Electrochemistry Lecture -19 PDF

Chapter 19 Lecture presentation class 1 Electrochemistry © 2015 Pearson Education, Ltd.  Redox Reactions Galvanic Cells Standard Reduction Potentials Batteries Corrosion Electrolysis Electrometallurgy...

Chapter 19 Lecture presentation class 1 Electrochemistry © 2015 Pearson Education, Ltd.  Redox Reactions Galvanic Cells Standard Reduction Potentials Batteries Corrosion Electrolysis Electrometallurgy © 2015 Pearson Education, Ltd.  Oxidation–Reduction (Redox) Reactions in which electrons are transferred from one atom to another are called oxidation–reduction reactions. – Redox reactions for short Atoms that lose electrons are being oxidized; atoms that gain electrons are being reduced. 2 Na(s) + Cl2(g) → 2 Na+Cl–(s) Na → Na+ + 1 e– oxidation Cl2 + 2 e– → 2 Cl– reduction © 2015 Pearson Education, Ltd.  Oxidation and Reduction Oxidation is the process that occurs when – the oxidation number of an element increases; – an element loses electrons; – a compound adds oxygen; – a compound loses hydrogen; or – a half-reaction has electrons as products. Reduction is the process that occurs when – the oxidation number of an element decreases; – an element gains electrons; – a compound loses oxygen; – a compound gains hydrogen; or – a half-reaction has electrons as reactants. © 2015 Pearson Education, Ltd.  Electrochemistry Electrochemistry is the study of redox reactions that produce or require an electric current. The conversion between chemical energy and electrical energy is carried out in an electrochemical cell. Spontaneous redox reactions take place in a voltaic cell. – Also known as a galvanic cell Nonspontaneous redox reactions can be made to occur in an electrolytic cell by the addition of electrical energy. © 2015 Pearson Education, Ltd.  Redox Reactions by the Half-Reaction © 2015 Pearson Education, Ltd.  Galvanic cell or Voltaic cell The experimental apparatus for generating electricity through the use of a spontaneous reaction is called a galvanic cell or voltaic cell, after the Italian scientists Luigi Galvani and Alessandro Volta, who constructed early versions of the device. © 2015 Pearson Education, Ltd. © 2015 Pearson Education, Ltd.  A Spontaneous Redox Reaction: Zn(s) + Cu2+(aq) → Zn2+ + Cu(s) © 2015 Pearson Education, Ltd.  Electrochemical Cells Oxidation and reduction half-reactions are kept as separate in half-cells in an electrochemical cell. To constitute an electrical circuit: Electron flow through a wire along with Ions (electrolyte) flowing through a solution via the salt bridge. The flow of electrons require a conductive solid electrode to allow the transfer of electrons either through: An external circuit or Metal or graphite electrode An electrochemical cell requires the exchange of ions between the two half-cells of the system via a salt bridge. © 2015 Pearson Education, Ltd.  Voltaic (Galvanic) Cells: Spontaneous Redox Reactions Electrical current: The amount of electric charge that passes a point in a given period of time – Whether as electrons flowing through a wire or as ions flowing through a solution Redox reactions involve the movement of electrons from one substance to another. – Therefore, redox reactions have the potential to generate an electric current. Voltaic (galvanic) cells produce an electrical current from spontaneous redox reactions. – To use that current, we need to separate the place where oxidation is occurring from the place where reduction is occurring. © 2015 Pearson Education, Ltd.  Electrodes of an Electrochemical Cell Anode – Electrode where oxidation occurs – Anions attracted to it – Connected to positive end of battery in an electrolytic cell – Loses weight in electrolytic cell Cathode – Electrode where reduction occurs – Cations attracted to it – Connected to negative end of battery in an electrolytic cell – Gains weight in electrolytic cell Electrode where plating takes place in electroplating © 2015 Pearson Education, Ltd.  