PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 18 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state university

CHAPTER 18 ELECTROCHEMISTRY

- Is the study of the relations between chemical reactions and electricity - Electrochemical processes involve the transfer of electrons from one substance to another ELECTROCHEMISTRY

- Also called redox reactions - Involve transfer of electrons from one species to another Oxidation - loss of electrons Reduction - gain of electrons - Ionic solid sodium chloride (Na + and Cl - ions) formed from solid sodium and chlorine gas 2Na(s) + Cl 2 (g) → 2NaCl(s) - The oxidation (rusting) of iron by reaction with moist air 4Fe(s) + 3O 2 (g) → 2Fe 2 O 3 (s) OXIDATION-REDUCTION REACTIONS

- There is no transfer of electrons from one reactant to another reactant Examples BaCO 3 (s) → BaO(s) + CO 2 (g) Double-replacement reactions NONREDOX REACTIONS

OXIDATION NUMBER (STATE) The concept of oxidation number - provides a way to keep track of electrons in redox reactions - not necessarily ionic charges Conventionally - actual charges on ions are written as n+ or n- - oxidation numbers are written as +n or -n Oxidation - increase in oxidation number (loss of electrons) Reduction - decrease in oxidation number (gain of electrons)

OXIDATION NUMBERS 1. Oxidation number of uncombined elements = 0 Na(s), O 2 (g), H 2 (g), Hg(l) 2. Oxidation number of a monatomic ion = charge Na + = +1, Cl - = -1, Ca 2+ = +2, Al 3+ = Oxygen is usually assigned -2 H 2 O, CO 2, SO 2, SO 3 Exceptions: H 2 O 2 (oxygen = -1) and OF 2 (oxygen = +2)

4. Hydrogen is usually assigned +1 Exceptions: -1 when bonded to metals (+1in HCl, NH 3, H 2 O and -1in CaH 2, NaH) 5. Halogens are usually assigned -1 (F, Cl, Br, I) Exceptions: when Cl, Br, andI are bonded to oxygen or a more electronegative halogen (Cl 2 O: O = -2 and Cl = +1) 6. The sum of oxidation numbers for - neutral compound = 0 - polyatomic ion = charge (H 2 O = 0, CO 3 2- = -2, NH 4 + = +1) OXIDATION NUMBERS

CO 2 The oxidation state of oxygen is -2 CO 2 has no charge The sum of oxidation states of carbon and oxygen = 0 1 carbon atom and 2 oxygen atoms 1(x) + 2(-2) = 0 x = +4 CO 2 x-2 for each oxygen OXIDATION NUMBERS

CH4CH4 x+1 1(x) + 4(+1) = 0 x = -4 OXIDATION NUMBERS

NO3-NO3- x-2 1(x) + 3(-2) = -1 x = +5 OXIDATION NUMBERS

- Just the oxidation or the reduction is given - The transferred electrons are shown Oxidation Half-Reaction - Electrons are on the product side of the equation Reduction Half-Reaction - Electrons are on the reactant side of the equation HALF-REACTIONS

For the redox reaction Cu 2+ (aq) + Zn(s) → Cu(s) + Zn 2+ (aq) Zn is oxidized (oxidation number changes from 0 to +2) Cu is reduced (oxidation number changes from +2 to 0) The oxidation half-reaction is: Zn(s) → Zn 2+ (aq) + 2e - The reduction half-reaction is: Cu 2+ (aq) + 2e - → Cu(s) HALF-REACTIONS

Oxidizing Agent - Is the reduced species (accepts electrons from another species) Reducing Agents - Is the oxidized species (donates electrons to another species) For the redox reaction Cu 2+ (aq) + Zn(s) → Cu(s) + Zn 2+ (aq) Cu is reduced so is the oxidizing agent Zn is oxidized so is the reducing agent OXIDIZING AND REDUCING AGENTS

Half-Reaction Method Acidic Solutions BALANCING REDOX EQUATIONS

- Write separate oxidation and reduction half-reactions - Balance all the elements except hydrogen and oxygen in each - Balance oxygen using H 2 O(l), hydrogen using H + (aq), and charge using electrons (e - ) - Multiply both half-reactions by suitable factors to equalize electron count - Combine the balanced half-reactions BALANCING REDOX EQUATIONS

Balance the following redox reaction in acid medium MnO 4 - (aq) + Fe 2+ (aq) → Fe 3+ (aq) + Mn 2+ (aq) Answer MnO 4 - (aq) + 5Fe 2+ (aq) + 8H+(aq) → 5Fe 3+ (aq) + Mn 2+ (aq) + 4H 2 O(l) BALANCING REDOX EQUATIONS

Half-Reaction Method Basic Solutions BALANCING REDOX EQUATIONS

- Balance equation as if it were acidic - Note H + ions and add same number of OH - ions to both sides - Cancel H + and OH - (=H 2 O) with H 2 O on other side BALANCING REDOX EQUATIONS

