dot point help 9.6.6. (1 Viewer)

[stargaze]

- Explain that acidic environments accelerate corrosion in non-passivating metals


Perform a first-hand investigation to compare and describe the rate of corrosion of metals in different acidic and neutral solutions

Um, for the prac one... whats t he actual results (cos i didnt finish it)

thanks~

 

[rama_v]

Metals corroded much faster in acidic solutions. In neutral solutions there was some corrosion and almost no corrosion in alkaline solutions.

 

[mitochondria]

But why?

I should make awareness to the fact that the dot point says explain.

stargaze, hint hint, this is electrochemistry. Have a look at the factors that affect an electrolytic cell and see if you can come up with a reason why this is so (and perhaps also the equations involved in a corrosion, say, rusting). Come back if you can't figure it out ^^

 

[LostAuzzie]

Theres two main reasons why Acidic environments accellerate corrosion:

1. Take a look at the Standard reduction potentials for Oxygen in acidic and neutral environments:
1/2 O2 + H2O + 2e -> 2OH 0.40 V
1/2 O2 + 2H + 2e -> H2O 1.23 V
From the above results you see the standard potential for oxygen in acidified conditions is more positive and thus oxygen in this circumstance is more likely to undergo reduction, the ellectrons coming from the oxidation of the metal.

2. As you would remember from Preliminary as well as Acids & Bases, Metals will displace hydrogen from solution, Adding to the corrosion by oxidation.

The success one book makes mention of a third thing involving OH ions but my teacher reckons I should just dismiss it and stick with my reasoning

Since the dot point specifically mentions non-passivating metals, I would make mention of the fact that passivating metals form the inert layer resistant to most chemical attack. Non passivating metals do not have this protection.

With regards to the prac:

As previously mentioned, the most corrosion occurred in acidic solution, which I explain using the above two examples, in regards to equations you could quote the two standard potentials for oxygen in acidic and neutral solutions, and you could use something like this for the second reasoning:
Fe(s) + H(aq) -> Fe(aq) + H(g) (It would be more understandable with +'s and -'s for ions so I apologise for that)
In neutral there is a balance between OH and H so that average corrosion occurs, while in basic solutions (excess OH not enough H) there is very little corrosion.

I hope that helps somewhat

 

[Haku]

could u elaborate on ur second point on metal and acids?

what u mean by displace?

 

[LostAuzzie]

Similar to metal displacement, Hydrogen will come out of solution while the metal will go into the solution (become an ion)
I apologise it isnt really all that clear without the + or - to indicate ions.

 

[Haku]

hehe thanks, alot clearer.

but would the displacement occur if the metal have a lower oxidation potential than hydrogen? ie: cu with reduction potential of 0.34

 

[LostAuzzie]

Nope the hydrogen would stay in solution. I probably should have stated 'more active metals' but we are mainly focusing on Iron which is above Hydrogen.

 

[Haku]

wait, if it oxidises iron it would gain an electron and come out as gas

but nothing would happen if it is mixed with cu solid right? cause there is no redox taking place.

 

[LostAuzzie]

Exactly, no redox takes place when copper is placed in the acid.

 

[rama_v]

Theres two main reasons why Acidic environments accellerate corrosion:

1. Take a look at the Standard reduction potentials for Oxygen in acidic and neutral environments:
1/2 O2 + H2O + 2e -> 2OH 0.40 V
1/2 O2 + 2H + 2e -> H2O 1.23 V

The success one book makes mention of a third thing involving OH ions but my teacher reckons I should just dismiss it and stick with my reasoning

Yeah, I came across that the other day. I think what the success one book is getting at is that when you have hydrogen ions in the water, the first reaction that you've written is accelerated as per Le Chatlier's principle. This is because, if you take a look at the HSC data sheet, all the reactiosn are written as equilibirum reactions, so when the H⁺ ions react with the OH^- ions, more OH^- ions are produced to counteract the change, driving the reaction forward. I think this is what the second equation basically shows. Hence the larger reduction potential in the acidic water

 

[Haku]

Yeah, I came across that the other day. I think what the success one book is getting at is that when you have hydrogen ions in the water, the first reaction that you've written is accelerated as per Le Chatlier's principle. This is because, if you take a look at the HSC data sheet, all the reactiosn are written as equilibirum reactions, so when the H⁺ ions react with the OH^- ions, more OH^- ions are produced to counteract the change, driving the reaction forward. I think this is what the second equation basically shows. Hence the larger reduction potential in the acidic water

lol, i am kinda slow, read it 4 times and still dun get u.

i think just by explaining that the oxidation potential of a acidicified environment containing H+ ions is higher than a neutral solution. Thus this will more readily oxidise the metal, causing a redox reaction with the metal being oxidised to form metal ions, and will combine with anions in water to form salths, ie: AgS..

 

[stargaze]

Ok, factors that affect electrolytic cell are:
- Voltage
- Concentration of Electrolyte
- Electrode surface area
- Electrode distance from each other

Referring to this, I think that by the acidic concentration of the electrolyte increasing, the:
- Voltage goes up (as Oxygen is reduced, thus providing a higher voltage as a result of the greater level of H+ ions that come from acid)
- Concentration of acid (H+ ions) increase, thus causing the above

An increasing voltage --> increasing rate of electrolysis --> increasing rate of corrosion

Is that ok?

 

[l-mercedes-l]

This was kind of consfusing but

extra H+ in the acidic solution acts as a CATALYST for the oxidation of metal.

O2 + 4H+ + 4e- --> 2H20

this reaction needs electrons from somewhere and for some strange reason the solid metal is happy to oblige.

Metal(s) --> Metal(aq)n+ + ne-

so then what you have happening goes something like this...

eg/ 2Fe(s) + 3/2 O2(g) + nH20(l) + 4H+(aq) --> Fe2O3.nH2O(s) + 4H+

At depth where CO2 conc is higher and the water more acidic, corrosion is ACCELERATED by hydrogen ions acting as CATALYSTS.

I'd love to have written this message in pink... colours always help me to understand chemistry coz its much more confusing in black and white. I think that this is right... If its not can someone let me know... lol

-Saedie

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