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- 1. Chemical Kinetics Reaction Mechanisms,Phases and the Movement of Molecules
- 2. PHASE DIAGRAMS (N.AND I.) ■New and Improved! ■Phase Change Diagrams only account for temperature ■New Phase Diagrams account for Pressure, as well. ■Terms: ■Critical Point ■Triple Point *This change comes as a result of our study of Gas Laws!
- 3. PHASE DIAGRAMS Temperature → Pressure → A B C D Solid Liquid Vapor Thelarge white dot (T) represents the triple point. The triple point is the equilibrium* point of all three phases– solid, liquid, gas. Point C – The Critical Point– at temperatures higher than this, there can only the gas phase T
- 4. MORE TRIPLE POINT ■Alongany line, the two phases represented are at equilibrium. Temperature → Pressure → A B C D Solid Liquid Vapor
- 5. CHEMICAL LINE UP Temperature→ Pressure → A B C D Solid Liquid Vapor Each Line has a “meaning”: Tp -A: Solid-Liquid Eq. Tp -B: Solid-Gas Eq. Tp -C: Liquid-Solid Eq. Tp -D: Supercooled H2O
- 6. EQUILIBRIUM ■The state inwhich two or more processes are occurring at the same time: ■Phase Changes ■As one molecule melts, another freezes ■As one molecule boils, another condenses ■Triple Point: ■ALL of the phase changes occur at the SAME time, at Equilibrium
- 7. RELATIONSHIPS (FOR WATER) ■1atm: ■BP = 100 C ■FP = 0.016 C (273.16K) ■0.0063 atm: ■BP = 0.0098 C ■FP = 0.0098 C Temperature → Pressure → A B C D Solid Liquid V a p o r 1 atm 100 C 0.0063 atm 0.0098 C 0 C High-Density Ice 1000 atm
- 8. SOME IDEAS TOCONSIDER ■Adding pressure causes a shift to the more dense phase. ■Adding temp. causes a shift to the less dense phase ■“Gibbs Phase Rule”
- 9. WATER’S UNIQUE PROPERTY ■Whatis the most dense phase of water? Gibbs’ Phase Rule explains why ice skates work the way they do!!! The blade of a skate provides pressure to the ice, which should be near 0 C This causes a shift to the more dense phase. This provides a film of water which allows the skate to slide!
- 10. OTHER PHASE CHANGEGRAPHS ■Carbon Dioxide ■What can this tell us about CO2 ? Temperature → Pressure → Solid Liquid Vapor 1 atm 0 C
- 11. THERMODYNAMICS ■The study ofheat/energy as it moves through a system ■Phase changes: Melting = Freezing? ■Yes-- in one, energy is being added (melting); in the other, it’s being taken away (freezing) ■Calculated using Calorimetry ■Uses masses and temperature changes to determine energy flow
- 12. CALORIMETRY ■Energy can becalculated using the formula: Q = m x c x ∆T Where: Q = energy (joules, kilojoules, or calories) m = mass in g c = Specific Heat (more) ∆T = Change in Temperature (K or o C)
- 13. SPECIFIC HEAT ■Term givenfor the amount of energy needed to raise 1 g of a substance a given temperature: ■For water, c = 1 cal / g o C ■Also c = 4.18 j / g o C ■From which we can deduce: ■1 cal = 4.18 j
- 14. FOR ICE? STEAM? ■Ice’sSpecific Heat: ■2.09 j/g o C ■Steam’s Specific Heat: ■1.84 j/g o C
- 15. ENERGY OF APHASE CHANGE ■Heat of Fusion: The amount of heat required to MELT a solid; ■Also, the heat given off by freezing a liquid ■For H2 O: 333 j / g ■Heat of Vaporization: The amount of heat needed to BOIL a liquid; ■Also, the heat given off by condensing a gas ■For H2 O: 2230 j / g
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- 17. WE KNOW: ■Chemical Bondsform due to electron movement to a state of lower energy. ■NOW: we begin learning about the forces that drive reactions.
- 18. CHEMICAL KINETICS HASFOUR PRIMARY PARTS: ■Kinetic Molecular Theory ■Reaction Mechanisms ■Reaction Rates ■Phases and Phase Changes (already started!)
- 19. KINETIC MOLECULAR THEORY ■Saysthat all particles are moving at all times ■Also called collision theory ■“Proven” (sorta) through our studies of gases.
- 20. POSTULATES OF KMT ■Allmolecules are extremely small in mass. ■All molecules are in constant motion. ■There are so many particles, even in a small sample, that we can use statistics (mean, median, mode, etc.) to make generalizations about particle movement.
