* Your assessment is very important for improving the workof artificial intelligence, which forms the content of this project
Download + NO 2
Stability constants of complexes wikipedia , lookup
Electrochemistry wikipedia , lookup
Equilibrium chemistry wikipedia , lookup
Woodward–Hoffmann rules wikipedia , lookup
Determination of equilibrium constants wikipedia , lookup
Photoredox catalysis wikipedia , lookup
Supramolecular catalysis wikipedia , lookup
Marcus theory wikipedia , lookup
Ultraviolet–visible spectroscopy wikipedia , lookup
Chemical thermodynamics wikipedia , lookup
Chemical equilibrium wikipedia , lookup
Industrial catalysts wikipedia , lookup
Physical organic chemistry wikipedia , lookup
Enzyme catalysis wikipedia , lookup
George S. Hammond wikipedia , lookup
Reaction progress kinetic analysis wikipedia , lookup
KINETICS CHEMISTRY— INTRODUCTION KINETICS • STUDY OF REACTION RATES • HOW FAST DOES IT HAPPEN? WHAT VARIABLES INFLUENCE THE RATE? WHAT IS THE PATH THE REACTION TAKES TO CONVERT REACTANTS TO PRODUCTS? • RATE • CHANGE IN AN AMOUNT OVER A PERIOD OF TIME • EX. DISTANCE TRAVELED, SPACE TRAVEL REACTION RATES • CHANGE IN THE CONCENTRATION OF A CHEMICAL COMPOUND IN THE REACTION OVER A PERIOD OF TIME • FOCUS ON ONE REACTANT OR ONE PRODUCT IN THE REACTION • WANT TO KNOW RATE OF DISAPPEARANCE FOR REACTANT/RATE OF APPEARANCE FOR A PRODUCT • RATE = ∆X ∆T UNITS = M/TIME EXAMPLE 1: A + B → D + E • A) RATE OF DISAPPEARANCE FOR A • RATEA = -∆[A] ∆T • B) RATE OF APPEARANCE FOR D • RATED = ∆[D] ∆T EXAMPLE 2: 2AB → A2 + B2 Time = 0 seconds 15.0 M 0M 0M Time= 60 seconds 5.0 M 5.00M 5.00M • DETERMINE THE RATE OF DISAPPEARANCE OF AB • DETERMINE THE RATE OF APPEARANCE OF A2 AND B2 REACTION RATE AT ANY MOMENT IN TIME • SLOPE OF LINE TANGENT TO THE REACTION CURVE COLLISION THEORY • REACTANT MOLECULES MUST COLLIDE TO PRODUCE A CHEMICAL REACTION • THE CONCENTRATIONS OF REACTANTS AFFECT THE # OF COLLISIONS AMONG REACTANTS • FOR REACTIONS OCCURRING IN ONE STEP—RATE OF REACTION IS PROPORTIONAL TO PRODUCT OF REACTANT CONCENTRATIONS • RATE = K[A] [B] • RATE OF ANY REACTION STEP DEPENDENT ON COLLISION FREQUENCY VARIABLES AFFECTING REACTION RATE 1) COLLISION RATES BETWEEN REACTANTS 2) % OF COLLISIONS WITH REACTANTS ARRANGED IN PROPER ORIENTATION TO PRODUCE REACTION. 3) % OF COLLISIONS WITH ENERGY ENERGY (ACTIVATION ENERGY) TO PRODUCE REACTION. WHEN DO COLLISION RATES INCREASE? • INCREASE IN CONCENTRATIONS OF REACTANTS • TEMPERATURE INCREASES •WHY? MOST COLLISIONS DO NOT RESULT IN A CHEMICAL REACTION! • SMALL PERCENTAGE OF COLLISIONS ACTUALLY CONVERT REACTANTS TO PRODUCTS. WHY? 1) MOLECULAR ORIENTATION • RANDOM ORIENTATION • NOT ALL COLLISIONS HAVE CORRECT ORIENTATION 2) MOLECULAR ENERGY AT COLLISION • MOLECULES HAVE DIFFERENT KINETIC ENERGIES • COLLISION ENERGY IS ENERGY SOURCE TO GET A REACTION STARTED ACTIVATION ENERGY (EA) • THE AMOUNT OF COLLISION ENERGY NEEDED TO OVERCOME EA SO THE REACTION CAN OCCUR • AMOUNT OF ENERGY NEEDED FOR A CHEMICAL REACTION TO HAPPEN, ENERGY NEEDED TO CONVERT REACTANTS TO PRODUCTS. ACTIVATION ENERGY--ENDOTHERMIC WHEN WILL REACTIONS OCCUR? 