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S&S part (d)(ii) (1 Viewer)

quarm

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ok...this is the question im most unsure of in the whole test

i mean, i still havnt figured out how the question can be worth 4 marks

i just went on about how stuff corroded faster in a beaker with HCl and didnt corrode as fast in acetic acid and even slower in pure water

hope that's right coz i made it up on da spot

i mean...i wrote more on part (i) than (ii) even tho (i) is worth half the marks as (ii)

most of my fwends wrote onli a paragraph too

did any1 else find that question hard?
 

ash_s137

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Um I added in why... like greater electrode potential when in acidic environment using two dif eqns which i couldn't be bothered to write here right now.
 

McLake

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Yeah, weird 4 mark q. I also talked about why acid speeds up corrosion. I talked to others and they said that's what they thought so ...
 

spice girl

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anyone thought about the fact that it is only SOME metals corrode faster in acid?

Some don't corrode in anything at all.

For all I know, copper isn't affected by acid.
 

McLake

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Just to add to what spice girl said, the question actually asked about MATERIALS in acid, not just METALS. Did anyone consider this?
 

Weisy

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and pracs done in class only used metals and alloys
 

idontknow

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i thought the question was based towards the interpretation of the results and how the results showed the differing rates of corrosion. My test was with the indictator for the Fe ions. But 4 marks was alot to give for an interpretation so i also went on to explain how the results are justified chemically using relevant half equations.
 

McLake

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Well we did our prac with wood, fabric and cermaics. Wood DID "corrode" slightly (hmm, seeing problem with the word corrode ...)
 

BlackJack

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I created a theory on how H+ reacted with OH- in the reduction half-equation, and this allows the metal ions to disperse throughout the solution without interfering with further corrosion on the surface... then talked out how H3O+ was a quicker transporter of charge, because it was smaller and could transfer the H+ ion to anothe H2O atom, speeding up the net 'speed' (like how sound travels faster in metals because the atom impart energy quickly to the next atom...).

I just stated that only some metals supported this hypothesis because it depends on their reactivity in the first place to 'oxidise' into the solution.
 

McLake

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Originally posted by profernity
I mean:
H+ ions can oxidise Fe:

Fe + 2H+ --> Fe2+ + H2 + 2e-
I'm sure those e- shouldn't be there ...
 

profernity

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just to demonstrate how quickly i emptied my brain after chem, i had meant to say (third time lucky):

hydrogen ions can reduce Fe

Fe + 2H+ + 2e- --> Fe2+ + H2

goddamn. i still don't know if that works, but i remember it's in pathways or something.
 

McLake

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Originally posted by profernity
just to demonstrate how quickly i emptied my brain after chem, i had meant to say (third time lucky):

hydrogen ions can reduce Fe

Fe + 2H+ + 2e- --> Fe2+ + H2

goddamn. i still don't know if that works, but i remember it's in pathways or something.
Isn't it: Fe + 2H+ --> Fe2+ + H2 ?
 

spice girl

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Originally posted by BlackJack
I created a theory on how H+ reacted with OH- in the reduction half-equation, and this allows the metal ions to disperse throughout the solution without interfering with further corrosion on the surface... then talked out how H3O+ was a quicker transporter of charge, because it was smaller and could transfer the H+ ion to anothe H2O atom, speeding up the net 'speed' (like how sound travels faster in metals because the atom impart energy quickly to the next atom...).

I just stated that only some metals supported this hypothesis because it depends on their reactivity in the first place to 'oxidise' into the solution.
Well, basically H+ is an oxidant, and OH- is a reductant. I don't think water takes part in the actual redox interaction at all: sometimes we write H2O as a reactant because it is more abundant than the corresponding oxidant/reductant:

e.g. In basic medium, we write H2O as the "oxidant" - like:
H2O + e- -> (1/2)H2 + OH-, but it is actually a combination of:

1. H2O -> H+ + OH- : this is the dissociation reaction
2. H+ + e- -> 1/2H2 : this is actual redox reaction - reduction of H+

According to kinetic theory, reaction rate is proportionate to the concentration of reactants. Thus in oxidation reactions, reaction rate is proportionate to [H+]

As to why "rust" doesn't form, it doesn't because the reaction (or precipitation reaction): Fe2+ + 2OH- <-> Fe(OH)2{s} equilibrium lies to the left because of the absence of OH- ions in an acidic solution.

:confused:
 

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