If ${\mathrm{E}}_{{\mathrm{Fe}}^{2+}/\mathrm{Fe}}^{\mathrm{o}}$ = -0.441 V and  ${\mathrm{E}}_{{\mathrm{Fe}}^{3+}/{\mathrm{Fe}}^{2+}}^{\mathrm{o}}$ = 0.771 V, the standard emf of the reaction: 

Fe + 2Fe³⁺→ 3Fe²⁺ will be:

Electrode potential is the potential difference between the -

1. Electrode and the electrolyte.

2. Anode and Cathode.

3. Anode and Atmosphere.

4. Cathode and Atmosphere.

The cell that will measure the standard electrode potential of a copper electrode is:

The standard electrode potential for Sn⁴⁺/Sn²⁺ couple is +0.15 V and that for the Cr³⁺/Cr couple is -0.74 V. These two couples in their standard state are connected to make a cell. The cell potential will be: 

The standard reduction potential for Fe²⁺|Fe and Sn²⁺|Sn electrodes are -0.44 V and -0.14 V respectively. For the cell reaction,

Fe²⁺ + Sn  →  Fe + Sn²⁺, the standard Emf is - 

1. +0.30 V

2. 0.58 V

3. +0.58 V

4. -0.30 V

If a salt bridge is removed from the two half cells, the voltage:

1. Drops to zero.

2. Does not change.

3. Increases gradually.

4. Increases rapidly.

The cell potential will be:

A hypothetical electrochemical cell is shown below.
A|A⁺(x M) || B⁺(y M)|B
The Emf measured is +0.20 V. The cell reaction is:

1. A⁺ + B → A + B⁺

2.  A⁺ + e^- → A ; B⁺ + e^- → B

3. The cell reaction cannot be predicted.

4. A + B⁺ → A⁺ + B

The difference between the electrode potentials of two electrodes when no current is drawn through the cell is called:
1. Cell potential.
2. Cell emf.
3. Potential difference.
4. Cell voltage.

Consider the following relations for emf of an electrochemical cell:

The correct relation among the given options is: