+A⟶P k [ A ]3 Termolecular A+A+B ⟶P k [ A ]2 [ B ] Termolecular A+B+C ⟶P k [A][B][C] Termolecular less likely than bimolecular 4-molecule collision even rarer Chemistry Lecture #11, p. 20 Notes Not many possible rate laws for elementary reactions Just those in the table above! Problem: For general reaction, we do not know if it is elementary Multistep reactions are common Leads to rate laws that cannot be written “by inspection” That is, they are not obvious Explains statement (Lecture #10): “ m, n not necessarily equal to !, # ” Chemistry Lecture #11, p. 21 Given a generic rxn whatistheprocess to determinethe rxn mechanism Process to Determine Reaction Mechanism Procedure Ex: Proposed mechanism Measure rates experiment !! Propose mechanism Step 1 2A I+C Elementary reactions Check consistency !" Step 2 I+B A+D Compare measured rate law with that from proposed Sum A+B C+D Overall reaction mechanism reaction With rate constants !! and !" and I ≡ intermediate Chemistry Determining a rxn mechanism is challenging can takinetii.ci is c sineering Lecture #11, p. 22 Rate Law of Proposed Mechanism !! Step 1 2A I+C Rate1 = #! $ A " Can just write !" down because Step 2 I+B A+D AND they are elementary Rate2 = #" $ I $ B reactions Sum A+B C+D What is the overall rate law for the reaction in terms of reactants, !! , and !" ? We must solve system of differential equations to get overall rate Here, we will assume one elementary reaction is rate limiting ⇒ One step is so slow, it sets the overall reaction rate ⇒ Assume Step 1 is rate-limiting reaction: !! ≪ !" Chemistry slow fast Lecture #11, p. 23 Rate Law of Proposed Mechanism !! Step 1 2A I+C slow Rateoverall = #! $ A " !" Step 2 I+B A+D fast Overall rate law for proposed mechanism Next, compare with experimental rate law Sum A+B C+D Consistent? If yes, great! But not proof of mechanism More than one mechanism could be consistent E If no... Propose another mechanism and try again Note: More complicated to solve if Reaction 2 is rate limiting Usually solved by assuming Reaction 1 is in equilibrium (Lecture #12) Chemistry havent But we treatedequilibrium yet Seeoptionalproblem on PS 11 Lecture #11, p. 24 The finaltopic on chemical kinetics is Catalysts Substances that increase reaction rate but are neither produced nor consumed in overall reaction How does catalyst affect rate law? *A + ,B →.C + 0D Rate = # [A]# B $ # 2 = 3⋅exp −9%⁄:2 Catalyst must affect #! How? Increase 3 ⇒ better orient molecules for reaction Can OR Lower 9% ⇒ lower energy of transition state or allow new mechanism Chemistry Eftp.piseartriexpinhettiaFinisarinesdef ti Sotypically catalysts lower Ea Lecture #11, p. 25 Catalysts: Graphically &%'%()*' 9% < 9% Catalyst can lead to multistep reaction, leading to multiple peaks Slowest step has highest peak in curve (i.e., hardest to get over) Chemistry Lecture #11, p. 26 Catalysts: Types Homogeneous: reactants, products, and catalysts are in the same phase ex’s: acid catalysis (protons) enzymes (next page) Heterogeneous: reactants, products, and catalysts are in different phases ex: solid catalysts for gaseous reactants/products Precious metals (Pt, Pd, Rh) catalyze: NO, NO2 ⟶ N2 + O2 “NOx” askcarmechanic.com Chemistry Lecture #11, p. 27 Natural Catalysts: Enzymes Example: Nitrogen fixation Industry vs. Enzymes N2 + 3 H2 ⟶ 2 NH3 Important for synthetic fertilizers Feeds the world! Cyanobacteria use enzyme Massive scale in industry nitrogenase Haber–Bosch