Using the Electrochemical Series Questions
You need to use the Electrochemical Series Data to answer these questions.
1. A cell is set up with a Ni²⁺ / Ni half-cell and a Cr³⁺ / Cr²⁺ half-cell.
a) Write the notation for the cell.
Pt(s) | Cr²⁺(aq), Cr³⁺(aq) || Ni²⁺(aq) | Ni(s)
b) Write the ionic equation for the reaction.
c) Identify the oxidising agent.
d) Calculate the E.M.F. of the cell.
-0.25 – (-0.41) = +0.16 V
e) State which half-cell is the anode/cathode.
Anode (Oxidation): Cr³⁺/Cr²⁺ half-cell.
Cathode (Reduction): Ni²⁺/Ni half-cell.
f) State the direction of electron flow.
From the Cr³⁺/Cr²⁺ half-cell (Left) to the Ni²⁺/Ni half-cell (Right).
g) State the purpose of the salt bridge.
To maintain charge balance and complete the circuit by allowing ions to move.
2. A cell is set up with a S.H.E. and a [Ni(NH₃)₆]²⁺ / Ni half-cell.
a) Write the notation for the cell.
Pt(s) | H₂(g) | H⁺(aq) || [Ni(NH₃)₆]²⁺(aq), NH₃(aq) | Ni(s)
(S.H.E. placed on the left as it is the standard reference).
b) Write the ionic equation for the reaction.
2H⁺ + Ni + 6NH₃→H₂ + [Ni(NH₃)₆]²⁺
c) Calculate the E.M.F. of the cell.
d) State the direction of electron flow.
From the Nickel complex half-cell (Right) to the S.H.E. (Left).
3. A cell is set up with a Pb²⁺ / Pb half-cell and a Fe³⁺ / Fe²⁺ half-cell.
a) Write the notation for the cell.
Pb(s) | Pb²⁺(aq) || Fe³⁺(aq), Fe²⁺(aq) | Pt(s)
b) Write the ionic equation for the reaction.
c) Identify the reducing agent.
d) Calculate the E.M.F. of the cell.
+0.77 – (-0.13) = +0.90 V
e) State which half-cell is the anode/cathode.
Anode (Oxidation): Pb²⁺/Pb half-cell.
Cathode (Reduction): Fe³⁺/Fe²⁺ half-cell.
f) State the direction of electron flow.
From the Pb half-cell (Left) to the Fe half-cell (Right).
4. A cell is set up with a Cr³⁺ / Cr₂O₇²⁻ half-cell and a Cr²⁺ / Cr half-cell.
a) Write the notation for the cell.
Cr(s) | Cr²⁺(aq) || Cr₂O₇²⁻(aq), H⁺(aq), Cr³⁺(aq) | Pt(s)
b) Write the ionic equation for the reaction.
Cr₂O₇²⁻ + 14H⁺ + 3Cr→2Cr³⁺ + 7H₂O + 3Cr²⁺
c) Calculate the E.M.F. of the cell.
+1.33 – (-0.91) = +2.24 V
d) State which half-cell is the anode/cathode.
Anode: Cr²⁺/Cr half-cell.
Cathode: Dichromate half-cell.
e) State the direction of electron flow.
From the Cr half-cell (Left) to the Dichromate half-cell (Right).
5. A cell is set up using an Ag⁺ / Ag half-cell and a Mn²⁺ / MnO₄⁻ half-cell.
a) Write the notation for the cell.
Ag(s) | Ag⁺(aq) || MnO₄⁻(aq), H⁺(aq), Mn²⁺(aq) | Pt(s)
b) Write the ionic equation for the reaction.
MnO₄⁻ + 8H⁺ + 5Ag→Mn²⁺ + 4H₂O + 5Ag⁺
c) Calculate the E.M.F. of the cell.
+1.52 – (+0.80) = +0.72 V
d) State which half-cell is the anode/cathode.
Anode: Ag/Ag⁺ half-cell.
Cathode: Manganate half-cell.
e) State the direction of electron flow.
From the Ag half-cell (Left) to the Manganate half-cell (Right).
6. A cell is set up using a secondary silver standard half-cell and a mystery half-cell. The voltmeter reads -1.05V.
a) State the identity of the mystery half-cell.
Ni²⁺ / Ni
(Calculation: E_right – E_standard = -1.05V. So X – 0.80 = -1.05, X = -0.25V).
b) Write the notation for the cell.
Ag(s) | Ag⁺(aq) || Ni²⁺(aq) | Ni(s)
c) Write the ionic equation for the reaction.
Spontaneous reaction:
2Ag⁺ + Ni→2Ag + Ni²⁺
d) State which half-cell is the anode/cathode.
Anode: Ni²⁺/Ni half-cell.
Cathode: Ag⁺/Ag half-cell.
e) State the direction of electron flow.
From the Ni half-cell to the Ag half-cell.
