Chemical Kinetics

Chemical Kinetics - Kinetics: Rates of Chemical Reactions...

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Unformatted text preview: Kinetics: Rates of Chemical Reactions Rate of a reaction: change in concentration / unit time can relate to any product or reactant [X] = "molar concentration of X" [reactant] decreases with time [product] increases with time product concentration reactant time Measuring reaction rates: • appearance of a product • disappearance of a reactant • Detection of change: direct detection of a chemical species Indirect methods: change in color, pH, pressure measuring a product: rate = ∆ [product] concentration is increasing (+) ∆ time measuring a reactant: rate = - ∆ [reactant] concentration is dereasing (- ) ∆ time Use (+) and (-) signs on concentration terms to make rate (+) Different ways to measure rate: 1 • average rate: over a specified time interval, from a specified starting time to a specified ending time. • instantaneous rate: slope of line tangent at a specified time. • initial rate: linear approximation of the first few moments. The effect of stoichiometry on rate: N 2 (g) + 3 H 2 (g) → 2 NH 3 (g) to compare rates: H 2 disappears 3 times as fast as N 2 Their absolute rates are NOT THE SAME ! Reaction rate with respect to each compound: rate (N2) = -∆ [N 2 ] rate (H2) = -∆ [H 2 ] ∆ time ∆ time rate (NH3) = ∆ [NH 3 ] THESE ARE ALL ∆ time DIFFERENT RATES! To establish a common rate of reaction, divide H 2 rate by 3 and NH 3 rate by 2 (their coefficients in the balanced chemical eqn.) reaction rate = -∆ [N 2 ] = - 1 ∆ [H 2 ] = + 1 ∆ [NH 3 ] ∆ time 3 ∆ time 2 ∆ time Example: The concentration of a reaction product was measured spectroscopically during the course of a reaction. The concentration 2 was 0 at time t = 0, and the concentration was 0.0350 M at time t = 15.0 minutes. What was the rate of the reaction during this interval, in M/sec ? Collision theory of reactions: particles must collide must collide with sufficient energy must collide with proper orientation Things that affect rate: concentration of the reactants: more molecules means more collisions more collisions means more chances to "get it right" temperature : kinetic energy T ↑ v ↑ 1/2 mv 2 ↑ as v ↑ number of collision per time ↑ as v ↑ energy of collisions ↑ number of collisions with correct orientation ↑ presence of a catalyst lower energy intermediates lower energy pathway 3...
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This note was uploaded on 01/18/2012 for the course CHEM 307 taught by Professor Dr.owenmcdougal during the Fall '11 term at Boise State.

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Chemical Kinetics - Kinetics: Rates of Chemical Reactions...

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