Process “Fixes” N2 at 25°C at 1 bar! Consumes 1-2% of world’s energy! Uses heterogeneous catalysts (Fe) 500°C at 300 bar Chemistry Lecture #11, p. 28 “England and all civilized nations, stand in deadly peril.” —William Crookes, 1898 (a) starvation (b) population (c) pollution (d) war Chemistry Lecture #11, p. 29 The Crisis Crookes predicted: many deaths by starvation in 1930s Why? Humanity running out of fertilizer Background Plants need nitrogen to grow Traditional farming: crop rotation animal manure composting Starting in 1840s: fertilizers imported to Europe Chemistry Lecture #11, p. 30 Fertilizers Bird guano 1840s–1870s Chincha Islands, Peru Chilean saltpeter NaNO3 1860s–1930s Atacama Desert, Chile Crookes predicted supply would end Challenged chemists to produce synthetic fertilizers to avoid mass starvation Haber and Bosch addressed this challenge Chemistry Lecture #11, p. 31 Fritz Haber and Carl Bosch N2 + 3H2 2NH3 Fritz Haber Karlsruhe Institute of Technology 1909 Fixed nitrogen with Os catalyst Yield: 125 mL of NH3 per hour Carl Bosch BASF 1910–1914 Scaled up process Fritz Haber,1868–1934 Carl Bosch,1874–1940 (wikipedia.org) (wikipedia.org) Identified cheaper Fe catalysts Solved huge technical challenges Crazy fact: Developed high-pressure reactors Half of the nitrogen in our 20 tons per day bodies comes from the Haber–Bosch process In 2018, 230 million tons produced Haber and Bosch won the Nobel Prize in Chemistry in 1918 and 1931 Chemistry Lecture #11, p. 32 Synthetic Fertilizers Ammonium sulfate 3NH3 + H2SO4 (NH4)2SO4 First synthetic fertilizer made by Bosch / BASF Lower nitrogen content (21%) than other fertilizers Ammonium nitrate NH3 + HNO3 NH4NO3 Higher nitrogen content (34%) But explosive! Nitric acid NH3 + 2O2 H2O + HNO3 Made by the Ostwald process from ammonia Main raw material for most common synthetic fertilizers But also used to make explosives, e.g., TNT (trinitrotoluene) Chemistry Lecture #11, p. 33 Oppau Plant (Ludwigshafen, DE) Fertilizer explosion at BASF Plant 1921 500–600 people killed Another 2000 people injured Cause? Common to mix NH4NO3 with (NH4)2SO4 (1:1) Adobe Stock Tests showed this mixtre not explosive below 60% nitrate But NH4NO3 is strongly hygroscopic (absorbs water) Becomes “plaster like” substance during storage BASF workers needed pickaxes to break it up before shipping Also, used small charges of dynamite! Adobe Stock Believed that fertilizers were not well mixed Pocket of NH4NO3 exploded when dynamite detonated Source: wikipedia.org Chemistry Lecture #11, p. 34 Photograph from Popular Mechanics Lecture #11, p. 35 But Why Is High Pressure Needed for NH3? Haber–Bosch process N2 + 3H2 ⇌ 2NH3 Reaction is exothermic at room temperature But kinetics makes it SLOW Large activation energy Catalysts lower Ea But they still require high T to get reasonable rates Why do we need to increase pressure? Need knowledge of equilibria to answer this question Next time! Chemistry Lecture #11, p. 36 What We Learned What role does temperature play in kinetics? " # = % exp(−-$ ⁄.#) Arrhenius equation -$ = Activation energy Given two spontaneous reactions, why can one be fast and one slow? -% can be large or small Independent of whether Δ- is exothermic or endothermic Given: Rate = "[A]! [B]" [C]# ⋯ , what do 8, 9, : mean? Rate laws complicated for multistep reactions Defined molecularity for elementary reactions Rate laws easily written down for elementary reactions Chemistry