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Buffers and Equilibria context Watch

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    (Original post by ps1265A)
    I also have another question: the dissociation of an acid HA is endothermic, what effect will this have on the Ka if the concentration of the acid increased. I thought that because it's more concentrated, more H+ dissociated, so the conc of products becomes greater, therefore Ka because bigger. But the answer is no change.
    Only one thing affects equilibrium constants and that's temperature ...
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    (Original post by ps1265A)
    Sorry, I mean NaOH
    To what?

    ammonia?
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    (Original post by charco)
    To what?

    ammonia?
    I'm going to start again.

    If we wanted to form a buffer solution say with a weak acid, I can add NaOH to CH3COOH which means I will have a salt of the acid (sodium ethanoate) and a weak acid (ethanoic acid)

    Now let's say we want to make a buffer with a weak base. I can add NH3 with HCl to form NH3Cl which is a salt of my weak base. Why can I not do this?


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    (Original post by charco)
    Only one thing affects equilibrium constants and that's temperature ...
    Ah, we have to consider Equilibria again! Just out of curiosity, why isn't what I said above correct?


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    (Original post by ps1265A)
    I'm going to start again.

    If we wanted to form a buffer solution say with a weak acid, I can add NaOH to CH3COOH which means I will have a salt of the acid (sodium ethanoate) and a weak acid (ethanoic acid)

    Now let's say we want to make a buffer with a weak base. I can add NH3 with HCl to form NH3Cl which is a salt of my weak base. Why can I not do this?


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    You can (although ammonium chloride is actually NH4Cl) provided the weak base is in excess ..
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    (Original post by charco)
    You can (although ammonium chloride is actually NH4Cl) provided the weak base is in excess ..
    Is that to prevent further substitution reactions?


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    (Original post by ps1265A)
    Is that to prevent further substitution reactions?


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    No, it's because a basic buffer consists of a weak base and the salt of a weak base ...

    ... is all.
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    (Original post by charco)
    No, it's because a basic buffer consists of a weak base and the salt of a weak base ...

    ... is all.
    Thanks


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    (Original post by charco)
    Only one thing affects equilibrium constants and that's temperature ...
    Does this mean, say if I was titrating a weak acid with NaOH, despite the volume added of NaOH and the pH rise, the Ka will still stay the same


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    (Original post by ps1265A)
    Does this mean, say if I was titrating a weak acid with NaOH, despite the volume added of NaOH and the pH rise, the Ka will still stay the same


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    Ka is a constant. That means that it doesn't change.

    Only one thing affects equilibrium constants and rate constants (which are related) and that's temperature ...
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    Good work students


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    (Original post by charco)
    Ka is a constant. That means that it doesn't change.

    Only one thing affects equilibrium constants and rate constants (which are related) and that's temperature ...
    Thanks! I think I understand the context of buffers really well now!

    I'm stuck on one part of the context: If we have a strong acid, is the concentration of H+ the SAME as the concentration of the acid (or) is the concentration of H+ makes up most of the concentration of the acid meaning that there is literally 0moldm-3 of the acid ITSELF in the solution?

    Also, is the concentration of A- the same as that of H+? So if we had a strong acid and 95% was dissociated, would this mean we have 95% concentration of H+ and A- EACH?

    Last question, kW is ALWAYS 1.00x10-14 at 298K, but varies at different temperatures. BUT water nevertheless has the same H+ and OH- concentration.
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    (Original post by ps1265A)
    Thanks! I think I understand the context of buffers really well now!

    I'm stuck on one part of the context: If we have a strong acid, is the concentration of H+ the SAME as the concentration of the acid (or) is the concentration of H+ makes up most of the concentration of the acid meaning that there is literally 0moldm-3 of the acid ITSELF in the solution?
    yes,it is the same as the stated original acid concentration

    and yes, there are no acid molecules left in the solution.



    Also, is the concentration of A- the same as that of H+?
    yes





    So if we had a strong acid and 95% was dissociated, would this mean we have 95% concentration of H+ and A- EACH?


    yes
    Last question, kW is ALWAYS 1.00x10-14 at 298K, but varies at different temperatures. BUT water nevertheless has the same H+ and OH- concentration.
    yes
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    (Original post by charco)
    yes,it is the same as the stated original acid concentration

    and yes, there are no acid molecules left in the solution.





    yes





    yes
    Thanks


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    (Original post by charco)
    yes,it is the same as the stated original acid concentration

    and yes, there are no acid molecules left in the solution.





    yes





    yes
    Is a buffer essentially formed at the half-neutralisation point between a weak acid and NaOH? Is this a 2:1?

    Why is it said that we need "excess" acid and a strong base?
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    (Original post by ps1265A)
    Is a buffer essentially formed at the half-neutralisation point between a weak acid and NaOH? Is this a 2:1?

    Why is it said that we need "excess" acid and a strong base?
    If you haven't got more acid than base then you haven't got any left after reaction!
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    (Original post by Pigster)
    Easy peasy way of losing a mark.

    Buffers do not keep pH constant, they minimise pH changes on addition of small amounts of an acid or a base.
    I was speaking qualitatively. Compared to adding the same volume of acid/base to a regular solution then relatively buffers do keep the pH constant. But you're right the pH would change a tiny bit
 
 
 
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