1) MUST HAVE A COLLISION 2) COLLISION MUST HAPPEN WITH THE CORRECT MOLECULAR ORIENTATION TO GENERATE A REACTION 3) COLLISION ENERGY ≥ EA ARRHENIUS EQUATION • RATE CONSTANT AND REACTION RATE ARE TEMPERATURE DEPENDENT. • ENABLES THE ACTIVATION ENERGY FOR A REACTION TO BE DETERMINED BASED ON THE RELATIONSHIP BETWEEN REACTION RATE AND TEMPERATURE. ARRHENIUS EQUATION •LNK = -EA ( 1/T ) + LNA R • • • • • K = RATE CONSTANT EA = ACTIVATION ENERGY (J) R = 8.314 J/MOLK T = KELVIN Z = PROPORTIONALITY CONSTANT, CHANGES BASED ON REACTION ARRHENIUS EQUATION • DIFFERENT FORM OF EQUATION CAN BE USED TO OBSERVE HOW TEMPERATURE CHANGES AFFECT THE RATE CONSTANT (K) •LN (K1/K2) = EA (1/T2 – 1/T1) R EXAMPLE 1 • CALCULATE ACTIVATION ENERGY (EA) FOR HI DECOMPOSITION WITH THE FOLLOWING DATA. Temperature (K) Rate Constant (M/s) 573 2.91 x 10-6 673 8.38 x 10-4 773 7.65 x 10-2 COLLISION THEORY • IN ORDER FOR TWO PARTICLES TO REACT CHEMICALLY, THEY MUST COLLIDE. NOT ONLY MUST THEY COLLIDE, BUT IT MUST BE AN “EFFECTIVE COLLISION.” THAT IS, THEY MUST HAVE THE CORRECT AMOUNT OF ENERGY AND COLLIDE WITH THE PROPER ORIENTATION IN SPACE. • ANY FACTOR WHICH INCREASES THE LIKELIHOOD THAT THEY WILL COLLIDE WILL INCREASE THE RATE OF THE CHEMICAL REACTION. Presented by Mark Langella, PWISTA.com FACTORS WHICH AFFECT THE RATE OF A CHEMICAL REACTION (BONDS MUST BREAK) • SURFACE AREA/ CONTACT AREA (OPPORTUNITY FOR COLLISIONS) • CONCENTRATION ( INCREASE FREQUENCY) • TEMPERATURE ( INCREASE FREQUENCY) • CATALYST ( EFFECTIVE COLLISIONS) • NATURE OF REACTANTS ( EFFECTIVE COLLISIONS) Presented by Mark Langella, PWISTA.com WHAT CAN INFLUENCE REACTION RATES? 1) TEMPERATURE 2) CONCENTRATION 3) CATALYST 4) SURFACE AREA 5) VOLUME/PRESSURE 6) REACTANT PROPERTIES 12_300 T1 T2 > T1 T2 0 0 Ea Energy Plot showing the number of collisions with a particular energy at T1& T2, where T2 > T1 -- Boltzman Distribution. Presented by Mark Langella, PWISTA.com RATE LAW (CONT.) • A+B→ C+D • RATE = K [A]M[B]N • RATE = RATE OF DISAPPEARANCE OF REACTANTS • K = RATE CONSTANT, SPECIFIC TO REACTIONS AND TEMPERATURE • M = REACTION ORDER IN TERMS OF A • N = REACTION ORDER IN TERMS OF B • M + N = OVERALL REACTION ORDER REACTION ORDER • INDICATES HOW CONCENTRATION CHANGES AFFECT CHANGES IN THE REACTION RATE • ORDERS: 0, 1, 2 • OVERALL ORDER OF REACTION = Σ INDIVIDUAL ORDERS OF EACH REACTANT • ORDER OF A REACTION IN TERMS OF A REACTANT ≠ REACTANT’S COEFFICIENT IN CHEMICAL EQUATION REACTION ORDERS (CONT.) • ZERO-ORDER REACTION • RATE NOT DEPENDENT ON REACTANT’S CONCENTRATION • CONSTANT REACTION