# Spontaneous reaction between chromium(III) and aluminum

I'm trying to rationalize this problem. It's been a while since I've done redox problems. So here goes, and any input would be great. Basically - am I right?

Here is what happens when you store chromium ion in an aluminum container:

$\ce{Cr^3+ + Al -> Cr + Al^3+ ~~~~~~E^{o}_{cell}=}$ ${+ 0.94}$$\ce{V}$

$\ce{Cr^3+}$: oxidizing agent

$\ce{Al}$: reducing agent

$\ce{Cr}$: conjugate reducing agent

$\ce{Al^3+}$: conjugate oxidizing agent

This reaction is spontaneous because we are going from strong to weak (just like in a spontaneous acid-base reaction).

So if we are going from strong to weak, all these statements are true:

$\ce{Cr^3+}$ is a stronger oxidizing agent than $\ce{Al^3+}$.

$\ce{Al}$ is a stronger reducing agent than $\ce{Cr}$.

• And what is your question? – Hexacoordinate-C Jul 21 '16 at 19:34
• @Shadock - am I right? – Dissenter Jul 21 '16 at 19:35
• For me the reaction you are considering is not correct. Because the reduction potiential of Chromium is -0.74V and the one of the Aluminium is -1.66V. – Hexacoordinate-C Jul 21 '16 at 19:49
• @Shadock - it is still a spontaneous reaction. The standard cell potential is positive. – Dissenter Jul 21 '16 at 19:58
• Oops I haven't read the problem correctly. Let me think – Hexacoordinate-C Jul 21 '16 at 20:14