Hard Oxidation-Reduction Reactions Quiz

Chapter Objectives

• Describe at least three types of corrosion and identify chemical reactions responsible for corrosion.
• Define oxidation and reduction.
• Write and balance half-reactions for simple redox processes.
• Describe the differences between galvanic and electrolytic cells.
• Use standard reduction potentials to calculate cell potentials under both standard and nonstandard conditions.
• Use standard reduction potentials to predict the spontaneous direction of a redox reaction.
• Calculate the amount of metal plated, the amount of current needed, or the time required for an electrolysis process.
• Distinguish between primary and secondary batteries.
• Describe the chemistry of some common battery types and explain why each type of battery is suitable for a particular application.
• Describe at least three common techniques for preventing corrosion.

Corrosion

• Corrosion is the degradation of metals by chemical reactions with the environment.
• Uniform corrosion occurs evenly over a large portion of the surface area of a metal.
• Galvanic corrosion occurs when two different metals contact each other in the presence of an appropriate electrolyte.
• Crevice corrosion occurs when two pieces of metal touch each other, leaving a small gap or crevice between the metals.
• Different metals corrode differently.
• Aluminum has a greater tendency to corrode than iron, but corrosion of aluminum is not problematic compared to iron.
• The aluminum oxide corrosion product forms a protective layer on the surface of aluminum metal.
• The iron oxide corrosion product flakes off the surface of iron, exposing fresh iron to corrosion.
• Corrosion occurs in a variety of forms. The chain shows uniform corrosion. The grill cover shows crevice corrosion where the handle is attached.

Oxidation-Reduction Reactions and Galvanic Cells

• Special conditions must be present before iron reacts with oxygen to form iron(III) oxide.
• Rust formation is a slow process, so the basics of electrochemistry must be investigated using more easily observed reactions.
• Reactions that transfer electrons between reactants are known as oxidation-reduction or redox reactions.
• Oxidation is the loss of electrons from some chemical species.
• Reduction is the gain of electrons to some chemical species.
• For an oxidation-reduction reaction to occur, one reactant must be oxidized and one reactant must be reduced.
• Oxidation cannot occur without reduction.
• When copper wire is placed in a silver nitrate solution, a redox reaction occurs.
• A reaction is observed to occur because the solution changes color and crystals form on the surface of the copper wire.
• The solution's blue color is indicative of Cu²⁺ ions in solution.
• Cu²⁺ is formed when a copper atom loses two electrons.
• The copper metal is oxidized.
• The crystals forming on the surface of the copper wire are silver metal.
• Silver is formed when a silver cation gains an electron.
• The silver cation is reduced.
• For the reaction between silver cation and copper metal, two half-reactions are written.
• One for the oxidation of copper and one for the reduction of silver.
• Neither half-reaction can occur without the other.
• The half-reactions as written indicate that Ag⁺ only accepts one electron whereas Cu loses two electrons.
• The electron transfer must balance, so the reduction half-reaction is multiplied by 2.
• Add the two half-reactions together, the electrons cancel out, leaving the net ionic equation for the redox reaction.
• The species undergoing oxidation is referred to as a reducing agent.
• The Cu was oxidized and is the reducing agent.
• The Cu facilitated the reduction of Ag⁺ by losing electrons.
• The species undergoing reduction is referred to as an oxidizing agent.
• The Ag⁺ was reduced and is the oxidizing agent.
• The Ag⁺ facilitated the oxidation of Cu by gaining electrons.

Building a Galvanic Cell

• A galvanic cell is any electrochemical cell in which a spontaneous chemical reaction can be used to generate an electric current.
• The name electrochemistry comes from the observation of electric currents in galvanic cells.
• To harness electricity from a galvanic cell, each half-reaction is prepared separately in half-cells.
• Cu metal immersed in Cu²⁺ solution is one half-cell.
• Ag metal immersed in Ag⁺ solution is the second half-cell.
• A salt bridge is crucial to a galvanic cell. The salt bridge allows ions to flow between each half-cell, completing the circuit.
• Current flows by the migration of ions in solution.
• To transfer current between the half-cells, a salt bridge is used.
• The salt bridge contains a strong electrolyte that allows either cations or anions to migrate into the solution where they are needed to maintain charge neutrality.
• A metal wire cannot transport ions and cannot be used.
• For a salt bridge composed of NH₄Cl: NH₄⁺ will flow into the Ag⁺ beaker to offset the removal of Ag⁺ from solution. Cl⁻ will flow into the Cu²⁺ beaker to offset the production of Cu²⁺ in solution.
• The circuit is completed by connecting wires to each metal strip.
• A voltage potential of 0.46 V will be measured for the described cell.