- 21. MORE POSTULATES OFKMT ■Collisions are perfectly elastic (all energy remains kinetic) ■The kinetic energy of the system is directly related to the temperature (higher temp. = more E!) ■… So what?
- 22. WHAT IT MEANS: ■Heatis a function of the energy of motion of particles ■When something gets hot, its particles are moving faster than before. ■Everything, whether you like it or not, is moving (at a molecular level). ■Chemical reactions and interactions are all a result of chemicals colliding with one another… which leads us to…
- 23. REACTION MECHANISMS ■Chemical KineticTheory says that NO reaction can occur without the effective collision of two or more molecules. ■Contact ■Electron Clouds Must “touch” ■Proper Alignment ■Molecules Must Align correctly in order to react ■High Enough Energy to Begin Reacting ■Activation E! ■This WHOLE concept is called: COLLISION THEORY
- 24. REACTION MECHANISMS ■Are definedas the underlying action in a chemical reaction ■2C8 H18 + 25 O2 16CO2 + 18 H2 O ■What are the odds that both Octanes and ALL 25 O2 collide at the same time? ■Problem: this is a highly exothermic reaction, rapid reaction-- How can this be?
- 25. ANSWER: ACTIVATED COMPLEXES ■Moleculesthat exist ONLY DURING a reaction ■Parts of molecules that have either: ■Broken apart from reactants ■Formed together to make products ■CO and OH (0) exist during this reaction, among others. ■Act. Complexes Video
- 26. SO WHAT? ■If weagree that kinetic theory is true, and… ■Activated Complexes exist, then… ■We can control how fast some chemical reactions occur.
- 27. RATE DETERMINING STEP ■ActivatedComplexes say that reactions occur in steps. ■Some steps are slower than others, which limit how fast a reaction can occur. ■Video here
- 28. REACTION RATES ■Demonstration Reaction: ■Zn+ NH4 NO3 ZnO + N2 + H2 O ■Slow Reaction was Changed to a quick one! ■Reaction Rate can be Affected! ■Slow Rate: Heat given off ■Fast Rate: EXPLOSION!! ■Watch this guy…
- 29. FACTORS AFFECTING REACTION RATE ■Thereare FOUR you need to know: ■Nature of the Reactants ■Temperature of the System ■Concentration of Reactants and/or Products ■Catalysts
- 30. NATURE OF REACTANTS ■Dependson what types of things are being reacted. ■Primarily a measure of the reactivity of the involved chemicals. ■The more reactive each participant, the more quickly the reaction occurs. ■Reactivity increases as E! increases ■Affects the number of effective collisions ■Includes Phase Differences: More later
- 31. TEMPERATURE ■In a vastmajority of Reactions, a higher temp. Results in a faster reaction ■An increase in effective collisions from having higher E! available ■To effectively deliver temperature: ■Add Heat: Fire, Sun, etc. ■Add H2 O: Has a high Specific Heat
- 32. CONCENTRATION AND REACTION RATE ■Anincrease in the Conc. of Reactants= Faster Reaction ■An increase in the Conc. of Products= Slower Reaction ■Think of a Jr. High dance…
- 33. CONCENTRATION AND REACTION RATE(CONTINUED) ■Surface Area Increase = An Increase in Conc. ■More Molecules Available to Collide ■Phases have variable surface area*: ■Liquid-Liquid ■Liquid-Gas (bubbling) ■Gas-Gas ■Liquid-Solid ■Gas-Solid ■Solid-Solid
- 34. CATALYSTS AND REACTIONRATE ■Catalysts always increase reaction rate. ■Defined: ■Chemicals that act to speed up a reaction without being used up themselves in the reaction. ■We do not fully understand the function of Catalysts ■Catalysts MAY break down during a reaction-- but will always re-form by the end of it.
- 35. REVERSIBLE REACTIONS ■Chemical Reactionsthat can occur in either direction in nature. ■Under certain circumstances, reactants become products ■Under different circumstances, products turn back into reactants ■Can (and will) achieve Equilibrium A + B C + D Vid Here
- 36. LECHATELIER’S PRINCIPLE ■LeChatelier saidthat a reversible reaction can be affected by adding a stress. ■Stresses include: ■Concentration of Chemicals ■Temperature ■Addition of a Catalyst
- 37. LECHATELIER’S RESULTS ■Concentration ofChemicals ■More reactant added ■Shifts reaction to the products side ■More product added ■Shifts reaction to the reactants side ■Temperature ■Increasing temp causes: ■Exothermic reactions to move toward reactants ■Endothermic reactions to move towards products
- 38. LAST FOR LECHATELIER ■ACatalyst will ALWAYS move a reaction towards the products side ■Catalysts always act to speed up reactions
- 39. CONTROLLING REACTIONS INTHE FIRST PLACE? ■We can control how fast reactions take place. ■We know that collisions are the root of all chemical reactions. ■We know that particles are in constant motion. ■Can we make a reaction occur? Stop a reaction from occurring?
