This is the chemistry assignment 🖕🖕

The given standard reduction potentials are:
For magnesium: $Mg^{2+}(aq) + 2e^- \longrightarrow Mg(s)$, $E^\circ_{red} = -2.375\,V$
For aluminum: $Al^{3+}(aq) + 3e^- \longrightarrow Al(s)$, $E^\circ_{red} = -1.660\,V$

For magnesium to displace aluminum ions, magnesium metal ($Mg(s)$) must be oxidized, and aluminum ions ($Al^{3+}(aq)$) must be reduced.
The oxidation half-reaction for magnesium is the reverse of its reduction:
$Mg(s) \longrightarrow Mg^{2+}(aq) + 2e^-$
The standard oxidation potential ($E^\circ_{ox}$) is the negative of the standard reduction potential:
$E^\circ_{ox}(Mg) = -(-2.375\,V) = +2.375\,V$
The reduction half-reaction for aluminum ions is already given:
$Al^{3+}(aq) + 3e^- \longrightarrow Al(s)$, $E^\circ_{red}(Al) = -1.660\,V$

The standard cell potential is the sum of the standard oxidation potential and the standard reduction potential:

Since the standard cell potential ($E^\circ_{cell}$) is positive ($+0.715\,V$), the reaction is spontaneous under standard conditions. This means magnesium will displace aluminum ions from the solution.

The overall balanced reaction is:

Magnesium will displace aluminum ions from the solution.
$\boxed{\text{Yes, Magnesium will displace Aluminum ions.}}$