Terminology for Galvanic Cells

• Electrodes are the electrically conducting sites at which either oxidation or reduction occurs.
• Oxidation occurs at the anode.
• Reduction occurs at the cathode.
• Cell notation - a shorthand notation for the specific chemistry of an electrochemical cell.
• Cell notation lists the metals and ions involved in the reaction.
• A vertical line, |, denotes a phase boundary.
• A double vertical line, ||, denotes a salt bridge.
• The anode is written on the left, the cathode on the right.
• General form of cell notation anode | anode electrolyte || cathode electrolyte | cathode

Standard Reduction Potentials

• To compare the oxidation-reduction trends of species used in electrochemistry, all half-cell potentials are written as reductions.
• A table of standard reduction potentials lists the potential of any half-reaction when connected to a SHE.
• All materials are 1 M for aqueous species and 1 atm partial pressure for gases.
• Although the half-reactions are listed as reductions in the table, one half-reaction in any electrochemical cell must be an oxidation and, therefore, reversed from what appears in the table.
• The cell potential sign must be changed when writing the half-reaction as an oxidation.
• Some half-reactions have positive potentials, whereas others have negative potentials.
• All potentials are measured with a SHE connected to the positive terminal.
• If the voltage is positive, the SHE is the anode, the oxidation site.
• A positive standard reduction potential means the half-reaction proceeds as written (reduction occurs).
• If the voltage is negative, the SHE is the cathode, the reduction site.
• A negative standard reduction potential means the half-reaction proceeds as an oxidation.
• The tendency for the chemicals involved in a half-reaction to be an oxidation or reduction depends on the value of the reduction potential.
• A large, positive value for the standard reduction potential implies the substance is reduced readily and a good oxidizing agent.
• A large, negative value for the standard reduction potential implies the substance is oxidized readily and a good reducing agent.
• For a galvanic cell, the half-reaction with the more positive reduction potential will be the cathode. The half-reaction with the more negative reduction potential will be the anode.
• The standard reduction potential for any pair of half-reactions, E°cell, is calculated from the standard reduction potentials for the cathode and anode.
• E°cell = E°red - E°ox

Example Problem 13.1

• Use standard reduction potentials to identify the anode and cathode and determine the cell potential for a galvanic cell composed of copper and iron. Assume standard conditions.
• Confirm that the potential of the following galvanic cell is 0.462 V: Cu(s) | Cu²⁺ (1 M) || Ag⁺ (1 M) | Ag(s)

Nonstandard Conditions

• The cell potential at nonstandard conditions is calculated using the Nernst equation.
• E=E° - (RT/nF) lnQ
• Q is the reaction quotient, F is the Faraday constant, and n is the number of electrons transferred in the reaction.
• F = 96,485 J V⁻¹ mol⁻¹ or 96,485 C mol⁻¹

Example Problem 13.2

• Assume that you have a cell that has an iron(II) concentration of 0.015 M and an H⁺ concentration of 1.0 × 10⁻³ M. The cell temperature is 38°C, and the pressure of hydrogen gas is maintained at 0.04 atm. What would the cell potential be under these conditions?

Cell Potentials and Free Energy

• Corrosion is a spontaneous process and has a negative Gibbs free energy change.
• The Gibbs free energy change for an electrochemical reaction can be calculated from the standard reduction potential.
• ΔG° = -nFE°
• n is the number of electrons transferred and F is Faraday's constant

Example Problem 13.3

• Suppose that we wish to study the possible galvanic corrosion between zinc and chromium, so we set up the following cell: Cr(s) | Cr²⁺(aq) || Zn²⁺(aq) | Zn(s).
• What is the chemical reaction that takes place and what is the standard free energy change for that reaction?

