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    Hi,
    I have done an A2 investigation between magnesium and hydrochloric, sulphuric and ethanoic acid and need to work out the rate constants and then the arrhenius equation.
    I did different concentrations of each acid so got a different rate for each concentration so am I going to have several rate constants?
    If I have help for one acid, I can work out the rest myself
    For ethanoic acid I got:
    Concentration Rate
    2.4 0.4
    2.1 0.325
    1.8 0.238
    1.5 0.163
    1.2 0.1
    I worked out the order for the reaction is second order.
    I really am stuck so could do with some help please!
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    I dont know what you're investigating from your post... But if we represent acid as [A] and magnesium as [B], then the rate equation will be rate=k[A]^m[B]^n, where m and n are orders you presumably know. Just rearrange to k=rate/[A]^m[B]^n to find k. So for ethanoic acid, k=0.4/2.4^m*[B]^n.
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    by rate constant, do you mean Kc???? if so isn't that Kc = [dissociation of acid]/[acid]
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    (Original post by boromir9111)
    by rate constant, do you mean Kc???? if so isn't that Kc = [dissociation of acid]/[acid]
    :nah:
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    oh, then it's what you said "LearningMath".... rate = k[acid] and switch around to get k.... problem solved.
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    (Original post by LearningMath)
    I dont know what you're investigating from your post... But if we represent acid as [A] and magnesium as [B], then the rate equation will be rate=k[A]^m[B]^n, where m and n are orders you presumably know. Just rearrange to k=rate/[A]^m[B]^n to find k. So for ethanoic acid, k=0.4/2.4^m*[B]^n.
    Thanks that's helped, I've found out that Mg is zero order so I don't think it would be included in the equation so would this be correct?
    k = 0.4/2.4^2 = 0.0694

    Then I do it for all the other concentrations?
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    (Original post by sophia27)
    Thanks that's helped, I've found out that Mg is zero order so I don't think it would be included in the equation so would this be correct?
    k = 0.4/2.4^2 = 0.0694

    Then I do it for all the other concentrations?
    yep
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    Thanks I'm so happy now!
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    (Original post by sophia27)
    Thanks that's helped, I've found out that Mg is zero order so I don't think it would be included in the equation so would this be correct?
    k = 0.4/2.4^2 = 0.0694

    Then I do it for all the other concentrations?
    k doesnt vary with concentration, but you'l have experimental error so it may be wise to calculate k for all concentrations and average.
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    (Original post by boromir9111)
    yep
    Ok, sorry to be a pain but I'm stuck again!
    I think I've given the wrong information, I think I have to use the results from when I changed the temperature and here is what I got:
    Temperature Rate
    50 0.375
    40 0.213
    30 0.163
    20 0.125
    10 0.113

    The concentration of the acid is 1.2moles. The order for ethanoic acid in this experiment is 1st order and magnesium zero order.
    To put the numbers in the arrhenius equation, shouldn't the rate constant be constant for all the temperatures?
    Sorry again, I think I have done my experiment wrong and don't really know what I'm doing
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    (Original post by boromir9111)
    by rate constant, do you mean Kc???? if so isn't that Kc = [dissociation of acid]/[acid]
    kc is the equilibrium constants! u dont have an upcoming chem exam do you??? lol im only messing i only really understood kc just before i went into the exam, no joke my friend explained it to me!! and i rocked it lol
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    (Original post by sophia27)
    Ok, sorry to be a pain but I'm stuck again!
    I think I've given the wrong information, I think I have to use the results from when I changed the temperature and here is what I got:
    Temperature Rate
    50 0.375
    40 0.213
    30 0.163
    20 0.125
    10 0.113

    The concentration of the acid is 1.2moles. The order for ethanoic acid in this experiment is 1st order and magnesium zero order.
    To put the numbers in the arrhenius equation, shouldn't the rate constant be constant for all the temperatures?
    Sorry again, I think I have done my experiment wrong and don't really know what I'm doing
    The rate constant should vary with temperature, it doesnt vary with concentration! What are you actually trying to achieve with the arrhenius equation?! I presume you want to find activation enthalpy, in which case you need to calculate the rate constant at the temperatures you've shown above, and then plot lnk against 1/T and take the gradient:

