magnesium metal plus silver acetate in water describes a redox displacement in which magnesium transfers electrons to silver(I) ions. In typical general-chemistry conditions, silver acetate is treated as a soluble source of \(\mathrm{Ag^+}\) and acetate \(\mathrm{CH_3COO^-}\), and magnesium is treated as an active metal that readily oxidizes.
Balanced reaction equations
Molecular equation in aqueous solution
Magnesium acetate remains in solution as ions, while silver forms a metallic deposit.
Complete ionic and net ionic equations
Acetate is a spectator ion under the usual assumption that no strong complexation or precipitation involving acetate dominates the chemistry.
Oxidation states and half-reactions
Electron accounting clarifies the redox roles: magnesium increases in oxidation state from 0 to \(+2\), while silver decreases from \(+1\) to 0.
- Oxidation half-reaction: \(\mathrm{Mg(s) \rightarrow Mg^{2+}(aq) + 2\,e^-}\)
- Reduction half-reaction: \(\mathrm{2\,Ag^+(aq) + 2\,e^- \rightarrow 2\,Ag(s)}\)
- Overall electron balance: 2 electrons transferred per magnesium atom oxidized
Species roles in solution
| Species | Role | Chemical change |
|---|---|---|
| \(\mathrm{Mg(s)}\) | Reducing agent (electron donor) | \(\mathrm{Mg}\) becomes \(\mathrm{Mg^{2+}}\) and enters solution |
| \(\mathrm{Ag^+(aq)}\) (from \(\mathrm{AgCH_3COO}\)) | Oxidizing agent (electron acceptor) | \(\mathrm{Ag^+}\) becomes \(\mathrm{Ag(s)}\) and deposits as metal |
| \(\mathrm{CH_3COO^-(aq)}\) | Spectator ion | Persists in solution; pairs with \(\mathrm{Mg^{2+}}\) as magnesium acetate |
Electrochemical driving force
Standard reduction potentials provide a quantitative justification for spontaneity. Representative values at \(25^\circ\mathrm{C}\) are \(\mathrm{Ag^+/Ag}\) with \(E^\circ \approx +0.80\ \mathrm{V}\) and \(\mathrm{Mg^{2+}/Mg}\) with \(E^\circ \approx -2.37\ \mathrm{V}\). The cell potential for the net ionic process is therefore strongly positive.
The corresponding standard Gibbs energy change for \(\mathrm{Mg(s) + 2\,Ag^+(aq) \rightarrow Mg^{2+}(aq) + 2\,Ag(s)}\) is very negative, with \(n = 2\) electrons transferred.
The equilibrium constant is correspondingly enormous:
Laboratory-scale consequence
Silver metal deposition on the magnesium surface and visible consumption of magnesium are expected, often accompanied by a dark gray or shiny coating as silver plates out.
Visualization of electron transfer and products
Common macroscopic outcomes
- Silver metal formation: a gray-to-silvery coating or crystals on the magnesium surface
- Magnesium consumption: thinning or pitting of the magnesium strip as \(\mathrm{Mg^{2+}}\) forms
- Solution composition shift: increasing \(\mathrm{Mg^{2+}}\) with acetate remaining in solution
Summary statement
Magnesium metal plus silver acetate in water produces a spontaneous redox displacement in which \(\mathrm{Mg(s)}\) is oxidized and \(\mathrm{Ag^+}\) is reduced, yielding \(\mathrm{Mg(CH_3COO)_2(aq)}\) and \(\mathrm{Ag(s)}\) with the net ionic reaction \(\mathrm{Mg(s) + 2\,Ag^+(aq) \rightarrow Mg^{2+}(aq) + 2\,Ag(s)}\).