Balance the following redox reaction in basic medium MnO 4 - (aq) + C 2 O 4 2- (aq) → MnO 2 (s) + CO 3 2- (aq) Answer 2MnO 4 - (aq) + 3C 2 O 4 2- (aq) + 4OH - (aq) → 2MnO 2 (s) + 6CO 3 2- (aq) + 2H 2 O(l) BALANCING REDOX EQUATIONS

ELECTRODE - Conducts electrons into or out of a redox reaction system Examples platinum wire carbon (glassy or graphite) indium tin oxide (ITO) Electroactive Species - Donate or accept electrons at an electrode

ELECTRODE Chemically Inert Electrodes - Do not participate in the reaction Examples Carbon, Gold, Platinum, ITO Reactive Electrodes - Participate in the reaction Examples Silver, Copper, Iron, Zinc

CHEMICAL CHANGE Spontaneous Process - Takes place with no apparent cause Nonspontaneous Process - Requires something to be applied in order for it to occur (usually in the form of energy)

VOLTAIC (GALVANIC) CELL - Spontaneous reaction - Produces electrical energy from chemical energy - Can be reversed electrolytically for reversible cells Example Rechargeable batteries Conditions for Non-Reversibility - If one or more of the species decomposes - If a gas is produced and escapes

- A spontaneous redox reaction generates electricity - One reagent is oxidized and the other is reduced - The two reagents must be separated (cannot be in contact) - Each is called a half-cell - Electrons flow through a wire (external circuit) VOLTAIC (GALVANIC) CELL

Oxidation Half reaction - Loss of electrons - Occurs at anode (negative electrode) - The left half-cell by convention Reduction Half Reaction - Gain of electrons - Occurs at cathode (positive electrode) - The right half-cell by convention VOLTAIC (GALVANIC) CELL

Salt Bridge - Connects the two half-cells (anode and cathode) - Filled with gel containing saturated aqueous salt solution (KCl) - Ions migrate through to maintain electroneutrality - Prevents charge buildup that may cease the reaction process VOLTAIC (GALVANIC) CELL

For the overall reaction Cu 2+ (aq) + Zn(s) → Cu(s) + Zn 2+ (aq) VOLTAIC (GALVANIC) CELL Anode Oxidation Zn(s) → Zn 2+ (aq) + 2e - Cathode Reduction Cu 2+ (aq) + 2e - → Cu(s) Salt bridge (KCl) Cl - K+K+ Voltmeter - + e-e- e-e- Cu electrode Zn electrode Cu 2+ Zn 2+

Voltage or Potential Difference (E) - Referred to as the electromotive force (emf) - Is the voltage measured - Measured by a voltmeter (potentiometer) connected to electrodes Greater Voltage - More favorable net cell reaction - More work done by flowing electrons (larger emf) POTENTIALS VOLTAIC CELL

Voltage or Potential Difference (E) - Work done by or on electrons when they move from one point to another Units: volts (V or J/C) Work (J) = E (V) x q (C) POTENTIALS OF VOLTAIC CELL

Charge Charge (q) of an electron = x C Charge (q) of a proton = x C C = coulombs Charge of one mole of electrons = (1.602 x C)(6.022 x /mol) = x 10 4 C/mol = Faraday constant (F) q = n x F(n = number of moles) POTENTIALS OF VOLTAIC CELL

Current - The quantity of charge flowing past a point in an electric circuit per second Units Ampere (A) = coulomb per second (C/s) Charge (C) = current (A) x time (s) POTENTIALS OF VOLTAIC CELL

STANDARD POTENTIALS Electrode Potentials - A measure of how willing a species is to gain or lose electrons Positive Voltage (spontaneous process) - Electrons flow into the negative terminal of voltmeter (flow from negative electrode to positive electrode) Negative Voltage (nonspontaneous process) - Electrons flow into the positive terminal of voltmeter (flow from positive electrode to negative electrode) Conventionally - Negative terminal is on the left of galvanic cells

Standard Reduction Potential (E o ) - Used to predict the voltage when different cells are connected - Potential of a cell as cathode compared to standard hydrogen electrode - Species are solids or liquids - Activities = 1 - We will use concentrations for simplicity Concentrations = 1 M Pressures = 1 atm STANDARD POTENTIALS

Standard Hydrogen Electrode (SHE) - Used to measure E o for half-reactions - Connected to negative terminal - Pt electrode - Acidic solution in which [H + ] = 1 M - H 2 gas (1 bar) is bubbled past the electrode H + (aq, 1 M) + e- ↔ 1/2H 2 (g, 1 atm) Conventionally, E o = 0 for SHE STANDARD POTENTIALS

