CHEM 1001 Lecture Notes - Lecture 99: Protein Kinase B, Enzyme, Forward Rate

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8 Apr 2016
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
Kinetics, Equilibrium, Thermodynamics, Electrochemistry.
:  - Unimolecular and Bimolecular.

-Kinetics: Rate of change of concentrations. Rate units: M/s = mol / (L*s).
Amount/time.
-Reaction rates are always positive. Reaction rates are NOT constant.
-Sn2+ + 2Fe3+ --> Sn4+ + 2Fe2+ Rate of formation of Sn4+ = ∆[Sn4+] / ∆t.
Rate of formation of Fe2+ = ∆[Fe2+] / ∆t. 2 moles of Fe2+ and 1 mole Sn4+ so rate of
reaction of Fe2+ is twice faster. Properly balance would be: Sn4+ = ∆[Sn4+] / ∆t. And
1/2Fe2+ = ∆[Fe2+] / ∆t. Notice the reaction rate is based on the moles in the reaction.
-:
-Concentrations of reactants, products, other species (catalysts).
-Temperature
-Solvent, phase, pressure.
-Other reactions.
-The  rate, Δ[H2O2]/Δt, is determined by the slope between two points,
as in the previous table.
-The !rate, ∆[H2O2]/∆t, is determined by the slope (tangent) at
one point. The instantaneous rate is the correct value.
-Between the two, the instantaneous rate is the correct value.
-Slide 25 - Rate of reaction decreases as H2O2 is consumed.
-Two types of rate laws: ": Reaction rate as function of concentration.
-# : Concentration as function of time.
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-An input of energy is normally needed: Collisions with other molecules.
Absorbing radiation (light).
- Definition: The half-life is the time it takes for one half of a reactant to
be converted into product.
-$!: These are elementary reactions between two
molecules. NO + O3 → NO2 + O2 HO2 + HO2→ H2O2 + O2
-The reaction must be written so that all stoichiometric coefficients are
integers.
-Correct: 2HO2 → H2O2 + O2
-Incorrect: HO2 → 1⁄2H2O2 + 1⁄2O2
-Rate law for Bimolecular reactions: The rate of reaction is directly
proportional to the concentrations of each of the reacting molecules.
NO + O3 → NO2+ O2
%reaction rate &'()*+,-'&.()*,(*+,
2HO2 → H2O2 + O2
% reaction rate &'(+*+,-'&.(*+,/+
k is the 0.
NO + O3 → NO2 + O2
%reaction rate &'()*+,-'&.()*,(*+,
Units of concentration are usually M, but not always.
Units of time are seconds, or minutes, or hours, etc.
The reaction rate has units of concentration/time.
The rate constant has units of 1/concentration time.
k is different for each reaction and may depend on temperature, solvent, phase, and
pressure.
This is an example of a: Second-order reaction.
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-Typically (but not always), the rate law is of the form:
rate of reaction = k[A]^m [B]^n.
-[A] and [B] represent reagent molarities. ., , and are constants for any
particular reaction.
There may be more or fewer than two reagents in the rate law.
The reagents are usually, but not always, reactants or products.
and are often small, positive integers. But they may be zero, fractional, and/or
negative.
The Constants m,n and k: &.123/1$3/
The reaction is of order m for A, order n for B, etc. The overall
reaction order is m + n.
k > 0 for every reaction that goes. If k < 0, then reverse rxn occurs.
.is a proportionality constant.
.is called the rate constant.
The faster the reaction, the larger the value of k.
.depends on temperature, pressure, solvent, phase.
The units for .depend upon the reaction order.
Examples of reaction order:
Reaction rate = .[N2O4]
First order in N2O2
First order overall
Reaction rate = .[NO][O3]
First order in NO
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