r eaction r ates. r ates of c hemical c hange any measurable change in an activity expressed as a...
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REACTION RATES
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RATES OF CHEMICAL CHANGE Any measurable change in an activity
expressed as a function of time is a rate. Chemical Kinetics – the study of reaction
rates and the factors that affect them. CHEMICAL RATES MUST BE DETERMINED
EXPERIMENTALLY! Measureable changes for chemical reactions
are things like change in mass, volume of an evolved gas, change in the acidity of the solution, and change in color (concentration).
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MEASURABLE CHANGES - EXAMPLE
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RATES Concentrations must be measured often Reaction rates decrease with time (pendulum) RATES ARE POSITIVE, EVEN IF SUBSTANCE IS BEING
USED UP Rates for equilibrium reactions become constant
(forward rate = reverse rate) Rates for non- equilibrium reactions decrease to
zero (reaction goes to completion) The rate of change for each reactant or product
must be divided by its coefficient in the balanced chemical equation.
Rate = -1/a x Δ[reactant]/Δt = 1/b x Δ[product]/Δt a and b represent the coefficients of the balanced reaction The negative in front of 1/a is to keep the rate positive
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MISCELLANEOUS RATE INFORMATION Ionic bonds are usually broken quickly
Covalent bonds are usually broken slowly
Homogeneous – all reactants are in the same phase
Heterogeneous – reactants are different phases
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Rate = -1/a x Δ[reactant]/Δt = 1/b x Δ[product]/Δt2Br -(aq) + H2O2(aq) + 2H30+(aq) Br2(aq) + 4H2O(l)
2N2O5(s) 4NO2(g) + O2(g)
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Time t (s) [H3O+] [Br2]
0 0.0500 085 0.0298 0.010195 0.0280 0.0110105 0.0263 0.0118
2Br -(aq) + H2O2(aq) + 2H30+(aq) Br2(aq) + 4H2O(l)
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GRAPHICAL REPRESENTATIONS 1
H2(g) + I2(g) 2HI(g)
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GRAPHICAL - RATE DECREASES WITH TIME
2N2O5(s) 4NO2(g) + O2(g)
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CHEMICAL KINETICS – RATES AND MECHANISMS OF CHEMICAL REACTIONS Mechanisms – a step by step description of a
chemical reaction
Collision theory – particles must collide and must have proper angles and energy
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FACTORS AFFECTING REACTION RATE Nature of reactants
Temperature of the system
Concentration of the reactants
Use of catalysts
Surface area
Pressure (partial pressure) is the same as concentration if reactants are all gases (Avogadro’s principle)
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INTERMEDIATE PRODUCTS (ACTIVATED COMPLEX)
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MECHANISMS All reaction mechanisms have a rate
determining step, a slow step.
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TEMPERATURE Each 100C (10K) increase generally doubles
the rate, temperature increases the effectiveness of collisions
Higher temperature, more kinetic energy, more likely to reach the activated complex
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COLLISIONS MUST HAVE PROPER ORIENTATION
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CONCENTRATION Increase in concentration usually increases
the rate except for some homogeneous phase reactions or if that reactant is not involved in the mechanism
Higher concentration – greater frequency of collisions, more reactants with sufficient energy to reach the activated complex
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CATALYSTS Catalyst – speeds up a reaction without being
altered itself, provides a new mechanism, lowers Eact
By the way, an inhibitor slows down a reaction (preservatives )
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ENZYMES ARE SPECIFIC ORGANIC CATALYSTS
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SURFACE AREA With greater surface area, reactants are
more likely to collide! Sugar cube vs. granular sugar
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PARTIAL PRESSURES FOR GASES If the partial pressure of a gas increases, the
concentration increases (more gas, more pressure) and therefore the rate will also increases for reactions in which all reactants are gases.
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HOW CAN REACTION RATE BE DESCRIBED?
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RATE LAWS The rate law describes the way in which reactant
concentration affects reaction rates. General rate law rate ∝ [A]n = k [A]n The exponent (n) is the order of reaction with respect to A and must be determined experimentally. EXPONENTS ARE NOT simply THE COEFFICIENTS OF BALANCED REACTIONS AS WE DO WITH EQUILIBRIUM. (although they could be the same) 1ST order, n = 1, double the concentration, double the rate. 2nd order, n = 2, double the concentration, quadruple the rate. ETC. “0 “order, the reactant is not involved in the reaction rate. Simply “used up” as the reaction proceeds.
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RATE LAW FOR MULTIPLE REACTANTSrate = k [A]n [B]mn is the order in [A]m is the order in [B]overall order = m + nUse experimental data to compare change in rate change in concentration.
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DETERMINING RATE LAWS FROM DATA EXAMPLE 1
Experiment [Br2] [C3H60] Rate in M/s
1 .1 M .1 M 1.64 x 10-5
2 .2 M .1 M 1.65 x 10-5
3 .1 M .2 M 3.29 x 10-5
C3H60 + Br2 C3H5OBr + HBr
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DETERMINING RATE LAWS FROM DATA EXAMPLE 2
Experiment [NO] [Br2 ] Rate in M/s
1 .003 M .003 M 4.04 x 10-4
2 .006 M .003 M 1.62 x 10-3
3 .003 M .006 M 8.08 x 10-4
2NO (g ) + Br2 (g ) 2NOBr(g )
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ZERO ORDER GRAPHICALLY
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1ST ORDER GRAPHICALLY
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2ND ORDER GRAPHICALLY
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