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Why wouldn't half-reactions in a battery just go on?

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  1. Jan 12, 2016 #1
    Have been trying to better understand how batteries work. Forgive me if the question's naive.

    A half-reaction on one of the electrodes in a battery produces free electrons (for example) and consumes anions (or produces cations). I understand why, if there's an external wire for electrons to travel on, there needs also to be a salt bridge for ions to travel on inside the battery, and if there isn't, the electrons will soon stop flowing due to separation of charge building up.

    However, since each half-reaction on its own is charge-balanced, what stops the half-reaction from just continuing to happen even when the circuit is open? Why for example in an alkaline battery Zn won't continue to react with OH- abundantly coming from the electrolyte?
     
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  3. Jan 12, 2016 #2

    462chevelle

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    Without the wire would you have electrons flowing from one to the other?
     
  4. Jan 12, 2016 #3

    Borek

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    Your question isn't naive at all. Take a look how the batteries are designed - half of the effort is a selection of the correct half reactions, the other half is the construction that allows the charge to flow without allowing reagents to get in a direct contact.
     
  5. Jan 14, 2016 #4

    James Pelezo

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    Here's a post made Dec 6th on Galvanic cells that might help understand the issues...

    I'll probably be castigated for this, but there seems to be a considerable misunderstanding in the function and chemistry of electrochemical cells, when they are actually very simple... Please forgive my being wordy, but in the spirit of supporting the forum an chemical education, I would like to add my interpretation of basic concepts about electrolysis and galvanic processes for those who would wish to know how such systems function. It does not presume to include all details of explaining electrochemical processes; just basic concepts.

    For both Electrolytic and Galvanic/Voltaic Cells the electrodes in electrochemical cells are defined in terms of the chemistry that takes place at a specified electrode; i.e., 'Reduction Rxn' = Cathode and 'Oxidation Rxn' = Anode.

    Galvanic Cells:
    For instructional purposes, there are two Galvanic processes; 'Controlled Galvanic Process' and 'Uncontrolled Galvanic Process'. The Uncontrolled Galvanic Process is one in which oxidation and reduction reactions occur simultaneously in one cell. Example: Given a container of Copper(II) Sulfate ions in solution and inserting a Zinc metal bar directly into the solution results in reduction of Cu+2(aq) ions to Cuo(s) sticking to the surface of the zinc bar. Over time, the copper sticking to the zinc bar will form a coating that will prevent further reduction of copper ions to copper metal and the Galvanic/Voltaic process stops. For a controlled Galvanic/Voltaic process, the chemical processes are separated into anodic and cathodic cells allowing the system to sustain a continuous current of electric charge flow until the anodic material completely dissolves into solution leading to a 'dead battery'.

    Example of Controlled Zn/Cu Galvanic/Voltaic process (Refer to the diagram at the beginning of this thread):
    The Galvanic/Voltaic Cell for the copper/zinc system is a spontaneous reaction process giving current flow when connected. Copper ions in solution are reduced at the copper bar electrode by gaining (reduction) 2 electrons... Cu+2 + 2e- => Cuo(s) leaving the copper electrode deficient in electrons => positive electrode (cathode). In the zinc cell side, The Zinc electrode is being oxidized to Zinc(II) ions that are delivered into solution. Zno(s) => Zn+2 + 2e-. This oxidation half reaction leaves excess electrons in the Zinc bar and => negative electrode (anode). The voltaic cell will discharge until all of the anode is dissolved and no oxidation half reactions occur and the cell is a 'dead battery'.

    As for the salt bridge, its function is to maintain balance of charge as the Galvanic process proceeds. In the anode side of the cell where oxidation is occurring, there is an increase in positive charge due to the cations being delivered into solution. The Negative ions of the salt bridge therefore migrate to the anodic cell to neutralize the build up of positive charge. The Positive ions of the salt bridge migrate to the cathode cell to replace the positive charge loss when cations in solution are reduced to neutral causing a loss of positive charge in the cathodic cell solution.

    Electrolytic Cells:
    The electrodes in the Electrolytic Cell are defined in the same way; i.e., Oxidation => Anode & Reduction => Cathode. The difference is the chemistry of the ions in the solution cause the anode to assume a positive charge and the cathode to assume a negative charge; opposite that of the Galvanic Process. Example: NaCl(melt) => Na+(l) + Cl-(l). The Electrolytic Cell is non-spontaneous and requires an outside potential to drive the reactions and is therefore connected to a Galvanic Cell (battery) giving one electrode in the electrolytic cell a positive charge and the other a negative charge. The Na+ ions migrate to the electrode connected to the (-) electrode of the battery and undergo reduction by gaining an electron (Na+ + e- => Nao(s) and the chloride ions undergo oxidation to give chlorine gas (Cl2) ... 2Cl- + 2e- => Cl2(g). Commercially this is referred to as a Downs Cell.
     
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