Redox Review & Electrochemistry Reduction-Oxidation Reactions

Redox Review & Electrochemistry Reduction-Oxidation Reactions • All electron chemical reactions involving the transfer of electrons are classified as Reduction-Oxidation (Redox) reactions. • Oxidation and reduction occur simultaneously Reference:Redox.html09-UN161aRedoxReactNaCl.jpg Reduction-Oxidation Terminology • Reduction: substance gains electrons from another species. • Oxidation number/state of atom gaining electrons is reduced (e.g., goes down). • In an electrochemical cell the reaction happens at the cathode. • In a ½ cell reaction, electrons are “reactants.” • Oxidation: substance loses electrons from another species. • Oxidation number/state of atom losing electrons increases. • In an electrochemical cell the reaction happens at the anode. • In a ½ cell reaction, electrons are “products.” • Oxidizing agent is an electron acceptor; it causes another species to be OXIDIZED, but itself (agent) is being reduced. • Reducing agent is an electron donor; it cause another species to be REDUCED but it (agent) is being oxidized. Review Problem: Determine the oxidation state for Cr in Cr2O72- Balancing Redox Reactions: Half-Reaction Method • A redox reaction MUST not only be mass balanced, but also charged balanced. • The chemical reaction is split into to half reactions to determine which species is being oxidized and which species is being reduced. • Charge balancing determines the number of electrons transferred between species for each half reaction. • Process: • The redox reaction is split into two separate half-reactions in order to determine which species is being oxidized and which species is being reduced. • The oxidation half-reaction has electrons as products. • The reduction half-reaction has electrons as reactants. • Each half-reaction needs to be mass & charge balanced. Balancing Redox Reactions: Half-Reaction Method • Environment concerns: The experimental environment (e.g. neutral, acidic or basic) that the redox reaction is occurring must be taken into consideration when balancing redox reactions. • Acidic or Neutral • Balance Oxygen atoms by adding one water molecule for every O atom needed. • Balance Hydrogen atoms by adding H+ ions. • Basic • Balance Oxygen atoms by adding one OH- ion for every O atom needed. • Balance Hydrogen atoms by adding H+ ions to counter balance OH- ions to make H2O. • A mass and charged balance redox equation is obtained by adding the two half-reactions together. • Electrons must not appear in the final balanced chemical equation. Problems: Balancing Redox Reactions 1. MnO4− (aq) + Cr3+ (aq) ⟶ Mn2+ (aq) +Cr2O72− (aq) Acidic 2. H2 (g) + Cu2+ (aq) ⟶ Cu (s) Neutral 3. MnO4− (aq) + Cr(OH)3 (s) ⟶ MnO2(s)+CrO42−(aq) Basic What is Electrochemistry? • Electrochemistry is about the chemical reactions that produce electricity and/or that are associated with the passage of an electrical current. • Electron transfer reactions are chemical reactions involved in the transfer of electrons between atoms. • Example: • Single displacement type and combustion reactions; • Some synthesis and decomposition reactions. Electrochemistry Terminology Quantity Term Definition Electrical charge Charge on a proton Electrical current The movement of electrons (i.e., charge). • An electrode surface area dictates the number of electrons that can flow • Larger batteries produce larger currents. Electrical potential (E) Energy (i.e. force) needed to move charge (i.e. electrons) Electrical field The potential acting upon other charges in a defined area. • The voltage needed to drive electrons through the external circuit. • Amount of force pushing the electrons through the wire is called the electromotive force, emf. Unit 1.602 x 10-19 Coulomb (C) Ampere (A) = 1 C/sec Volt (V) = J/C Electrochemical Cell: Redox in Action • Electrons are “driven” from anode to cathode by an electromotive force (emf). • Oxidation occurs at the anode (-). • Reduction occurs at the cathode (+). • Electrochemical Cell types • Voltaic (galvanic or Daniel) cells have positive cell potential. • Spontaneous: DGo (-); K>1 & oEcell (+) • Chemical change produces a current (i.e. Energy is released). • Everyday batteries are voltaic cells. • Electrolytic cells have a negative cell potential • Non spontaneous: DGo (+); K<1 & oEcell (-) • Uses electrical energy to overcome the energy barrier and causes a non spontaneous reaction to occur. Reference: chemistry.libretexts Standard Electrochemical Potential (E°cell) • E is an intensive physical property of matter; therefore, E values are NOT modified (i.e. even if you needed to multiply the half-reactions to cancel out the number of electrons to balance the redox reaction the E value is NOT changed). • The difference in potential energy between the anode and the cathode in an electrochemical cell is the potential, E. • Standard potential of an electrochemical cell is represented by the symbol, E°. • Standard conditions: T is 25 °C; pressure is 1 atm for gases and solutions have a 1 M concentration. • E°cell is computed from the cell’s half-reaction obtained from a Redox potential table (E°). • All values are listed as REDUCTION potentials. • How to compute an electrochemical standard potential: • Method 1: The value sign is NOT switched for the Eocell oxidation reaction. Eocell = Eocathode - Eoanode Eocell = 0.34V – (-0.76)V = 1.10V • Method 2 : The value sign IS switched for the Eocell oxidation reaction. Eocell = Eocathode + Eocell = 0.34V Eoanode + (0.76)V = 1.10V Electrochemical Cell Notation: Shorthand Representation • Shorthand • The oxidation half-cell is on the left side. • The reduction half-cell is on the right side. electrode | electrolyte || electrolyte | electrode • Single vertical line, | indicates a phase barrier. • If multiple electrolytes are in same phase, a comma is used rather than a vertical line |. • Often use an inert electrode. • Double vertical lines, || indicates the presence of a salt bridge. Reference: socratic.org Reduction Half-Reaction (Eocell), V F2(g) + 2 e−→ 2 F−(aq) 2.87 Ag+(aq) + 1 e−→ Ag(s) 0.80 I2(s) + 2 e−→ 2 I−(aq) 0.54 Mg2+(aq) + 2 e−→ Mg(s) -2.37 Problems: A. Determine the standard electrochemical potential, Ecello . B. Determine the reducing agent & oxidizing agent. C. Determine the spontaneity of the reaction. Justify your answer. D. Write the shorthand notation for the electrochemical cell. Mg2+(aq) + 2 Ag(s) → Mg(s) + 2 Ag+(aq) Problem: Given the electrochemical cell: Cr|Cr3+ (aq)||Cu2+(aq)|Cu 1. Write the half-reactions and overall redox reaction for this cell. 2. Determine the Ecello of the reaction. 3. Sketch an electrochemical cell of this reaction and label all appropriate parts (e.g., cathode & reduction ½ reaction, anode & oxidation ½ reaction, salt bridge, flow of electron). 4. Is this an example of a voltaic or electrolytic cell? 5. Predict the reactions DGo value. Reduction Half - Reaction Eredo (V) Cu2+(aq) + 2 e−→ Cu(s) + 0.34 Cr(aq) + 3 e−→ Cr3+(aq) - 0.74 Thermochemical (DG°), Electrochemical (Eocell) and Equilibrium (K) Relationships Eocell DG° K Spontaneous Reaction? Product favored <1 (negative) >1 (positive) >1 yes Reactant favored <1 (positive) <1 (negative) <1 no =0 =1 Equilibrium = 0 Problem: Determine the equilibrium constant value and the DGo value in kJ for the electrochemical cell: 2 Ag+(aq) + Fe(s) 2 Ag(s) + Fe2+(aq) at 25 oC. Reduction Half - Reaction Eredo (V) Fe2+(aq) + 2 e−→ Fe(s) - 0.45 Ag+ (aq) + 1 e−→ Ag (s) + 0.80 DGo = -nF Ecello Relating Thermodynamics to Cell Potential • The relationships between the terms of ΔGo, K, and Eocell is given in the diagram below. • Given any one of the three quantities, the other two can be calculated and used to determine whether a process is spontaneous or not. Reference: opentstax Resulting in the Nernst Equation: E cell = Ecello – (RT/nF) lnQ Problem: Determine the Ecell value & spontaneity of the following electrochemical cell at room temperature (25 oC). Co (s) + Fe2+ (aq) → Co2+ (aq) + Fe (s) [Fe2+] = 1.94 M [Co2+] = 0.15 M E cell = Ecello – (RT/nF) lnQ Reduction Half - Reaction Eredo (V) Fe2+ (aq) + 2 e−→ Fe(s) - 0.45 Co2+ (aq) + 2e−→ Co (s) - 0.28 Electrolysis Electrolysis of water • Electrolytic cells are nonspontaneous reactions (i.e., Ecello is negative) where electrical energy (stored) is used to drive the reaction. • The use of electric energy (stored) to drive a nonspontaneous reaction forward the process referred to as electrolysis. • Examples: • Electrolysis of water (H2O) into its two elements – H2(g) & O2(g). • Refining of metals • Plating Reference: Chemistry.Lbre.texts.org Plating of metals Reference: sparknotes.com The Quantitative Side of Electrolysis • The amount of current (i.e., with units of amperes (A = 1 C/sec) ) that flows through an electrolytic cell is related the moles of electrons passing through the current. • Stoichiometry of the reaction indicates the moles of electrons flowing between reactants and products. • Faraday’s Law: • The amount of metal deposited during an electrolysis process is directly proportional to the charge (i.e., coulombs, (Q ) of the metal’s cation, the current & the length of time of the process. • Charge (Q) = current (I) x time • Example: Consider electrolysis of aqueous Copper (II) ion: Cu2+ (aq) + 2e- → Cu(s) It will take 2 mole electrons to produce 1 mole copper metal. • A Faraday (i.e., 96,485 C/mole electrons) relate moles of electrons to current Problem: A 0.010-mm layer of chromium must be deposited on an automobile part having a total surface area of 3.30 m2 from a solution of containing chromium(III) ions. • How many minutes would it take to deposit this layer of chromium, if the current was 33.46 A? • The density of chromium (metal) is 7.19 g/cm3. • Faraday = 96,485 C/mol e• 1 mole Cr = 52.00 grams • Q = I x time (sec) An Overview on Batteries • Batteries are electrochemical cells that produce electrical energy. • Two categories of batteries: • Primary: • Single cell use or non rechargeable batteries • Galvanic/voltaic • Example: alkaline batteries • Secondary: • Multiple use or rechargeable batteries • Galvanic/voltaic • Examples: Lead-acid (car); Ni-Cd; Li; Ni-MH (nickel metal hydride) Illustrations of Batteries: Alkaline Alkaline batteries (1950) were designed as direct replacements for zinc-carbon (dry cell) batteries. anode: Zn(s)+2OH−(aq) ⟶ ZnO(s) +H2O(l) +2e− cathode: 2MnO2(s) + H2O(l) + 2e− ⟶ Mn2O3(s) + 2OH−(aq) overall: Ecell° = Reference: Openstax Zn(s) + 2MnO2(s) ⟶ ZnO(s) + Mn2O3(s) +1.43 V Eredo = − 1.28 V = + 0.15 V Illustrations of Batteries: Car Battery – Lead-Acid • A lead-acid battery has six cells in series with sulfuric acid as the electrolyte (30% H2SO4). • Anode is Pb. • Pb(s) + SO42−(aq) → PbSO4(s) + 2 e− • Cathode is coated with PbO2. • PbO2(s) + 4 H+(aq) + SO42−(aq) + 2 e− → PbSO4(s) + 2 H2O(l) • Each battery cell has a voltage of 2.09 V. Reference: Openstax Illustrations of Batteries: NiCd • Construction: • A nickel-plated cathode and a cadmium plated anode with a potassium hydroxide electrolyte. • Jelly-roll design: • Material is place between the electrode (plates) to prevent shorting out when they (anode and cathode plates) are coiled prior to encasement. • “Rolled” design increases current out put when compared to similarsized traditional alkaline batteries. • Voltage 1-2 to 1.25 per cell • Rechargable for approximately 1000 times. Reference: Openstax • Anode: Cd(s) +2 OH−(aq) ⟶ Cd(OH)2(s) + 2e− • Cathode: NiO2(s) + 2H2O(l) + 2e− ⟶ Ni(OH)2(s) + 2OH−(aq) • Overall: Cd(s) + NiO2(s) + 2H2O(l) ⟶ Cd(OH)2(s) + Ni(OH)2(s) Illustrations of Batteries: Ni-Metal Hydride (NiMH) • Similar “jelly-roll” design as NiCd batteries. • Anode is a metal hydride, and it is oxidized. • M∙H(s) + OH−(aq) → M(s) + H2O(l) + e− • The cathode is Ni coated with NiO(OH); where NiO(OH) is reduced. • NiO(OH)(s) + H2O(l) + e− → Ni(OH)2(s) + OH− • The battery has as a cell voltage of ~ 1.30 V. • NiCad (Ni & Cd) batteries were standard battery in this category, but in recent years have been phased out due to disposal problems (e.g., cadmium toxicity) Reference: en.globtek.com/nimh-battery-safety-notes Illustrations of Batteries: Lithium • The reactions: • Anode: LiCoO2 ⇌ Lix−1CoO2 + xLi+ + xe− • Cathode: x Li+ + xe−+ x C6 ⇌ xLiC6 • Overall: LiCoO2 + xC6 ⇌ Lix−1CoO2 + xLiC6 • * Coefficients (x) representing moles, are no more than about 0.5 moles. • KOH electrolyte The charge flows between the electrodes as the lithium ions move between the anode and cathode. Reference:openstax • Popular • Battery voltage is about 3.7 V per cell, which is a large amount current. • Light weight comparable to other batteries • Produce a near constant voltage as discharge • In storage has looses voltage slowly. Electrolytic Cell: Fuel Cell • A fuel cell is a device that converts chemical energy into electrical energy. • Fuel cells are electrolytic cells; nonspontaneous reactions (Ecello is negative). • Require an outside, continuous fuel reservoir in order to continue to produce electrical energy. • Example: Hydrogen fuel cells. • Reaction: 2H2(g) + O2(g) ⟶ 2H2O(l) (exhaust) • Anode: 2H2(g) + 2O2−(aq) ⟶ 2H2O + 4e− • Cathode: O2(g) + 4e− ⟶ 2O2− (aq) • Oxygen from air reacts with hydrogen gas to produce water and electricity. • Voltage is about 0.9V. • Efficiency: 40% to 60% compared to typical combustion engine which is ~25% to 35%. Reference:openstax Corrosion: Natural Redox Reaction • Corrosion is the spontaneous degradation (i.e. oxidation) of metal in the environment. • Example: Rusting of metal is a corrosion reaction. • Rust is hydrated iron (III) oxide. • To have rusting need: • Moisture (e.g., water) must be present. • Required for flow between cathode and anode. • Electrolytes promote rusting. • Enhance current flow. • Acids promote rusting. • Lower pH = lower E°red Reference: openstax Strategies to Prevent Corrosion (Rusting) • Coating (e.g., plating and/or painting) surface to reduce contact with environment. • Plating: Sacrifice the plating layer. • Metals, like Al, form an oxide that strongly attaches to the metal surface, preventing the rest from corroding. • Cathodic protection. • Uses an active metal like zinc or magnesium for the anode where the cathode is the object (i.e., metal) that need to be protected from corroding (i.e., oxidizing) . Reference: openstax