Electrolysis Notes

Class Rules

• Raise your hand before asking questions.
• Be active and participative.
• Always pay attention and take notes.

Learning Objectives

• Define electrolysis.
• Outline the electrolysis of molten lead(II) bromide.
• Predict the products of electrolysis for ionic compounds (molten and in solution).

Redox Reactions

• REDOX is short for REDUCTION-OXIDATION.
• In a redox reaction, the atom donating the electrons is oxidized, and the one accepting them is reduced.

Electricity and Electron Flow

• Electricity is the flow of electrons around a circuit from negative (-) to positive (+).

Introduction to Electrochemistry

• Electrochemistry deals with:
  • The chemical changes produced by electric current (in an electrolytic cell).
  • The production of electricity by chemical reactions (in an electrochemical/galvanic cell).

Activity #1

• Research the differences between electrochemical/galvanic cells and electrolytic cells.

Electrochemical/Galvanic Cell vs. Electrolytic Cell

• Electrochemical/Galvanic Cell:
  • Converts chemical energy into electrical energy.
  • The redox reaction is spontaneous and is responsible for the production of electrical energy.
  • The two half-cells are set up in different containers, being connected through a salt bridge or porous partition.
  • The anode is negative, and the cathode is the positive electrode.
  • The reaction at the anode is oxidation, and that at the cathode is reduction.
  • Electrons are supplied by the species getting oxidized; they move from anode to the cathode in the external circuit.
• Electrolytic Cell:
  • Converts electrical energy into chemical energy.
  • The redox reaction is non-spontaneous, and electrical energy has to be supplied to initiate the reaction.
  • Both electrodes are placed in the same container in the solution of molten electrolyte.
  • The anode is positive, and the cathode is the negative electrode.
  • The reaction at the anode is oxidation, and that at the cathode is reduction.
  • The external battery supplies the electrons; they enter through the cathode and come out through the anode.

Spontaneous vs. Non-Spontaneous Reactions

• Non-spontaneous reaction (electrolytic cell): $2 NaCl(l) \rightleftharpoons 2 Na(s) + Cl_2(g)$
• Spontaneous reaction (electrochemical cell).

Electrolysis

• Electrolysis is the decomposition of a compound into its elements by an electric current.
• Uses:
  • Extract metals that are high in the reactivity series (cannot be extracted by heating their ores with carbon).
  • Produce non-metals, such as chlorine.
  • Purify some metals.
• Electrolysis is generally carried out in an electrolytic cell.

Electrolytic Cell Components

• The electrolyte is the compound that is decomposed; it is either a molten ionic compound or a concentrated aqueous solution of ions.
• Power supply must be direct current.
• The electrodes are rods, made from either carbon (graphite) or metal, which conduct electricity to and from the electrolyte.
  • The anode is the positive electrode.
  • The cathode is the negative electrode.

Electrolysis of Molten Lead(II) Bromide

• The molten lead bromide contains lead ions $(Pb^{2+})$ and bromide ions $(Br^-)$.
• Process:
  1. Electrons flow from the negative terminal of the battery to the cathode.
  2. Ions move to the electrode of opposite charge in the liquid.
  3. At the cathode (-), the $(Pb^{2+})$ ions accept electrons. Lead begins to appear below the cathode.
  4. At the anode (+), the $(Br^-)$ ions give up electrons. Red-brown bromine vapor bubbles off.
  5. Electrons flow from the anode to the positive terminal of the battery.
• Overall reaction: $PbBr<em>2(l) \rightarrow Pb(l) + Br</em>2(g)$

Electrolysis of Molten Ionic Compounds

• Metallic cations are reduced at the cathode to form metal.
• Non-metallic anions are oxidized at the anode to form non-metal.

Example: Electrolysis of Molten Sodium Chloride (NaCl)

• At the cathode: silvery beads (sodium) are observed.
• At the anode: bubbles of yellow-green gas (chlorine) are observed.

Exercise #1

• Determine the product observed at both cathode and anode for the electrolysis of:
  • Molten potassium bromide.
  • Molten magnesium oxide.