    The arrhenius equation:

    k= A e-Ea/RT

    lnk = lnA -(Ea/R)*(1/T)

    Put this in the form y=mx+c

    lnk = -(Ea/R)*(1/T) + lnA

    So you have gradient equal to Ea/R, and R is constant so you can find Ea

    This stuff isnt taught at A-level, i had to do quite a bit of research to understand it for my investigation. So if you havnt been introduced to it or researched it, you're going to have a hard time! Good luck though, and you do have the right measurements!
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    (Original post by sophia27)
    Ok, sorry to be a pain but I'm stuck again!
    I think I've given the wrong information, I think I have to use the results from when I changed the temperature and here is what I got:
    Temperature Rate
    50 0.375
    40 0.213
    30 0.163
    20 0.125
    10 0.113

    The concentration of the acid is 1.2moles. The order for ethanoic acid in this experiment is 1st order and magnesium zero order.
    To put the numbers in the arrhenius equation, shouldn't the rate constant be constant for all the temperatures?
    Sorry again, I think I have done my experiment wrong and don't really know what I'm doing
    rate constant changes with temperature, refer to arrhenius equation. That is how you can get activation energy. Activation energy is the one that is constant for each particular reaction.
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    (Original post by uni4med)
    kc is the equilibrium constants! u dont have an upcoming chem exam do you??? lol im only messing i only really understood kc just before i went into the exam, no joke my friend explained it to me!! and i rocked it lol
    It's actually called "acid dissociation constant" but this is measuring the initial rate(reactant) i believe so there is no need for Kc as i changed it around after realizing that but thanks for your reply and good luck with medicine, sure you will be fine!!!
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    (Original post by sophia27)
    Ok, sorry to be a pain but I'm stuck again!
    I think I've given the wrong information, I think I have to use the results from when I changed the temperature and here is what I got:
    Temperature Rate
    50 0.375
    40 0.213
    30 0.163
    20 0.125
    10 0.113

    The concentration of the acid is 1.2moles. The order for ethanoic acid in this experiment is 1st order and magnesium zero order.
    To put the numbers in the arrhenius equation, shouldn't the rate constant be constant for all the temperatures?
    Sorry again, I think I have done my experiment wrong and don't really know what I'm doing
    K is temperature dependant so you should get a different value for each temperature. For 50 k=0.375/1.2 = 0.3125. But why is it now first order when in the first post you said it was second order?
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    (Original post by sophia27)
    Ok, sorry to be a pain but I'm stuck again!
    I think I've given the wrong information, I think I have to use the results from when I changed the temperature and here is what I got:
    Temperature Rate
    50 0.375
    40 0.213
    30 0.163
    20 0.125
    10 0.113

    The concentration of the acid is 1.2moles. The order for ethanoic acid in this experiment is 1st order and magnesium zero order.
    To put the numbers in the arrhenius equation, shouldn't the rate constant be constant for all the temperatures?
    Sorry again, I think I have done my experiment wrong and don't really know what I'm doing
    You certainly shouldn't be worried about the rate changing with temperature, I'd be very worried if it didn't change with temperature. Cooking is a nice example of chemical change (and you get to eat your products, it's way better than chemistry). Stuff happens faster at higher temperatures.

    It should follow the Arrhenius equation as you said. That is:

    k = Ae^{\frac{-E_a}{RT}}

    The next step depends on how much maths you've done...if you know logs then:

    \ln{k} = \frac{-E_a}{RT} + \ln{A}

    That will give you a straight line plot from which you can work out the activation energy, E, which is all the information you need to work out the rate constants at different temperatures.

    If you don't know logs, then try taking some ratios of ks at different temperatures, bearing in mind that your E, R and A will stay the same. That will let you calculate k for any temperature (just sub it in).
 
 
 
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