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Rates of Reaction

  1. Jan 11, 2005 #1
    I have some quick questions that I need some help with:


    Why would a simple chemical reaction such as

    NO9g) + 1/2O2(g) -----------> NO2(g) most likley be most rapid at the start?


    In what ways can the rate of dissolving a lump of sugar be increased? (at least 3)


    Finally, what effect would there be on the rate of constant of, forward reaction, and reverse reaction if there was an increase in temperature?
  2. jcsd
  3. Jan 12, 2005 #2


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    ChemRookie, you are asking us to do your homwork for you. We can't do that.

    But we can help if you show what you have tried so far.

    At least 2 of the 3 questions are pretty straightforward, and the answers should be in your text. You should be able to answer them without understanding a thing about kinetics. To answer the first question, you must first understand the concept of a reaction rate, and why it is the way it is. Once you get this, #1 is pretty obvious.
  4. Jan 12, 2005 #3
    1) Think about what particles are present at the start - ie the concentration of the reactants are maximum at the start.

    2) Look up how surface area : volume ratio affects reaction rate

    3) Look up kinetic theory/Boltzman distribution.
  5. Jan 16, 2005 #4

    1) Would the answer be that the chemicals involved in this reaction have a electron configuration in which they easily give off their electrons..therefore, easily react. ? Would it be specific to what chemicals they are, or is there just a rule that reactions are faster at the beginning?

    2) obviously one is temperature, possibly also stirring faster? and maybe using a catalyst? (another substance to help it along)

    3) for number 3. I know that a forward reaction is a reaction where the concentration increases with time, and a reverse is the other way around. So, how would an increase in temp affect the rate of constant of those 2..well, for a forward reaction, the the concentration would not increase I guess..since the reaction would happen quicker? for a reverse reaction, I guess it would have no affect?
  6. Jan 17, 2005 #5


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    #3. The equilibrium constant changes as the temperature varies since the ratio Rate forward/Rate reverse changes, find in your text a relationship/derivation of the equilibrium constant from the rate costants. Also note that an increase in temperature would now facilitate both reactions, since the energy of the system is now closer to the activation energies of both reactions.......it is most often the case that the one reaction limited by kinetics will have its rate increased more in a relative sense.
  7. Mar 3, 2005 #6
    one more question, here:

    For the reaction

    CO + NO2 ----> CO2 + NO the activation energy for the forward reaction is 135kj/mol of CO reacted.

    a) determine the heat of reaction
    b) from the data given, and the deltaHr for the reaction, determine the activation energy for the reverse reaction.
    c) draw and label a potential energy diagram for the reaction

    can someone help me with a and b at least..thanks. someone told me I dont have enough data to do the question and looks that way to me too.
  8. Mar 6, 2005 #7
    can anyone answer if I have the neccesary info here? or any help with it

  9. Mar 6, 2005 #8
    a) Arrehenius Equation
    b) Hess' Law
  10. Mar 7, 2005 #9


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    And I would do (c) before doing (b).
  11. Mar 10, 2005 #10
    but, I dont have all of these variables.


    the question:
    and once I get the answer, how would I get the reverse reaction?
  12. Mar 13, 2005 #11

    T= 0degcel=293K?
    P= 8.314J/mol K
    E(a) = 135kj/mol
    A= ?
    what's the arrhenius constant?

    k = -Ea/RT + ln A
  13. Apr 1, 2009 #12
    The T should equals to 273k
  14. Apr 9, 2009 #13
    Question 1. According to reaction rate and experimentally approved stated that rate is direct proportion to the concentration.At the initial time there are the highest concentration as long as the reaction took places the concentration gradually decrease simultaneously with rate of reaction.
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