Electrolysis of Aqueous Ionic Compounds (Solution)

• Solution contains water, which produces ions: $H_2O(l) \rightarrow H^+(aq) + OH^-(aq)$
• At the cathode (check the reactivity series):
  • If the metal is more reactive than H, hydrogen gas forms.
  • If the metal is less reactive than H, metal forms.
• At the anode:
  • If the solution is concentrated and halide (contains halogen ion), halogen gas forms.
  • If the solution is dilute or not halide, oxygen gas forms.

Exercise #2

• Determine the product observed at both cathode and anode for the electrolysis of:
  • Concentrated solution of sulfuric acid, $H<em>2SO</em>4$.
  • Dilute solution of potassium chloride, $KCl$.

Learning Objectives

• Describe how ions move, compete, and react during electrolysis.
• Explain electrolysis as a redox reaction.
• Write half-equations for electrode reactions.

Molten Lead(II) Bromide - Ion Movement and Reactions

• Ions are free to move in molten lead(II) bromide.
• Opposite charges attract; positive lead ions $(Pb^{2+})$ move to the cathode (-), and negative bromide ions $(Br^-)$ move to the anode (+).
• At the cathode (-): Lead ions receive two electrons each and become lead atoms. The half-equation is: $Pb^{2+}(l) + 2e^- \rightarrow Pb(l)$
• At the anode (+): Bromide ions each give up an electron and become atoms, which pair up to form molecules. The half-equation is: $2Br^-(l) \rightarrow Br_2(g) + 2e^-$
• Ions gain electrons: reduction.
• Ions lose electrons: oxidation.
• Overall, electrolysis is a redox reaction; reduction occurs at the cathode, and oxidation occurs at the anode.
• Remember OILRIG: Oxidation Is Loss, Reduction Is Gain.

Electrolysis of Concentrated vs. Dilute Solutions

Concentrated Solution of Sodium Chloride

• Ions from water are also present: $(H^+)$ and $(OH^-)$.
• The solution contains $(Na^+)$ and $(Cl^-)$ ions from the salt and $(H^+)$ and $(OH^-)$ ions from water.
• The positive ions go to the cathode, and the negative ions go to the anode.
• At the cathode, the $(H^+)$ ions accept electrons since hydrogen is less reactive than sodium: $2H^+(aq) + 2e^- \rightarrow H_2(g)$
  • The hydrogen gas bubbles off.
  • The $(Na^+)$ and $(OH^-)$ ions remain, giving a solution of sodium hydroxide.
• At the anode, the $(Cl^-)$ ions give up electrons more readily than the $(OH^-)$ ions do: $2Cl^-(aq) \rightarrow Cl_2(aq) + 2e^-$
  • The chlorine gas bubbles off.
• Remember RAC: Reduction At Cathode.

Dilute Solution of Sodium Chloride

• The same ions are present as before, but the proportion of $(Na^+)$ and $(Cl^-)$ ions is lower.
• At the cathode, hydrogen 'wins' as before and bubbles off: $4H^+(aq) + 4e^- \rightarrow 2H_2(g)$
• At the anode, $(OH^-)$ ions give up electrons since not many $(Cl^-)$ ions are present: $4OH^-(aq) \rightarrow O<em>2(g) + 2H</em>2O(l) + 4e^-$

Half-Equations for Electrode Reactions

• A half-equation shows the electron transfer at an electrode.
• Steps in writing half-equations:
  1. Name the ions present and the products.
  2. Write each half-equation correctly:
     • Give the ion its correct charge.
     • Positive ions go to the cathode, and negative ions to the anode.
     • Write the correct symbol for the element that forms (e.g., Cl for chlorine, not CI).
     • The number of electrons in the equation should be the same as the total charge on the ion(s) in it.
  3. Add the state symbols.

Example

• Electrolysis of molten magnesium chloride:
  • Magnesium ions $(Mg^{2+})$ and chloride ions $(Cl^-)$ are present; Magnesium and chlorine form.
  • Reduction at the cathode: $Mg^{2+} + 2e^- \rightarrow Mg$ or $Mg^{2+}(l) + 2e^- \rightarrow Mg(s)$
  • Oxidation at the anode: $2Cl^- \rightarrow Cl<em>2 + 2e^-$ or $2Cl^-(l) \rightarrow Cl</em>2(g) + 2e^-$