Equilibrium Constants

• The cell potential can be used to calculate the equilibrium constant for an electrochemical reaction.
• E = (RT/nF) ln K
• n is the number of electrons transferred, R is the universal gas law constant, and F is Faraday's constant.
• The relationship between the cell potential and the equilibrium constant can be re-written in terms of the common (base 10) log.
• E = (2.303RT/nF) log K
• The equation can be simplified for reaction carried out at standard temperature, 25°C (298 K).
• E° = (0.0592/n) log K

Example Problem 13.4

• In a process called flash electroplating, a current of 2.50 × 10³ A passes through an electrolytic cell for 5.00 minutes. How many moles of electrons are driven through the cell?

Example Problem 13.5

• Suppose that a batch of parts is plated with copper in an electrolytic bath running at 0.15 V and 15.0 A for exactly 2 hours. What is the energy cost of this process if the electric utility charges the company $0.0500 per kWh?

Calculations Using Masses of Substances in Electrolysis

• A knowledge of current, how long the current flows, stoichiometry, and the number of electrons required to reduce a metal cation are used to answer the following questions.
• How much material is plated given a specific current for an allotted time or electrical energy expenditure?
• How long must a given current to pass through the cell to yield a desired mass of plated material?

Example Problem 13.6

• An electrolysis cell that deposits gold (from Au³⁺(aq)) operates for 15.0 minutes at a current of 2.30 A. What mass of gold is deposited?

Example Problem 13.7

• Suppose that you have a part that requires tin coating. You've calculated that you need to deposit 3.60 g of tin to achieve an adequate coating. If your electrolysis cell (using Sn²⁺) runs at 2.00 A, how long must you operate the cell to obtain the desired coating?

Batteries in Engineering Design

• Lithium-ion batteries have recently come under scrutiny because of a high-profile fire during the first commercial flight of the Boeing Dreamliner 787.
• What factors led to the Dreamliner's problems?
  • The Dreamliner requires higher levels of electrical power than conventional aircraft due to replacing the traditional hydraulic system for controlling the flight with an all-electronic alternative.
  • Increased fuel efficiency was a major goal, so it was important to minimize weight. Lithium-ion batteries have high energy capacity and low weight.
• The anode is a form of graphite into which lithium atoms have been incorporated, or “intercalated."
• The cathode is cobalt oxide that also has lithium incorporated into it.
• Two reasons energy density is high:
  • Li⁺/Li has one of the largest standard reduction potentials
  • Both lithium and carbon are relatively light
• High temperatures cause lithium-ion batteries to degrade fairly quickly, so engineering designs that use them generally must account for this.

Primary Cells

• Single-use batteries that cannot be recharged are primary cells.
• The most prevalent type of primary cell is the alkaline battery.
• An alkaline battery has a zinc electrode at which oxidation occurs.
• The cathode is derived from manganese(IV) oxide.
• The chemistry of an alkaline dry cell battery. The net reaction is shown above.
• The alkaline battery is termed a dry cell because the KOH electrolyte is in the form of a paste or gel.
• Lithium batteries are small and long lasting. They are used for medical devices like pacemakers.
• Lithium is the anode
• Manganese(IV) oxide is the cathode, but in this case the MnO₂ reacts with lithium ions.
• The lithium battery provides a stable current and electrical potential for long periods.
• Zinc-air batteries are also primary cells.
• Zinc is the anode.
• Oxygen reacts at the cathode.

Secondary Cells

• Rechargeable batteries are secondary cells or secondary batteries.
• Nickel-metal-hydride batteries are an example of secondary cells.
• The anode reaction is MH(s) + OH⁻(aq) → M⁺ + H₂O(l) + e⁻.
• The complex cathode reaction can be represented as NiO(OH)(s) + H₂O(l) + e⁻ → Ni(OH)₂(s) + OH⁻(aq)
• Nickel is converted from NiO(OH) to Ni(OH)₂ at the cathode.
• The cathode and anode materials are separated by an insulator throughout the battery.
• Water reacts with a metal alloy to form an absorbed hydrogen atom and hydroxide ion at the anode.
• The negative collector is in contact with the anode material.
• Nickel-metal-hydride batteries have become popular as rechargeable cells.
• The lead-acid storage battery found in cars is a secondary cell.
• The anode for a lead-acid battery is lead metal.
• The cathode for a lead-acid battery is lead oxide.
• The lead-acid storage battery consists of Pb anodes alternating with PbO₂ cathodes, all immersed in sulfuric acid.

