Topics 9 & 19

Redox Processes

oxidation

the loss of one or more electrons from a substance

Also, in its simplest level, can be considered as a reaction in which a substance combines with oxygen (gains oxygen) or removes hydrogen

reduction

the gain of one or more electrons

In its simplest level, can also be considered as a reaction in which a substance removes oxygen or add hydrogen

What do oxidation and reduction always do?

They always occur together (OIL RIG - Oxidation Is Loss & Reduction Is Gain)

oxidation number

a concept providing a way to keep track of electrons in redox reaction according to certain rules

• Covalent compounds - more electronegative element forming negative ion

• Oxygen always has oxidtaion number of -2 except in following two situations:

  • H₂O₂ (oxygen is assigned -1 in peroxides)

  • OF₂/F₂O (oxygen is assigned +2 in this compound)

• Hydrogen always has an oxidation number of +1 except when it’s part of a metal hydride:

  • Ex: NaH, CaH₂; (hydrogen is assigned -1)

• Fluorine always has an oxidation number of -1

carbonate

CO₃^-2

sulfate

SO₄^-2

What three elements always have the same charge (and what are their charges), even though they are in the d-block?

Zn (+2), Cd (+2), Ag (+1)

phosphate

PO₄^3-

nitrate

NO₃^1-

hydroxide

OH^-1

chlorate

ClO₃^-1

ammonium

NH₄⁺¹

How to balance redox reactions in acidic solution?

Don’t do the thing with OH^-

How to balance redox reactions in basic solution?

H⁺ ions appearing in final equations must be neutralized by adding an equivalent number OH^-1 ions (forming H₂O molecules) to each side of the equation

oxidizing agent

a substance that is able to oxidize other substances (undergoes reduction itself)

reducing agent

a substance that is able to reduce other substances (undergoes oxidation itself)

voltaic cell

a device used to obtain electrical energy from a spontaneous chemical reaction

• batteries are examples of voltaic cells

• also known as a galvanic cell

Voltaic cell diagram

• each beaker is referred to as a half-cell

• oxidation occurs in the half-cell on the left side of the diagram

• reduction occurs in the half-cell on the right side of the diagram

• electrodes: the solid metals that serve to transfer electrons from one half-cell to another (a conducting wire connects the electrodes)

  • anode: the negative electrode in the half-cell where oxidation occurs

  • cathode: the positive electrode in the half-cell where reduction occurs

• salt bridge: completes the circuit by preventing the build-up of charge by allowing ions to flow from one solution to another.

  • consists of a glass tube filled with a saturated solution of KNO₃ gel with a porous plug at each end.

• voltmeter: measures the potential difference in volts (V)

electrolytic cell

a device that uses electrical energy from an external source to make a non-spontaneous chemical reaction take place

• uses include:

  • recharging batteries

  • producing metals like sodium, potassium and aluminum

  • electroplating materials

  • producing hydrogen and oxygen gas from water

• unlike a voltaic cell, an electrolytic cell doesn’t contain two separate half-cells or a salt bridge

• an electrolytic cell requires an external power source because the reaction is non spontaneous

How do the anodes and cathodes of a electrolytic cell work?

• similar to a voltaic cell in that oxidation still occurs at the anode and reduction still occurs at the cathode

• different than a voltiac cell in that the anode is positively charged and the cathode is negatively charged in an electrolytic cell

electrolyte

a substance that conducts electricity in an electrolytic cell

what’s the standard hydrogen electrode?

• consists of a platinum electrode (inert) surrounded by hydrogen gas at 100 kPa (~1 atm) of pressures

• the platinum electrode is immersed in an aqueous solution of acid in which the concentration of H⁺ ions is exactly 1 M (1 mol*dm³)

• the temp is maintained at 298 K

• assigned a standard electrode potential of 0.00 V as a reference point (like C-12 isotope is the basis of comparison for all atomic masses)

How are the standard electrode potentials of other half-cells determined?

It’s determined by connecting the half-cell to the standard hydrogen electrode using a salt bridge and connecting wire under standard conditions

what is standard electrode potential ($E^{\theta}$)

the potential (“pull“ on the electrons from the oxidizing agent) of a half-reaction under standard state conditions, as measured against the potential of the standard hydrogen electrode.