Voltaic Cell The salt bridge is required to complete the circuit and maintain charge balance and cell neutrality. © 2015 Pearson Education, Ltd.  Galvanic Cells Anode: – The electrode where oxidation occurs – The electrode where electrons are produced – Is what anions migrate toward – Has a negative sign © 2015 Pearson Education, Ltd.  Galvanic Cells Cathode: – The electrode where reduction occurs – The electrode where electrons are consumed – Is what cations migrate toward – Has a positive sign © 2015 Pearson Education, Ltd.  Cell Notation Shorthand description of a voltaic cell is written as follows: electrode | electrolyte || electrolyte | electrode – Oxidation half-cell on the left; reduction half-cell on the right – Single | = phase barrier If multiple electrolytes in same phase, a comma is used rather than | Often use an inert electrode – Double line || = salt bridge Example: Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s) © 2015 Pearson Education, Ltd.  Voltaic Cell Anode = Zn(s) Cathode = Cu(s) The anode is oxidized to Cu2+ ions are reduced at Zn2+ ions. the cathode. Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s) © 2015 Pearson Education, Ltd.  https://www.youtube.com/watch?v=qpFC_Ecu_yQ © 2015 Pearson Education, Ltd.  Electrodes Typically, – the anode is made of the metal that is oxidized; and – the cathode is made of the same metal as is produced by the reduction. If the redox reaction we are running involves the oxidation or reduction of an ion to a different oxidation state, or the oxidation or reduction of a gas, we may use an inert electrode. – An inert electrode is one that does not participate in the reaction but just provides a surface for the transfer of electrons to take place on. © 2015 Pearson Education, Ltd.  Which Way Will Electrons Flow? Under standard conditions, zinc has a stronger tendency to oxidize than copper. – Electrons flow from anode to cathode. – Therefore, the electrons flow from zinc, making zinc the anode. Zn → Zn2+ + 2 e− E° = +0.76 Cu → Cu2+ + 2 e− E° = −0.34 © 2015 Pearson Education, Ltd.  Consider the following reaction: 2Ag+(aq)+Cu(s)⇌Cu2+(aq)+2Ag(s) Write the two half-reactions. Ag+(aq)+e−⇌Ag(s) Reduction Cu(s)⇌Cu2+(aq)+2e− Oxidation Identify the cathode and anode. Cu(s) is losing electrons thus being oxidized; oxidation occurs at the anode. Anode (where oxidation occurs): Ag+ is gaining electrons and thus is being reduced; reduction happens at the cathode. Cathode (where reduction occurs Construct the Cell Diagram. Cu(s)|Cu2+(aq)||Ag+(aq)| Ag(s) © 2015 Pearson Education, Ltd.  Consider the following two reactions:  Cu2+(aq)+Ba(s)→Cu(s)+Ba2+(aq)  2Al(s)+3Sn2+(aq)→2Al3+(aq)+3Sn(s) Split the reaction into half reactions a) Ba2+(aq) → Ba(s) + 2e- (Anode; where oxidation happens) Cu2+(aq) + 2e- → Cu(s) (Cathode; where reduction happens) b) Al3+(aq) → Al(s) + 3e- (Anode; where oxidation happens) Sn2+(aq) +2e- → Sn(s) (Cathode; where reduction happens) Construct a cell diagram for the following each reactions Ba2+(aq) | Ba(s) || Cu(s) | Cu2+(aq) Al(s) | Al3+(aq) || Sn2+(aq) | Sn(s) © 2015 Pearson Education, Ltd.  Consider a galvanic cell consisting of 2Cr(s)+3Cu2+(aq)⟶2Cr3+(aq) + 3Cu(s) Write the oxidation and reduction half-reactions and write the reaction using cell notation. Which reaction occurs at the anode? The cathode? By inspection, Cr is oxidized when three electrons are lost to form Cr3+, and Cu2+ is reduced as it gains two electrons to form Cu. Balancing the charge gives oxidation:2Cr(s)⟶2Cr3+(aq)+6e− reduction:3Cu2+(aq)+6e−⟶3Cu(s) Cell notation uses the simplest form of each of the equations, Cr(s)∣Cr3+(aq)∥Cu2+(aq)∣Cu(s)Cr © 2015 Pearson Education, Ltd.  Some oxidation- reduction reactions involve species that are poor conductors of electricity, and so an electrode is used that does not participate in the reactions. Frequently, the electrode is platinum, gold, or graphite, all of which are inert to many chemical reactions. © 2015 Pearson Education, Ltd.  Mg(s)+2H+(aq)⟶Mg2+(aq) +H2(g) oxidation: Mg(s)⟶Mg2+(aq) +2e− reduction:2H+(aq) +2e−⟶H2(g) The cell used an inert platinum wire for the cathode, so the cell notation is Mg(s)∣Mg2+(aq)∥H+(aq)∣H2(g)∣Pt(s) © 2015 Pearson Education, Ltd.  Write the following balanced reactions using cell notation. Use platinum as an inert electrode 3CuNO3(aq) 3CuNO3(aq)+Au(NO3)3(aq)⟶3Cu(NO3)2(aq) +Au(NO3)3(aq)⟶3Cu(NO3)2(aq)+Au(s +Au(s) 3 Cu0 - 6 e- → 3 CuII (oxidation) 2 AuIII + 6 e- → 2 Au0 (reduction) Pt(s)∣Cu+(aq),Cu2+(aq)∥Au3+(aq)∣Au(s) © 2015 Pearson Education, Ltd.  Describe in shorthand notation a galvanic cell for which the cell reaction is Cu(s)+2Fe3+(aq)→Cu2+(aq)+2Fe2+(aq) Oxidation: Cu(s)→Cu2+(aq)+2e– Reduction: Fe3+ +e–→Fe2+ Cu∣Cu2+∥Fe2+,Fe3+∣Pt Since both Fe2+ and Fe3+ are in solution, a Pt electrode is used.) © 2015 Pearson Education, Ltd.  Batteries: Leclanché Acidic Dry Cell Electrolyte in paste form – ZnCl2 + NH4Cl or MgBr2 Expensive, nonrechargeable, heavy, easily corroded © 2015 Pearson Education, Ltd.  Batteries: Alkaline Dry Cell Same basic cell as acidic dry cell, except electrolyte is alkaline KOH paste Longer shelf life than acidic dry cells and rechargeable, with little corrosion of zinc. © 2015 Pearson Education, Ltd.  Batteries: Lead Storage Battery Six cells in series Electrolyte = 30% H2SO4 Rechargeable, heavy © 2015 Pearson Education, Ltd.  NiCad Battery Electrolyte is concentrated KOH solution. Rechargeable, long life, light; however, recharging incorrectly can lead to battery breakdown. © 2015 Pearson Education, Ltd.  Lithium Ion Battery Electrolyte is concentrated KOH solution. Rechargeable, long life, very light, more environmentally friendly, greater energy density © 2015 Pearson Education, Ltd.  Fuel Cells They are like batteries in which reactants are constantly being added. – So they never run down! Anode and cathode are both Pt coated metal. Electrolyte is OH– solution. Anode reaction: 2 H2 + 4 OH– → 4 H2O(l) + 4 e– Cathode reaction: O2 + 4 H2O + 4 e– → 4 OH– © 2015 Pearson Education, Ltd.  Electrochemical Cells Overview In all electrochemical cells, oxidation occurs at the anode and reduction occurs at the cathode. In voltaic cells (spontaneous reactions; E°cell is positive), – the anode is the source of electrons and has a (−) charge; – the cathode draws electrons and has a (+) charge. In electrolytic cells (non spontaneous reactions; E°cell is negative), © 2015 Pearson Education, Ltd.  Standard reduction potential (E° or E° red): Reduction potential for a half-reaction under standard state conditions. When two half-reactions are connected to the one with the larger (more positive) E° red goes as a reduction On the other one (less positive E° red) goes as an oxidation © 2015 Pearson Education, Ltd.  Write the half-cell reaction and the overall cell reaction for the electrochemical cell: Zn | Zn2+ (1.0M) || Pb2+ (1.0M) | Pb Calculate the standard e.m.f for the cell if standard electrode potential (reduction) for Pb2+ | Pb and Zn2+ | Zn electrodes are – 0.126 V and – 0.763 V respectively Half reactions are- Oxidation half-reaction: Zn → Zn2+ + 2e– Reduction half-reaction: Pb2+ + 2e– → Pb Overall cell reaction: Zn + Pb2+ → Zn2+ + Pb E°cell = E°R – E°L E°cell = –0.126 – ( –0.763) E°cell = 0.637 V. © 2015 Pearson Education, Ltd.  Calculate the standard reduction potential of Ni2+|Ni electrode when the cell potential for the cell Ni | Ni2+ (1M)|| Cu2+ (1M) | Cu is 0.59 V (E° Cu2+| Cu = 0.34 V). Ni | Ni2+ (1M)|| Cu2+ (1M) | Cu The e.m.f of the cell E°cell = E°R – E°L E°cell = E° (Cu2+ | Cu) – E° (Ni | Ni2+) 0.59 V = 0.34 – E° (Ni | Ni2+) E° (Ni | Ni2+) = 0.34 – 0.59 = – 0.25 V © 2015 Pearson Education, Ltd.  Electrolysis Electrolysis is the process of using electrical energy to break a compound apart. Electrolysis is done in an electrolytic cell. Electrolytic cells can be used to separate elements from their compounds. © 2015 Pearson Education, Ltd. © 2015 Pearson Education, Ltd.  