Using data in Appendix E, calculate the standard emf for each of the following reactions: (a) H2(g) + F2(g) → 2 H+(aq) + 2 F-(aq) (b) Cu2+(aq) + Ca(s) → Cu(s) + Ca2+(aq) (c) 3 Fe2+(aq) → Fe(s) + 2 Fe3+(aq) (d) 2 ClO3-(aq) + 10 Br-(aq) + 12 H+(aq) → Cl2(g) + 5 Br2(l) + 6 H2O(l)

Verified step by step guidance

Step 1: Identify the half-reactions for each given reaction and write them down. For example, for reaction (a), identify the oxidation and reduction half-reactions.

Step 2: Use Appendix E to find the standard reduction potentials (E°) for each half-reaction. Remember that the standard reduction potential is given for the reduction process, so you may need to reverse the sign for oxidation reactions.

Step 3: Calculate the standard emf (E°cell) for each reaction by using the formula E°cell = E°cathode - E°anode. Ensure you correctly identify which half-reaction is the cathode (reduction) and which is the anode (oxidation).

Step 4: For reactions involving multiple electrons, ensure that the number of electrons lost in the oxidation half-reaction equals the number of electrons gained in the reduction half-reaction. You may need to multiply the half-reactions by appropriate coefficients to balance the electrons.

Step 5: Sum the balanced half-reactions to obtain the overall balanced redox reaction, and verify that the calculated E°cell corresponds to the balanced reaction.

Key Concepts

Here are the essential concepts you must grasp in order to answer the question correctly.

Standard Electrode Potential (E°)

The standard electrode potential (E°) is a measure of the tendency of a chemical species to be reduced, measured under standard conditions (1 M concentration, 1 atm pressure, and 25°C). It is expressed in volts and is crucial for calculating the electromotive force (emf) of electrochemical cells. Each half-reaction in a redox process has a specific E° value, which can be found in standard tables.

Nernst Equation

The Nernst equation relates the cell potential to the concentrations of the reactants and products in a redox reaction. It allows for the calculation of the emf under non-standard conditions by incorporating the reaction quotient (Q). The equation is given by E = E° - (RT/nF)ln(Q), where R is the gas constant, T is the temperature in Kelvin, n is the number of moles of electrons transferred, and F is Faraday's constant.

Balancing Redox Reactions

Balancing redox reactions involves ensuring that both mass and charge are conserved in the reaction. This process typically includes identifying oxidation and reduction half-reactions, balancing the number of electrons transferred, and adjusting coefficients to achieve equal numbers of atoms and charges on both sides of the equation. Properly balanced reactions are essential for accurate emf calculations.

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Balancing Basic Redox Reactions

Related Practice

Textbook Question

Using standard reduction potentials (Appendix E), calculate the standard emf for each of the following reactions: (b) Ni1s2 + 2 Ce4+1aq2 ¡ Ni2+1aq2 + 2 Ce3+1aq2

Textbook Question

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

Textbook Question

Using standard reduction potentials (Appendix E), calculate the standard emf for each of the following reactions: (d) 2 NO3-1aq2 + 8 H+1aq2 + 3 Cu1s2 ¡ 2 NO1g2 + 4 H2O1l2 + 3 Cu2+1aq2

Textbook Question

The standard reduction potentials of the following halfreactions are given in Appendix E:

Ag⁺(aq) + e^- → Ag(s)

Cu²⁺(aq) + 2 e^- → Cu(s)

Ni²⁺(aq) + 2 e^- → Ni(s)

Cr³⁺(aq) + 3 e^- → Cr(s)

(a) Determine which combination of these half-cell reactions leads to the cell reaction with the largest positive cell potential and calculate the value.

(b) Determine which combination of these half-cell reactions leads to the cell reaction with the smallest positive cell potential and calculate the value.

Textbook Question

A voltaic cell consists of a strip of cadmium metal in a solution of Cd(NO₃)₂ in one beaker, and in the other beaker a platinum electrode is immersed in a NaCl solution, with Cl₂ gas bubbled around the electrode. A salt bridge connects the two beakers. (a) Which electrode serves as the anode, and which as the cathode? (b) Does the Cd electrode gain or lose mass as the cell reaction proceeds? (c) Write the equation for the overall cell reaction.