Electrochemistry

Electrochemistry Chapter 20

IB Topics 9.4-9.5 • Text Pages 851-854, 876-881 • Oxford p. 56-57

Electrochemistry • Study of interchange of chemical and electrical energy. • Two main processes • Generation of electrical current from spontaneous chemical reactions • Use of a current to produce chemical change

Redox Reaction • 8H1+(aq) + MnO41-(aq) +5Fe2+ (aq) Mn2+ (aq) 5Fe3+ (aq) + 4H2O (l) Reduced OxidizedOxidizing ReducingAgentAgent • When the oxidizing and reducing agents are present in the same solution, no useful work is obtained. • Energy is released as heat

Half Cell • A metal in contact with a solution of its own ions

Separating the Oxidizing and Reducing Agents of a Redox ReactionAllows the energy to be “captured” in the form of electricity. Gaining Electrons Losing Electrons Negative charge builds up Positive charge builds up Current flows for an instant then ceases. Charge builds up in the two compartments

Solutions must be connected so that ions can flow to keep net charge in each compartment at zeroThe salt bridge keeps the solutions electrically neutral by allowing cations and ions to flow freely between the cells.

Galvanic (Voltaic ) Cell • Device in which chemical energy is converted to electrical energy. • Anode—compartment in which oxidation occurs. • Cathode—compartment in which reduction occurs. • Electrons flow from anode to cathode

Cell Potential/Electromotive force • The reaction in a voltaic cell is always redox. • Oxidizing agent “pulls” electrons through a wire from reducing agent. • Ecell • Emf • Units are volts (V) • Measured by voltmeter

Galvanic (Voltaic) Cells

Practice • 2H1+ (aq) + Zn (s)  Zn2+ (aq)+ H2(g) Al3+ (aq) + Mg (s)  Al (s) + Mg2+ (aq) Show the following for each: -Half reaction at each electrode -Direction of electron flow

Standard Electrode Potentials • Topic 19.1.1-19.1.4 • Oxford p. 58, 59 • Text p. 855-870

Cell Potential/Electromotive force • The reaction in a voltaic cell is always redox. • Oxidizing agent “pulls” electrons through a wire from reducing agent. • Ecell • Emf • Units are volts (V) • Measured by voltmeter

Standard Hydrogen Electrode • The hydrogen half cell is used as a standard • Standard conditions of 1 atm, 298 K, and 1.0 M • The Standard electrode potential of the hydrogen half cell is zero volts • E cell = 0 volts Pt electrode is inert. It is in contact with H+(aq). This is called a standard hydrogen electrode.

The Electrochemical Series • Elements higher in the chart are more reactive than H and have a negative reduction potential • Elements lower in the chart are more reactive than H and have a positive reduction potential • Electrons flow from more negative to less negative 15

Standard Reduction Potentials • Total potential of the cell is 0.76 V. • Standard conditions [H+] = 1 M and PH2 = 1atm • 2H1+ (aq) + Zn (s)  H2(g) + Zn2+ (aq) • Cathode reaction--Reduction: 2H1+ (aq) + 2e-  H2(g) E°cell = 0 V • Anode Reaction--Oxidation: Zn (s)  Zn2+ (aq) + 2e- E°cell = 0.76 V

Voltaic Cell Potential • E°cell = E° red cathode- E° red anode • E°cell = E° red cathode+ E° ox anode

Sample Problem 1 • Consider a galvanic (voltaic) cell based on the following reaction • Draw the cell • Write the half reactions • Assign SRP’s • Give the balanced cell reaction • Calculate E°cell Al3+ (aq) + Mg (s)  Al (s) + Mg2+ (aq)

Sample Problem 2 • Give the balanced reaction • Calculate E cell

Ecelland Gibb’s Free Energy (Δ G) • Voltaic cells with positive Ecell values provide energy to do work  have - ΔG values