The E o for Ag + (aq) + e - ↔ Ag(s) is V Implies - if a sample of silver metal is placed in a 1 M Ag + solution, a value of V will be measured with S. H. E. as reference STANDARD POTENTIALS

Silver does not react spontaneously with hydrogen 2H + (aq) + 2e - → H 2 (g)E o = V Ag + (aq) + e - → Ag(s)E o = V Reverse the second equation (sign changes) Ag(s) → Ag + (aq) + e - E o = V Multiply the second equation by 2 (E o is intensive so remains) 2Ag(s) → 2Ag + (aq) + 2e - E o = V Combine (electrons cancel) 2Ag(s) + 2H + (aq) → 2Ag + (aq) + H 2 (g) E o = V STANDARD POTENTIALS

Consider Cu 2+ (aq) + Zn(s) → Cu(s) + Zn 2+ (aq) Cu 2+ (aq) + 2e - → Cu(s)E o = V Zn 2+ (aq) + 2e - → Zn(s) E o = V Reverse the second equation (sign changes) Zn(s) → Zn 2+ (aq) + 2e - E o = V Combine (electrons cancel) Cu 2+ (aq) + Zn(s) → Cu(s) + Zn 2+ (aq) E o = V E o is positive so reaction is spontaneous Reverse reaction is nonspontaneous STANDARD POTENTIALS

- Half-reaction is more favorable for more positive E o - Refer to series and E o values in textbook - For combining two half reactions, the one higher in the series proceeds spontaneously as reduction under standard conditions - The one lower in the series proceeds spontaneously as oxidation under standard conditions Formal Potential - The potential for a cell containing a specified concentration of reagent other than 1 M STANDARD POTENTIALS

- When a half reaction is multiplied by a factor E o remains the same - For a complete reaction E cell = E + - E - and E o = E + o - E - o E + = potential at positive terminal E - = potential at negative terminal STANDARD POTENTIALS

For the Cu – Fe cell at standard conditions Cu e - ↔ Cu(s) V Fe e - ↔ Fe(s) V E cell = V Galvanic (overall) Reaction Cu 2+ (aq) + Fe(s) ↔ Cu(s) + Fe 2+ (aq) STANDARD POTENTIALS

- Positive E implies spontaneous forward cell reaction - Negative E implies spontaneous reverse cell reaction If cell runs for long - Reactants are consumed - Products are formed - Equilibrium is reached - E becomes 0 - Reason why batteries run down STANDARD POTENTIALS

∆G AND K eq - Recall that spontaneous reaction has a negative value of ∆G ∆G = -nFE n = number of moles of electrons transferred F = Faraday constant = x 10 4 C/mol E = cell potential (V or J/C) Under Standard Conditions ∆G o = -nFE o

NERNST EQUATION For the half reaction aA + ne - ↔ bB The half-cell potential (at 25 o C), E, is given by

NERNST EQUATION E o = standard electrode potential R = gas constant = J/K-mol T = absolute temperature F = Faraday’s constant = x 10 4 C/mol n = number of moles of electrons transferred

NERNST EQUATION - The standard reduction potential (E o ) is when [A] = [B] = 1M - [B] b /[A] a = Q = reaction quotient - Concentration for gases are expressed as pressures in atm - Q = 1 for [ ] = 1 M and P = 1 atm logQ = 0 and E = E o - Pure solids, liquids, and solvents do not appear in Q expression

NERNST EQUATION At cell equilibrium at 25 o C E = 0 and Q = K eq (the equilibrium constant) Or Positive E o implies K eq > 1 Negative E o implies K eq < 1

REFERENCE ELECTRODES - Provide known and constant potential Examples Silver-silver chloride electrode (Ag/AgCl) Saturated Calomel electrode (SCE)

INDICATOR ELECTRODES - Respond directly to the analyte - Two Classes of Indicator Electrodes Metal Electrodes - Surfaces on which redox reactions take place Examples Platinum Silver

INDICATOR ELECTRODES - Respond directly to the analyte - Two Classes of Indicator Electrodes - Ion-Selective Electrodes - Selectively binds one ion (no redox chemistry) Examples pH (H + ) electrode Calcium (Ca 2+ ) electrode Chloride (Cl - ) electrode

ELECTROLYSIS - Voltage is applied to drive a redox reaction that would not otherwise occur Examples - Production of aluminum metal from Al 3+ - Production of Cl 2 from Cl - - Electroplating

ELECTROLYSIS CELL - Consists of two electrodes in an electrolyte solution - Nonspontaneous reaction - Requires electrical energy to occur

CORROSION - Oxidation of a metal to produce compounds of the metal Examples - Rusting of iron (Fe forms Fe 2 O 3 ·xH 2 O - Copper in bronze forming copper(II) compounds (green color) Prevention - Painting or coating - Plating of iron with chromium - Anodic protection (metal is oxidized under controlled conditions) - Cathodic protection (a more reactive metal is placed in contact)