Electrolysis In electrolysis we use electrical energy to overcome the energy barrier of a nonspontaneous reaction, allowing it to occur. The reaction that takes place is the opposite of the spontaneous process. 2 H2(g) + O2(g) → 2 H2O(l) spontaneous 2 H2O(l) → 2 H2(g) + O2(g) electrolysis Some applications of electrolysis are the following: (1) Metal extraction from minerals and purification (2) Production of H2 for fuel cells (3) Metal plating © 2015 Pearson Education, Ltd.  Electrolytic Cells The source of energy is a battery or DC power supply. – The positive terminal of the source is attached to the anode. – The negative terminal of the source is attached to the cathode. Electrolyte can be either an aqueous salt solution or a molten ionic salt. Cations in the electrolyte are attracted to the cathode and anions are attracted to the anode. Cations pick up electrons from the cathode and are reduced; anions release electrons to the anode and are oxidized. © 2015 Pearson Education, Ltd.  Electrolysis and Electrolytic Cells Electrolysis of Aqueous Sodium Chloride © 2015 Pearson Education, Ltd.  Electrolysis and Electrolytic Cells Electrolysis of Molten Sodium Chloride © 2015 Pearson Education, Ltd.  Commercial Applications of Electrolysis Downs Cell for the Production of Sodium Metal © 2015 Pearson Education, Ltd.  Electroplating In electroplating, the work piece is the cathode. Cations are reduced at cathode and plate to the surface of the work piece. The anode is made of the plate metal. The anode oxidizes and replaces the metal cations in the solution. © 2015 Pearson Education, Ltd.  Commercial Applications of Electrolysis Electrorefining of Copper Metal © 2015 Pearson Education, Ltd.  Rusting At the anodic regions, Fe(s) is oxidized to Fe2+. The electrons travel through the metal to a cathodic region where O2 is reduced. – In acidic solution from gases dissolved in the moisture The Fe2+ ions migrate through the moisture to the cathodic region, where they are further oxidized to Fe3+, which combines with the oxygen and water to form rust. – Rust is hydrated iron(III) oxide, Fe2O3 · nH2O. The exact composition depends on the conditions. – Moisture must be present. Water is a reactant. It is required for ion flow between cathodic and anodic regions. Electrolytes promote rusting. – They enhance current flow. Acids promote rusting. – Lowering pH © 2015 Pearson Education, Ltd.  Corrosion Corrosion: The oxidative deterioration of a metal © 2015 Pearson Education, Ltd.  Corrosion Prevention of Corrosion For some metals, oxidation protects the metal (aluminum, chromium, magnesium, titanium, zinc, and others). For other metals, there are two main techniques. © 2015 Pearson Education, Ltd.  Corrosion Prevention of Corrosion 1. Galvanization: The coating of iron with zinc © 2015 Pearson Education, Ltd.  Corrosion Prevention of Corrosion 2. Cathodic Protection: Instead of coating the entire surface of the first metal with a second metal, the second metal is placed in electrical contact with the first metal: Attaching a magnesium stake to iron will corrode the magnesium instead of the iron. Magnesium acts as a sacrificial anode. © 2015 Pearson Education, Ltd.  Sacrificial Anode If a metal more active than iron, such as magnesium or aluminum, is in electrical contact with iron, the metal rather than the iron will be oxidized. This principle underlies the use of sacrificial electrodes to prevent the corrosion of iron. © 2015 Pearson Education, Ltd.  Ways to prevent corrosion passivation (formation of a thin oxide layer by treatment with an oxidizing agent) formation of an alloy (stainless steel) coating with a layer of a less active metal (tin cans) cathodic protection (use zinc or magnesium as a sacrificial metal) © 2015 Pearson Education, Ltd.