Using standard reduction potentials (Appendix E), calculate the standard emf for each of the following reactions: (c) Fe1s2 + 2 Fe3+1aq2 ¡ 3 Fe2+1aq2

Verified step by step guidance

1. Identify the half-reactions involved in the overall reaction. For the given reaction, the half-reactions are: \( \text{Fe} \rightarrow \text{Fe}^{2+} + 2e^- \) and \( \text{Fe}^{3+} + e^- \rightarrow \text{Fe}^{2+} \).
2. Look up the standard reduction potentials for each half-reaction from Appendix E. The standard reduction potential for \( \text{Fe}^{3+} + e^- \rightarrow \text{Fe}^{2+} \) is typically given, and you will need to reverse the sign for the oxidation half-reaction \( \text{Fe} \rightarrow \text{Fe}^{2+} + 2e^- \).
3. Calculate the standard emf (electromotive force) for the overall reaction by using the formula: \( E^\circ_{\text{cell}} = E^\circ_{\text{cathode}} - E^\circ_{\text{anode}} \).
4. Substitute the values of the standard reduction potentials into the formula. Remember that the cathode is where reduction occurs and the anode is where oxidation occurs.
5. Sum the potentials to find the standard emf for the overall reaction. Ensure that the stoichiometry of the electrons is balanced in the half-reactions before calculating the emf.

Key Concepts

Standard Reduction Potentials

Standard reduction potentials are measured voltages that indicate the tendency of a chemical species to gain electrons and be reduced. Each half-reaction has a specific potential, typically measured under standard conditions (1 M concentration, 1 atm pressure, and 25°C). These values are crucial for predicting the direction of redox reactions and calculating the overall cell potential.

Electrochemical Cell and EMF

An electrochemical cell consists of two half-cells where oxidation and reduction reactions occur. The electromotive force (EMF) is the voltage generated by the cell, which can be calculated using the standard reduction potentials of the half-reactions involved. The overall EMF can be determined by subtracting the reduction potential of the anode from that of the cathode.

Balancing Redox Reactions

Balancing redox reactions involves ensuring that both mass and charge are conserved in the reaction. This is done by identifying the oxidation and reduction half-reactions, balancing the number of electrons transferred, and adjusting the coefficients of the reactants and products accordingly. Properly balanced reactions are essential for accurate calculations of standard EMF.