- 40. CONTROLLING CHEMICAL REACTIONS ■Itis possible to prevent reactions from occurring, and it is possible to make reactions occur that normally wouldn’t: ■Activation Energy ■Enthalpy and Entropy
- 41. ACTIVATION ENERGY ■The EnergyRequired to START a Chemical Reaction ■E! act (abbr.) ■Examples: Gasoline; Ba(OH)2 + NH4 Cl ■Can be shown on an potential E! Graph ■Without enough E!, no reaction will occur.
- 42. REACTION RATE FORCES ■Additionally,nature controls whether or not reactions through two forces: ■Enthalpy, or ΔH ■The total amount of energy in an object/chemical ■Entropy, or ΔS ■The amount of disorder in an object or chemical
- 43. WHAT NATURE WANTS… ■Naturewants: ■Energy to be given off to the universe ■An decrease in ∆H (- ∆H) ■Entropy to increase ■An increase in ∆S (+ ∆S)
- 44. PREDICTING REACTION OCCURRENCE ■Giventhe ΔH and ΔS of a system, you can predict if a reaction will occur. ■If ΔH is negative and: ΔS is Positive, the reaction will occur ΔS is Negative, the reaction will occur if: ΔH has a greater absolute value than ΔS ■If ΔS is positive and: ΔH is positive, the reaction will occur if: ΔS has the larger absolute value
- 45. MORE PREDICITING ■If ΔHis positive and ΔS is negative the reaction will never occur. ■Although not “the whole story,” phases are related to ΔH and ΔS.
- 46. PHASES AND FORCES ■Recall: Solids:Very Low ΔS, Very Low ΔH Orderly, low E! Liquids: Low ΔS, Low H Semi-Orderly, moderate E! Gases: High ΔS, High ΔH Disorderly, high E! Plasma: Very High ΔS, Very High ΔH Ridiculous disorder, ridiculous E!
- 47. INTERMOLECULAR FORCES ■Forces thatact between 2 separate molecules. ■Phases are directly related to the strength of the Intermolecular Forces in a system. ■Solids- Strong IMF ■Plasma- Weak IMF
- 48. INTERMOLECULAR FORCES ■There are4 Primary Intermolecular Forces: ■London Dispersion Forces ■Dipole Forces ■Hydrogen Bonds ■Gravity
- 49. REMEMBER THIS? ■To aidein your understanding of IMF, check out this blast from the past: 0 2 0.9 3 1.8 1
- 50. LONDON DISPERSION FORCES ■Occurprimarily in Non-Polar Covalent Cpds., but do occur in all substances. ■Caused by the random “dispersion” of electrons within a compound. ■As e- move, they create weak positive and negative areas that are therefore attracted to each other. ■A temporary electro-magnetic force
- 51. DIPOLE FORCES ■Occur inpolar molecules and ionic compounds only. ■The (∂+) or (+) portion is attracted, quite strongly, to the (∂-) or (-) part. ■*This force is stronger as substances become more ionic in character.* ■A permanent, though variable (depending on the differential), electro-magnetic force
- 52. HYDROGEN BONDS ■Occur onlyin Polar Molecules that contain a hydrogen and highly electronegative element. ■F, O, Cl <in decreasing order> ■A permanent, average strength (because of the partial differential) electro-magnetic force H O H H H O Hydrogen Bond
- 53. GRAVITY ■The pull oftwo massive bodies towards each other ■For their size, atoms are quite massive ■Extremely weak force in spite of this ■Gravity is directly tied to mass; size, at the atomic level, is generally irrelevant
- 54. IMF AFFECT PHYSICALPROPERTIES ■According to periodic trends, Water should boil at a much lower temperature: Atomic Mass → Temp ●H2S ●H2Se ●H2Te ●(H2Po) H2O ●H2S ●H2Se ●H2Te ●(H2Po) ● H2O ● H2O boils at 100 C; H2Te boils at -2 C! Trends Say: Real Life Proves: