Help with Electrochemistry Lab: Al+3, Cu+2, Fe+3, Zn+2, KNO3

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SUMMARY

The discussion focuses on an electrochemistry lab involving aluminum ions (Al3+), copper ions (Cu2+), iron ions (Fe3+), zinc ions (Zn2+), and potassium nitrate (KNO3). Participants utilized filter paper soaked in KNO3 as a salt bridge to facilitate redox reactions between various ion combinations. The confusion arises regarding the identification of oxidation and reduction reactions, particularly when both half-cells contain positive ions. The half-reactions and their standard electrode potentials were provided, but clarification on the complete cell reactions and the role of the salt bridge was requested.

PREREQUISITES
  • Understanding of electrochemical cells and redox reactions
  • Familiarity with standard electrode potentials
  • Knowledge of half-reaction notation
  • Basic principles of using a salt bridge in electrochemistry
NEXT STEPS
  • Research the complete electrochemical cell reactions for Al3+, Cu2+, Fe3+, and Zn2+
  • Learn about the function and significance of salt bridges in electrochemical cells
  • Study the Nernst equation and its application in calculating cell potentials
  • Explore the interactions between nitrate ions and metal ions in electrochemical contexts
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Students and educators in chemistry, particularly those focusing on electrochemistry, laboratory technicians conducting redox experiments, and researchers interested in ion interactions in electrochemical systems.

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I'm very confused with an electrochemistry lab. For the lab, we used Al+3, Cu+2, Fe+3, Zn+2, and KNO3. The setup was wells with all of the solutions in them. We soaked a piece of paper(filter paper) in KNO3 for the reactions, then used the paper as a salt bridge between the solutions of the ions for all combinations(i.e. Cu+2 with Fe+3, Fe+3 with Zn+2 etc.). What I'm confused with is how to write the reactions as both cells have positive ions in them, making them both reduction reactions, unless I'm missing something horribly wrong. Can someone help me? I have no clue what the oxidation reactions are. This is what I have so far:
<br /> \subsection{\ce{Al^{+3}} and \ce{Cu^{+2}}}<br /> \paragraph{\ce{Cu^{+2}} half reaction}<br /> \ce{Cu^{+2} +2e^{-} -&gt; Cu} .34V<br /> \paragraph{\ce{Al^{+3}} half reaction}<br /> \ce{Al^{+3} + 3e^{-} -&gt; Al} -1.66V <br /> \subsection{\ce{Cu^{+2}} and \ce{Fe^{+3}}}<br /> \paragraph{\ce{Cu^{+2}} half reaction}<br /> \ce{Cu^{+2} +2e^{-} -&gt; Cu} .34V<br /> \paragraph{\ce{Fe^{+3}} half reaction}<br /> \ce{Fe^{+3} + e^{-} -&gt; Fe^{+2}} .77V<br /> \subsection{\ce{Fe^{+3}} and \ce{Zn^{+2}}}<br /> \paragraph{\ce{Fe^{+3}} half reaction}<br /> \ce{Fe^{+3} + e^{-} -&gt; Fe^{+2}} .77V<br /> \paragraph{\ce{Zn^{+2}} half reaction}<br /> \ce{Zn^{+2} + 2e^{-} -&gt; Zn} -.76V<br /> \subsection{\ce{Al^{+3}} and \ce{Fe^{+3}}}<br /> \paragraph{\ce{Al^{+3}} half reaction}<br /> \ce{Al^{+3} + 3e^{-} -&gt; Al} -1.66V <br /> \paragraph{\ce{Fe^{+3}} half reaction}<br /> \ce{Fe^{+3} + e^{-} -&gt; Fe^{+2}} .77V<br /> \subsection{\ce{Al^{+3}} and \ce{Zn^{+2}}}<br /> \paragraph{\ce{Al^{+3}} half reaction}<br /> \ce{Al^{+3} + 3e^{-} -&gt; Al} -1.66V <br /> \paragraph{\ce{Zn^{+2}} half reaction}<br /> \ce{Zn^{+2} + 2e^{-} -&gt; Zn} -.76V<br /> \subsection{\ce{Cu^{+2}} and \ce{Zn^{+2}}}<br /> \paragraph{\ce{Cu^{+2}} half reaction}<br /> \ce{Cu^{+2} +2e^{-} -&gt; Cu} .34V<br /> \paragraph{\ce{Zn^{+2}} half reaction}<br /> \ce{Zn^{+2} + 2e^{-} -&gt; Zn} -.76V<br />
 
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You seem to have not stated what counter half cell was used for each have reaction which you listed. You would not expect much with just a half cell plus a salt bridge. Did you have reactions between the Nitrate and a metal ion?
 

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