E o cell = E o red - E o ox. Oxidizing agent must be higher on the table than the reducing agent in order for the reaction to be spontaneous

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1 Chemistry 12 Electrochemistry III Name: Date: Block: 1. SRP Table Standard Reduction Potential (SRP) Table This table is used to predict whether a chemical species will spontaneously give electrons to or take electrons from another species. Oxidizing agents (take electrons) are on the side Reducing agents (give electrons) are on the side Oxidizing agents will only take electrons spontaneously from species below them in the right column. Reducing agents will only give electrons spontaneously to species above them in the left column. **Oxidizing agent must be higher on the table than the reducing agent in order for the reaction to be spontaneous** Example: Fe 3+ and Cl - Locate each species on the SRP table. Fe 3+ is a(n) agent. Cl- is a(n) agent. Which is higher? Will the reaction occur spontaneously? Write the balanced equation for the reaction. If two half- reactions can be added together to give a REDOX reaction, the voltages associated with the half- reactions can also be added. E o red = the reduction potential of the reduction reaction E o ox = the oxidation potential of the oxidation reaction E o cell = E o red - E o ox (Don t worry! There s an easier way to calculate it without having to memorize the formula!)

2 Calculate the cell potential of Identify (and balance) the two half reactions: Ag + + Al à Ag + Al 3+ ü Note 1: Since voltage is the work done per electron, do not multiple the E o value for the reduction of Ag + by 3. ü Note 2: If a half- reaction is reversed, the sign of its E o value is also reversed. What is the total cell potential? Is it spontaneous? Determine whether the following reactions will occur spontaneously. If so, write the balanced equation and calculate E o cell.. Ca and Zn 2+ I - and Al I - and Br 2 F 2 and Al 3+ Ag + and Sn I 2 and Cl -

3 The Strength of Oxidizing and Reducing Agents Will chlorine and sodium bromide spontaneously react to produce bromine and sodium chloride? If so, write the balanced net ionic equation for the reaction. Balanced formula equation: Complete ionic equation: Net ionic equation: Identify oxidizing agent and reducing agent: Will the reaction spontaneously occur? For each of the following, determine whether a reaction will occur and if so, write the balanced net ionic equation for the reaction and calculate E o cell.. 1. I 2 + CaF 2 2. Al + CuSO 4 3. Br 2 + NaCl 4. Sn + Al(NO 2) 3

4 Similarly to BL- acids and bases, oxidizing and reducing agents vary in strengths as well. The stronger a reducing agent (A- ) is, the weaker its complementary oxidizing agent (A) is. If a chemical species has a strong tendency to give an electron away, then it will have a weak tendency to take it back. The strength of an oxidizing agent is called its (E o ) or potential to be reduced and gain an electron. Oxidation potential measures the strength of reducing agents. When the reduction half- reactions is read backwards, they are oxidation half reactions. In other words, oxidation potentials of reducing agents are simply the opposite of the reduction potentials of their complementary oxidizing agents. Example: 1. Which is the stronger reducing agent, Co or Sr? 2. Which is the stronger oxidizing agent, Fe 3+ or Al 3+? 3. Which has the greater oxidation potential, Br - or I -? 4. Which has the greater reduction potential, Cr 3+ or Sn 2+? 5. What is the reduction potential of Ag +? 6. What is the oxidation potential of Ca? 7. A and B are hypothetical elements. A 2+ is a stronger oxidizing agent than B. Which has the greater oxidation potential, A or B -? When there are more than two chemicals available to react, the SRP table is used to predict which species will react. The predominant redox reaction will be between the strongest available oxidizing agent (highest on the left side of the table) and the strongest available reducing agent (lowest on the right side of the table).

5 Practice: 1. What will the predominant redox reaction be with a mixture of Cl 2, Ag +, Sn 2+ and I -? Identify the oxidizing agent(s) and the reducing agent (s). Use the SRP table to determine which species is the strongest oxidizing agent and which species is the strongest reducing agent. Write the two half- reactions and balance the transfer of electrons. Add the half- reactions together. 2. Sn 4+, Br -, Zn 2+ and Ag 3. CuBr 2 (aq) and Al (s) 4. Na +, Cu + and F - 5. Copper and bromine in a solution of iron (III) iodide Hebden Workbook Pg. 224 #36 38

1.7 REDOX. Convert these to ionic and half equations and you can see clearly how the electrons are transferred:

1.7 REDOX. Convert these to ionic and half equations and you can see clearly how the electrons are transferred: 1.7 REDOX Oxidation and Reduction: Oxidation and reduction reactions can be identified by looking at the reaction in terms of electron transfer: Our understanding of oxidation and reduction was limited