Fuel Cells

• A fuel cell is a voltaic cell in which the reactants can be supplied continuously and the products of the cell reaction are continuously removed.
• Most common type is based on the reaction of hydrogen and oxygen to produce water.
• Oxygen is reduced at the cathode.
• Hydrogen is oxidized at the anode.

Limitations of Batteries

• Corrosion is a major cause for the loss of performance in batteries.
• Protective plating of materials used in batteries is an attempt to limit the performance-diminishing effects of corrosion on batteries.

Electrolysis

• Electrolysis is the process of passing an electric current through an ionic solution or molten salt to produce a chemical reaction.
• Electrolytic cells are divided into two categories based on the nature of the electrodes used.
  • Passive electrolysis: the electrodes are chemically inert materials that simply provide a path for electrons.
  • Active electrolysis: the electrodes are part of the electrolytic reaction.
• Electrolysis changes the polarity of the electrodes in a system. For reduction, electrons are forced to the cathode. The cathode becomes the negative electrode. For oxidation, electrons are pulled from the anode. The anode becomes the positive electrode.
• In electrolysis, an external source of current drives a redox reaction that would otherwise not be spontaneous. The flow of ions completes the circuit.

Passive Electrolysis in Refining Aluminum

• Electrolysis provides the means to overcome the nonspontaneous reaction to separate aluminum from its oxide.
• The Hall-Heroult refining process uses carbon electrodes as inert sites for passive electrolysis.

Active Electrolysis and Electroplating

• The process of depositing a thin coat of metal on another metal by using electrolysis is electroplating.
• In some cases, the thin coating is cosmetic, or to provide some vital functionality for the coated piece, such as corrosion resistance or desirable conductive properties.
• Silver is plated onto electrical devices because silver is a good conductor and resistant to corrosion.
• The object being electroplated is the cathode.
• Anode Ag(s) + 2CN⁻(aq) → Ag(CN)₂⁻(aq) + e⁻
• Cathode Ag(CN)₂⁻(aq) + e⁻ → Ag(s) + 2CN⁻(aq)
• Silver is transferred from the anode to the cathode, coating the cathode in a thin layer of silver.
• The zero cell potential is not critical since an external current drives electrolysis.
• Barrel plating is often used to apply coatings to small parts.

Electrolysis and Stoichiometry

• For electroplating, it can be vitally important to use carefully controlled amounts of materials.
• Controlling the flow of electrons (current) in an electroplating operation provides a method to accurately limit the amount of material deposited.
• Electroplating is often used to prevent galvanic corrosion in an electrical apparatus in places where different metals come into contact with one another.

Current and Charge

• When current is measured in an electric circuit, the observation is the flow of charge for a period of time.
• The unit of current, the ampere (A), is defined as one coulomb per second: 1 A = 1 C s⁻¹.
• If a known current flows through a circuit for a known time, the charge can be easily calculated.
• Charge = current × time, Q = I × t
• Using Faraday's constant, F = 96,485 C mol⁻¹, and the calculated charge, the number of moles of electrons that pass through the circuit can be calculated.
• If the number of electrons required to reduce each metal cation is known, the number of moles of material plated can be calculated.
• Electricity use is often measured in terms of power. The SI unit for power is the watt (1 watt = 1 J s⁻¹).
• Electrical utilities normally determine consumption in kilowatt-hours, kWh (1 kWh = 3.60 × 10⁶ J).

Example Problem 13.4

• In a process called flash electroplating, a current of 2.50 × 10³ A passes through an electrolytic cell for 5.00 minutes. How many moles of electrons are driven through the cell?

Example Problem 13.5

• Suppose that a batch of parts is plated with copper in an electrolytic bath running at 0.15 V and 15.0 A for exactly 2 hours. What is the energy cost of this process if the electric utility charges the company $0.0500 per kWh?

Example Problem 13.6

• An electrolysis cell that deposits gold (from Au³⁺(aq)) operates for 15.0 minutes at a current of 2.30 A. What mass of gold is deposited?

Example Problem 13.7

• Suppose that you have a part that requires tin coating. You've calculated that you need to deposit 3.60 g of tin to achieve an adequate coating. If your electrolysis cell (using Sn²⁺) runs at 2.00 A, how long must you operate the cell to obtain the desired coating?