RATE • FIRST-ORDER REACTION • CONCENTRATE AFFECTS REACTION RATE • EXAMPLE: DOUBLE CONCENTRATION, DOUBLE THE RATE. • SECOND-ORDER REACTION • CONCENTRATION AFFECTS REACTION RATE • EXAMPLE: DOUBLE CONCENTRATION, QUADRUPLE THE RATE RATE CONSTANT (K) UNITS Reaction Order Basic Formula Units 0 Rate = k Ms-1 1 Rate = k [A] s-1 2 Rate = k [A]2 M-1s-1 3 Rate = k [A]3 M-2s-1 EXAMPLE 3: 2NO(G) + O2(G) 2NO2(G) BASED ON THE REACTION’S RATE LAW OF RATE = K(NO)2 (O2) CLASSIFY THIS REACTION’S ORDER. EXAMPLE 4: • DETERMINE THE RATE LAW, REACTION ORDER, AND RATE CONSTANT (K) FOR THE FOLLOWING REACTION AT A SPECIFIC TEMPERATURE---• 2NO(G) + 2H2(G) N2(G) + 2H2O(G) Experiment [NO]initial [H2]initial Rate initial 1 0.20M 0.30M 0.0900 M/s 2 0.10M 0.30M 0.0225 M/s 3 0.10M 0.20M 0.0150 M/s EXAMPLE 5: • DETERMINE THE RATE LAW FOR THE FOLLOWING REACTION---• NH4+(AQ) + NO2-(AQ) N2(G) + 2H2O(L) Experiment [NH4+]initial [NO2-]initial Rate initial 1 5 x 10-2 M 2 x 10-2 M 2.70 x 10-7 M/s 2 5 x 10-2 M 4 x 10-2 M 5.40 x 10-7 M/s 3 1 x 10-1 M 2 x 10-2 M 5.40 x 10-7 M/s INTEGRATED RATE LAW • ENABLES THE DETERMINATION A REACTANT’S CONCENTRATION AT ANY MOMENT IN TIME • ENABLES THE DETERMINATION OF THE TIME IT TAKES TO REACH A CERTAIN REACTANT CONCENTRATION • ENABLES THE DETERMINATION OF THE RATE CONSTANT OR REACTION ORDER 1ST ORDER INTEGRATED RATE LAW • ONLY USED WITH 1ST ORDER REACTIONS • FOCUS ON INITIAL CONCENTRATION AND ΔC FOR ONE REACTANT • INITIAL CONCENTRATION OF REACTANT KNOWN---- CAN DETERMINE REACTANT CONCENTRATION AT ANY TIME • INITIAL AND FINAL REACTANT CONCENTRATIONS KNOWN---CAN DETERMINE RATE CONSTANT 1ST ORDER INTEGRATED RATE LAW • RATE = -Δ[A] = K [A] ΔT -TAKE EQUATION AND INTEGRATE WITH CALCULUS TO GET…. • LN[A]T – LN[A]0 = - KT • [A]0 = INITIAL CONCENTRATION (T = 0) • [A]T = CONCENTRATION AFTER A PERIOD OF TIME EXAMPLE 1: A B + 2D • USING THE DATA PROVIDED FOR A 1ST ORDER REACTION, DETERMINE THE RATE CONSTANT AND [A] AT TIME = 5.0 X 102S. Time (s) [A] (M) 0 0.020 5.0 x 10 0.017 1.0 x 102 0.014 1.5 x 102 0.012 2.0 x 102 0.010 HALF-LIFE • RADIOACTIVE DECAY IS A 1ST ORDER PROCESS • HALF-LIFE (T1/2)— • TIME IT TAKES FOR HALF OF A CHEMICAL COMPOUND TO DECAY OR TURN INTO PRODUCTS • FOCUS ON REACTANT • CONSTANT, NOT DEPENDENT ON [ ] • RATE CHANGES WITH TEMPERATURE SO HALF-LIFE VARIES BASED ON TEMPERATURE EXAMPLE 2: • FIND THE HALF-LIFE FOR THE FOLLOWING REACTION WITH A REATE CONSTANT (K) OF 1.70 X 10-3 S-1 2ND ORDER INTEGRATED RATE LAW • USED ONLY FOR SECOND ORDER REACTIONS • FOCUS ON INITIAL CONCENTRATION AND ΔC FOR ONE REACTANT WITH REACTION 2ND ORDER WITH RESPECT TO IT. • INITIAL CONCENTRATION OF REACTANT KNOWN---- CAN DETERMINE REACTANT CONCENTRATION AT ANY TIME • INITIAL AND FINAL REACTANT CONCENTRATIONS KNOWN---CAN DETERMINE RATE CONSTANT 2ND ORDER INTEGRATED RATE LAW • RATE = -Δ[A] = K [A]2 ΔT -TAKE