ocr a f325 revision thread

If you didn't divide by (-ΔS), then the '-' (negative) would stay with the T, making the equation: (G-ΔH)/ΔS=-T. Another general example for this would be if the equation was C=D-AB, then to make the equation equal to AB you would move the D over (C-D=-AB) and then divide by (-A) to make it equal B: (C-D)/-A=B. Again, if the negative wasn't moved over, then the equation would be (C-D)/A=-B.

When you move the A in A+B=C, the sign of the A changes, because that is what is being moved over, so you get B=C-A. It would be like if you had 3+1=4, if you move the 1 over, you get 3=4-1, which is true. If it was the other way around (like B=A-C in the example), then the equation would become 3=1-4, which isn't true.

Basically, if you are moving something over that is by itself (like A, B or ΔH), then the sign changes, but if you are dividing (or multiplying) by something to split up a term (like AB, or CD, or TΔS), then the sign stays the same.

It's kw = [oh-][h+] isnt it?

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No I mean like if a reaction has H2O in e.g. H2O2 ==> H2O + 1/2O2 then do you include H2O in Kc? Cause there's one expression where apparently H2O is in so much excess it is considered to be constant and you don't include it

=S

Can anyone help me with this please??!

I was wondering if anyone could help me with 6d in the January 2011 paper.

'When dissolved in water, the enthalpy change of solution of the salt potassium fluoride, KF, is -15kJ mol^-1.

The salt rubidium fluoride, RbF, has an enthalpy change of solution in water of -24kJ mol^-1.

Suggest reasons for the difference between the enthalpy changes of solution of KF and RbF.'

The thing that I'm confused about isn't directly asked about in the question, but it is still bugging me and I've spent ages trying to work it out. Why is the enthalpy change of solution of RbF more exothermic than that of KF? K⁺ has a smaller ionic radius than Rb⁺, so KF should have more exothermic/negative enthalpy change of hydration and lattice enthalpy values than RbF. I know the mark scheme says that lattice enthalpy affects delta H solution more than delta H hydration does, but that shouldn't make a difference, should it, as KF is more exothermic in both cases?

Guys can someone explain to me how a buffer works??

I was wondering if anyone could help me with 6d in the January 2011 paper.

'When dissolved in water, the enthalpy change of solution of the salt potassium fluoride, KF, is -15kJ mol^-1.

The salt rubidium fluoride, RbF, has an enthalpy change of solution in water of -24kJ mol^-1.

Suggest reasons for the difference between the enthalpy changes of solution of KF and RbF.'

The thing that I'm confused about isn't directly asked about in the question, but it is still bugging me and I've spent ages trying to work it out. Why is the enthalpy change of solution of RbF more exothermic than that of KF? K⁺ has a smaller ionic radius than Rb⁺, so KF should have more exothermic/negative enthalpy change of hydration and lattice enthalpy values than RbF. I know the mark scheme says that lattice enthalpy affects delta H solution more than delta H hydration does, but that shouldn't make a difference, should it, as KF is more exothermic in both cases?

a buffer is made up of a weak acid and its conjugate base - which could be obtained from the salt of the weak acid
the weak acid partially dissociates
the salt fully dissociates
therefore, you have a high conc of weak acid and conjugate base, and few H+
a buffer system is formed - which is basically the dissociation of the weak acid
HA >< (my reversible arrows!) H+ + A-
on addition of small amounts of alkali, the few H+ present will react to form H20 as H+ + OH- goes to H20. this will cause HA to dissociate, shifting the equilibrium to the right to replace the H+ ions that reacted
on addition of small amounts of acid, this will react with the A- ions and so shift the equilibrium to the left, to decrease the excess H+ ions
hope that helps!!

_

thats what I got, the mark scheme made a mistake

I was wondering if anyone could help me with 6d in the January 2011 paper.

'When dissolved in water, the enthalpy change of solution of the salt potassium fluoride, KF, is -15kJ mol^-1.

The salt rubidium fluoride, RbF, has an enthalpy change of solution in water of -24kJ mol^-1.

Suggest reasons for the difference between the enthalpy changes of solution of KF and RbF.'

The thing that I'm confused about isn't directly asked about in the question, but it is still bugging me and I've spent ages trying to work it out. Why is the enthalpy change of solution of RbF more exothermic than that of KF? K⁺ has a smaller ionic radius than Rb⁺, so KF should have more exothermic/negative enthalpy change of hydration and lattice enthalpy values than RbF. I know the mark scheme says that lattice enthalpy affects delta H solution more than delta H hydration does, but that shouldn't make a difference, should it, as KF is more exothermic in both cases?

Worked example:

thank you soo much, is that all you need to know?? Soo I just apply that to the questions??

THANK you, and if it was not feasible , it would have been Cr+ what