Line Notation for Electrochemical Cells • Shorthand showing two half-cells connected by a salt bridge or porous barrier, : Zn(s)/Zn2+(aq)//Cu2+(aq)/Cu(s) anodecathode The electrodes are shown on the ends and the electrolytes for each side are shown in the middle. 21

Electrolysis HL • Topic 19.2.1-19.2.3 • Text p. 876-881 • Oxford p. 57-58

Electrolysis • Forcing electricity through a cell to produce a chemical change for which the cell potential is negative. • Causes a non-spontaneous reaction to occur. • Charging batteries, production of aluminum, Chrome plating, obtain reactive metals (Na) from ores

Power source • Anode/Cathode reversed • Electron flow reversed • Ion flow reversed

Electrolytic Cell • 1. Electrons are "produced" in the battery at the anode, the site of oxidation. • 2. The electrons leave the electrochemical cell through the external circuit. • 3. These negative electrons create a negative electrode in the electrolytic cell which attracts the positive Na+ ions in the electrolyte. Na+ ions combine with the free electrons and become reduced (2Na+ + 2e- → Na ) • 4. Meanwhile the negative Cl- become attracted to the positive electrode of the electrolytic cell. At this electrode chlorine is oxidized, releasing electrons (Cl-→ Cl2 + 2 e-) • 5. These electrons travel through the external circuit, returning to the electrochemical cell.

Factors Affecting the Discharge of Ions During Electrolysis • Cation discharged at cathode (negative electrode) • Anion discharged at anode (positive electrode). • In (aq) solutions, H+ and OH- will be present also. • Which ion discharged influenced by • Its position in electrochemical series • Concentration • Nature of electrode

Role of Water in Electrolysis • Aqueous solutions • Water can be oxidized or reduced • H2O (l) 2H+ (aq) + OH- (aq) • Cathode • 2H+ (aq) + 2e - H2 (g) Reduction • Anode • 4OH- (aq)  O2 (g) +2H2O (l) + 4e - Oxidation

Position in Electrochemical Series • The lower the metal ion in the series, the more readily it reduces (gains electrons) to discharge at the cathode. • Ex: Electrolysis of NaOH (aq)— • At cathode (negative electrode) • H2 or Na discharged? • H2 is lower in the series, so it is discharged! • Ex: Electrolysis of CuSO4 (aq)— • At anode (negative electrode) • H2 ions or Cu discharged? • Cu is lower in the series, so it is discharged! • H2 is usually reduced at the cathode unless the electrolyte solution contains easily reduced ions like Cu 2+ or Ag + • O2 is usually oxidized at the anode unless the electrolyte solution contains easily oxidized ions like Br - or I-

What will the products of the following aqueous electrolysis be?

Concentration • The more concentrated an ion in solution, the greater the production of the substance. • NaCl solution. • Both O2 and Cl2 evolved at positive electrode. • Dilute NaCl (aq) • Mainly O2 evolved. • Concentrated NaCl (aq • Mainly Cl2 evolved.

Nature of the Electrode • Normally safe to assume electrode is inert (graphite or platinum) and is not involved in the reaction. • If copper electrodes are used in the electrolysis of an aqueous CuSO4 solution, the positive electrode will be oxidized to release electrons and form Cu2+ while copper is simultaneously reduced at the negative electrode to form copper. The concentration of Cu2+ will remain constant.

Copper anode dissolves to produce Cu2+ ions which are then deposited on to the cathode Electroplating • Example—Copper plating • Anode • Copper metal • Cathode • Metal to be copper plated • Electrolyte • Copper II sulfate

Factors Affecting the Quantity of Products • # of electrons • Depends on current (flow of electric charge) and the amount of time it flows. • Time held constant • 2 x current=2 x electrons flowing = 2 X product forming • Current held constant • 2 x time=2 x electrons flowing = 2 X product forming • Charge on ion • Na+(l) +1e-  Na (l) • Formation of one mole of Na requires 1 mole of electrons • Pb2+(l) + 2e-  Pb (l) • Formation of one mole of lead requires 2 moles of elections.

Electrochemistry