Electrolysis and Electrochemical Cells Flashcards

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Electrolysis

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Electrolysis is the process of using electrical energy to break down or decompose a compound, usually an ionic compound in the molten or aqueous state. It occurs in an electrolytic cell where ions are discharged at electrodes to form elements or new substances.

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Electrolytic cell

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An electrolytic cell is the apparatus where electrolysis occurs and contains two electrodes, an electrolyte, and an external power source. The cell drives non-spontaneous reactions by forcing electrons to flow from the anode to the cathode.

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Battery role

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The battery (or power source) acts as an electron pump, removing electrons from the anode and supplying them to the cathode. This makes the anode positive and the cathode negative, driving ion discharge at the electrodes.

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Electrodes

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There are two electrodes: the anode (positive) which attracts anions and where oxidation occurs, and the cathode (negative) which attracts cations and where reduction occurs. Electrodes are often inert (graphite, platinum) or reactive metals depending on the process.

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Electrolyte

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An electrolyte is a molten ionic compound or aqueous solution that conducts electricity by containing free-moving ions. During electrolysis the electrolyte provides the cations and anions that are discharged at the electrodes.

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Oxidation and reduction

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Oxidation is the loss of electrons and occurs at the anode, while reduction is the gain of electrons and occurs at the cathode. During electrolysis ions are discharged by undergoing oxidation or reduction to form neutral atoms or molecules.

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Molten electrolysis

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In molten electrolysis the ionic compound is melted so ions are free to move and the elements can be separated, for example molten $NaCl$ contains $Na^+$ and $Cl^-$. At the anode $Cl^-$ is oxidised to $Cl_2$ gas and at the cathode $Na^+$ is reduced to $Na$ metal.

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Half-equation

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A half-equation shows the reaction at one electrode and includes electrons explicitly, with ions on the left and products on the right. For molten $NaCl$ the anode half-equation is $2Cl^- (l) \rightarrow Cl_2 (g) + 2e^-$ and the cathode half-equation is $Na^+ (l) + e^- \rightarrow Na (l)$.

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Overall equation

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To form the overall equation you balance electrons between half-equations, add left and right sides and cancel electrons. For molten $NaCl$ the balanced overall equation is $2NaCl (l) \rightarrow 2Na (l) + Cl_2 (g)$.

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Aqueous electrolysis

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In aqueous electrolysis water partially ionises to give $H^+$ and $OH^-$, so the electrolyte contains those ions plus any ions from the dissolved salt. Only one ion at each electrode is selectively discharged, determined by the electrochemical series and ion concentrations.

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Electrochemical series

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The electrochemical series ranks cations by how easily they are reduced; for cations an order is $K^+, Na^+, Ca^{2+}, Mg^{2+}, Zn^{2+}, Fe^{2+}, Pb^{2+}, H^+, Cu^{2+}, Ag^+$. Ions lower in the series are discharged preferentially over ions higher up, and concentration can alter which anion is discharged.

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Halide discharge

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For anions the order includes $SO_4^{2-}, NO_3^-, Cl^-, Br^-, I^-, OH^-$. In concentrated halide solutions (e.g., concentrated $NaCl$) the halide ion can be discharged in preference to $OH^-$ because of its higher concentration despite similar positions in the series.

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H+/OH- half-equations

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The common half-equations for aqueous electrolysis are $2H^+ (aq) + 2e^- \rightarrow H_2 (g)$ at the cathode and $4OH^- (aq) \rightarrow 2H_2O (l) + O_2 (g) + 4e^-$ at the anode. These are used to derive overall reactions when $H^+$ and $OH^-$ are the species discharged.

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Final electrolyte

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The final electrolyte is determined by the ions left in solution after discharge. For example, dilute $NaCl$ electrolysis leaves $Na^+$ and $Cl^-$ so the final electrolyte remains $NaCl (aq)$, while concentrated $NaCl$ leaves $Na^+$ and $OH^-$ giving $NaOH (aq)$, and dilute $CuSO_4$ can leave $H^+$ and $SO_4^{2-}$ producing $H_2SO_4 (aq)$.

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Reactive anode

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A reactive anode is made of a metal (e.g., Cu or Ag) that can oxidise and enter solution rather than producing anion discharge. For instance a copper anode will oxidise via $Cu (s) \rightarrow Cu^{2+} (aq) + 2e^-$, causing the anode to lose mass.

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Electrolytic purification

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Electrolytic purification uses an impure metal as the anode and pure metal as the cathode in a solution of the metal's salt, so the metal from the anode dissolves and pure metal plates onto the cathode. Impurities either remain in solution or fall off as anode sludge, producing a purified metal deposit.

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Electroplating

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Electroplating deposits a thin layer of metal onto an object by using the metal's salt as the electrolyte, the plating metal as the anode, and the object as the cathode. Metal ions at the anode oxidise to $M^{n+}$ and those ions are reduced at the cathode to form a metal coating, e.g., $Ag (s) \rightarrow Ag^+ (aq) + e^-$ and $Ag^+ (aq) + e^- \rightarrow Ag (s)$.

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Simple cell

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A simple cell converts chemical energy to electrical energy by placing two different metals in contact with an electrolyte; electrons flow from the more reactive metal to the less reactive metal. The more reactive metal becomes the anode and loses mass as it oxidises, and the larger the reactivity difference the larger the voltage produced.

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Simple cell equations

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In a simple cell the anode metal always ionises, e.g., $Zn (s) \rightarrow Zn^{2+} (aq) + 2e^-$. With a dilute $H_2SO_4$ electrolyte the cathode reaction is $2H^+ (aq) + 2e^- \rightarrow H_2 (g)$, giving the overall $Zn (s) + 2H^+ (aq) \rightarrow Zn^{2+} (aq) + H_2 (g)$.

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Hydrogen fuel cell

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A hydrogen fuel cell continuously converts chemical energy to electricity using $H_2$ and $O_2$ with an alkaline electrolyte. The anode half-reaction is $2H_2 (g) + 4OH^- (aq) \rightarrow 4H_2O (l) + 4e^-$, the cathode half-reaction is $O_2 (g) + 2H_2O (l) + 4e^- \rightarrow 4OH^- (aq)$, and the overall reaction is $O_2 (g) + 2H_2 (g) \rightarrow 2H_2O (l)$.

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Fuel cell advantages

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Hydrogen fuel cells are more efficient than combustion because a greater fraction of chemical energy becomes useful electricity, and their only direct product is water. Hydrogen can be produced renewably (e.g., electrolysis), making the system potentially low-carbon.

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Fuel cell disadvantages

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Hydrogen is difficult and expensive to produce, store, and transport safely because it is highly flammable and often requires high-pressure cylinders. Also producing hydrogen typically requires energy that may come from polluting sources, so fuel cells are not inherently pollution-free.