Lecture Recording:
Lecture Summary/Notes:
OK, so basically, we started this course with chemical kinetics!
Key Questions covered in the lecture:
-What makes a reaction fast or slow?
-How do we measure or define that (the rate of the reaction)?
Kinetics, unlike thermochemistry, is very UNPREDICTABLE!
To explain this - if you look at a curved function, at every point, the tangent is a different slope, right? The energy of a reaction (E) is like a curved function, with a rising hill, and then a drop, like this!
How do we define the rate of the reaction?
**We must define the conditions first (since they are always changing!)
WAYS to define the rate:
Average rate (change in concentration over the change in time).
Initial rate (slope of the tangent at x=0).
Both the average rate and the initial rate are not very accurate ways to measure rate... Initial rate depends on the initial conditions of the reaction!
SO WHAT DO WE USE????
The rate equation (or the definition of a rate)
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The order is usually a very small integer or simple fraction!
The Zero Order Reaction (where m = 0 in R= k[A]^m)
Looks like this!
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R= k[A]^0
R=k
Slope m (from y=mx+b) is -k (since the reactant is disappearing)
[A]t = -kt + [A]
[A]t = -kt + [A]
(the concentration of A at any time t = -kt + initial amount of reactant)
At this point, we started talking about Radioactive Decay, which is always a first order reaction! (first order reactions are the most common in nature). That brought us somehow to the topic of the integrated rate law.
I wish I had a classmate like you when I was in university
ReplyDeletelol :P
ReplyDeletethanks so much!
ReplyDeleteNo problemmm~~~ :D
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