EQUATION AND INTEGRATE WITH CALCULUS TO GET…. • 1 [A]T __ - 1__ = KT [A]0 • [A]0 = INITIAL CONCENTRATION (T = 0) • [A]T = CONCENTRATION AFTER A PERIOD OF TIME EXAMPLE 3: 2NO2(G) 2NO(G) + O2(G) • USING THE DATA PROVIDED, FIND THE RATE CONSTANT IF THE RATE LAW = K[NO2]2. Time (s) [NO2] 0.0 0.070 1.0 x 102 0.0150 2.0 x 102 0.0082 3.0 x 102 0.0057 EXAMPLE 4: • NO2 REACTS TO FORM NO AND O2 BY SECOND-ORDER KINETICS WITH A RATE CONSTANT = 32.6 L/MOLMIN. WHAT IS THE [NO2] AFTER 1 MINUTE IF THE INITIAL [NO2] = 0.15M? REACTION MECHANISM • PATHWAY OR SERIES OF STEPS THROUGH WHICH REACTANTS CONVERTED TO PRODUCTS • NOT ALL STEPS PROCEED AT THE SAME RATE • SUM OF THE STEPS = OVERALL REACTION • RATE OF ANY STEP IS DIRECTLY PROPORTIONAL TO THE REACTANT CONCENTRATIONS IN THE STEP. EX. 1: 2NO(G) + O2(G)2NO2(G) • STEP 1: 2NO N2O2 • STEP 2: N2O2 + O2 2NO2 __________________________ • 2NO + O2 2NO2 RATE-LIMITING STEP • STEP IN A REACTION MECHANISM THAT “LIMITS” HOW FAST PRODUCTS ARE FORMED. • “LIMITS” THE RATE OF REACTANTS CONVERTING TO PRODUCTS INTERMEDIATES • CHEMICAL COMPOUNDS FORMED AND CONSUMED IN A REACTION MECHANISM • APPEAR ON BOTH SIDES OF CHEMICAL EQUATION • TRANSITION COMPOUNDS BETWEEN REACTANTS AND PRODUCTS • UNSTABLE, ONLY EXIST A SHORT TIME DETERMINING THE RATE LAW FOR A REACTION USING A REACTION MECHANISM 1)OVERALL REACTION RATE LAW • ONLY DETERMINE EXPERIMENTALLY 2)OVERALL REACTION RATE LAW CAN BE APPLIED TO FIND THE STEPS IN A REACTION MECHANISM • RATE LAW INDICATES SLOWEST STEP IN THE MECHANISM • DETERMINE THE FAST STEPS REMAINING IN THE MECHANISM RATE LAWS FOR ELEMENTARY STEPS IN MECHANISM • PREDICTABLE • CAN USE COEFFICIENTS ONLY FOR THESE STEPS • AA + BB DD + EE RATE = K[A]A [B]B • OVERALL REACTION RATE = RATE OF SLOWEST STEP EX. 2: NO2 + CO NO + CO2 • STEP 1: 2NO2 NO + NO3 (SLOW) • STEP 2: NO3 + CO NO2 + CO2 (FAST) DETERMINE THE RATE LAW FOR EACH STEP. EX.3 BASED ON THE FOLLOWING REACTION MECHANISM…. • F2 + NO2 NO2F + F (SLOW) • NO2 + F NO2F (FAST) a)WRITE THE OVERALL REACTION. b)DETERMINE THE RATE LAW FOR EACH STEP c) WHAT IS THE RATE LAW FOR THE OVERALL REACTION? CATALYSTS • CHEMICAL COMPOUNDS (ATOMS, MOLECULES, IONS) THAT INCREASE ONLY THE REACTION RATE. • INCREASES BOTH FORWARD AND REVERSE RATES FOR A CHEMICAL REACTION • NOT CONSUMED IN THE REACTION, NOT ALTERED • PRESENT AT THE START AND END OF REACTION • HAVE NO EFFECT ON EQUILIBRIUM CONSTANTS (K), ΔH, OR ΔS CATALYSTS (CONT.) • CHANGE THE PATH A CHEMICAL REACTION TAKES TO GET REACTANTS TO PRODUCTS • LOWERS ACTIVATION ENERGY (EA) OR ENERGY NEEDED FOR THE REACTION TO START • ONLY SMALL AMOUNT OF THE COMPOUND NEEDED TO INCREASE THE RATE FOR A REACTION WITH A LOT OF REACTANT • RECYCLE AND REUSE CATALYST REACTION PATHWAY CATALYSTS (CONT.) • EXAMPLE: H2O2 DECOMPOSITION Catalyst Ea (kJ/mol) Reaction Rate None 75.3 1 Catalase 8 6.3 x 1011