• is an intensive property that doesn’t change even when the reaction is multipllied by an integer for balancing

• but, if reaction is reversed, the sign of $E^{\theta}$ is also reversed

How to calculate cell potentials using standard electrode potentials?

• the reaction with the more positive $E^{\theta}$ will undergo reduction

• the reaction with the more negative $E^{\theta}$ will undergo oxidation

When will a reaction be spontaneous using standard electrode potentials?

a reaction will be spontaneous if $E^{\theta}$_cell is positive

Use cell potential to calculate the standard change in Gibbs energy for a redox reaction

$$\Delta G^{\theta}=-nFE_{cell}^{\theta}$$

• n = amount, in moles of electrons transferred in balanced equation

• F = Faraday’s constant (96,500 Coulombs)/(1 mol e^-)

• $E_{cell}^{\theta}$ = standard cell potential

  • 1 Volt Coulomb (VC) = 1 Joule (J)

Predict and explain the products of electrolysis of dilute aqueous solutions

• if the electrolyte is a dilute aqueous solution, the water present can also be oxidized or reduced

  • 2H₂O (l) + 2e^- → H₂ (g) + 2OH^-1 (aq) $E^{\theta}$ = -0.83 V

    • this reduction reaction will occur at the cathode (-) unless the electrolyte contains the ions of a metal with an $E^{\theta}$ > -0.83 V.

  • 2H₂O (l) → O₂ (g) + 4H⁺ (aq) + 4e^- $E^{\theta}$ = -1.23 V

    • this oxidation reaction will occur at the anode (+) unless the electrolyte contains the ions of a nonmetal with an $E^{\theta}$ > -1.23 V

  • sulfates (SO₄^-2) tend not to oxidize:

    • S₂O₈^-2 (aq) + 2e^- → 2SO₄^2- (aq) $E^{\theta}$ = +2.01 V

Describe the use of electrolysis in electroplating

• Electroplating is achieved by passing an electrical current through a solution containing dissolved metal ions and the metal object to be plated

• Uses include:

  • gold plating has decorative and practical uses. Electrical connectors are often gold plated to provide a corrosion resistant conductive surface

  • rhodium plating is used to produce “white gold“

  • zinc plating is used with “galvanized“ nails. Iron is coated or plated with zinc. Zinc is more easily oxidized than iron. A protective coating of zinc oxide is formed when the zinc oxidizes, thus protecting the underlying iron.

  • tin plating is used to protect steel food cans from reacting with their contents (tin is less easily oxiidized than iron).

Describe the differences between a primary, secondary and fuel cell

• primary (voltaic cell): an electrochemical device that converts chemical energy from spontaneous redox reactions into electrical energy

• secondary cell: the chemical reactions are reversible, so the battery can be recharged

• fuel cell: an electrochemical device that converts the chemical potential energy in a fuel directly into electrical energy

  • fuel cells differ from voltaic cells in that fuel cells require a continuous supply of oxidant/fuel, while voltaic cells contain a finite amount of reactants contained within them

Describe the main features of a hydrogen fuel cell

• The hydrogen fuel cell consists of:

  • a reaction chamber with separate inlets for hydrogen (H₂) and oxygen (O₂) gas

  • an outlet for the product, H₂O

  • an electrolyte of aqueous sodium hydroxide

  • a semi-permeable membrane (PEM) that separates the hydrogen and oxygen gases (OH^-1/H⁺ can pass through)

  • hydrogen gas is oxidized at the anode: H₂ (g) → 2H⁺ (aq) + 2e^-

  • the electrons cannot travel through the PEM, so they leave through the external circuit

  • oxygen gas is reduced at the cathode: O₂ (g) + 4e^- + 4H⁺ (aq) → 2H₂O (l)

What are benefits of hydrogen fuel cells (compared to combustion reactions)

• water is the only reaction product so no greenhouse gases are produced

• the reaction takes place at room temp

• there are no harmful oxides of nitrogen produced

What are the risks/problems of hydrogen fuel cells?

• hydrogen is a highly flammable gas so production and storage are difficult

• most hydrogen is produced as a by-product of the crude oil industry

What are the main features of a direct methanol fuel cell

• A direct methanol fuel cell has the same components as the hydrogen fuel cell except it uses methanol to provide the H⁺ ions at the anode rather than hydrogen

  • At the anode the methanol and water react to form H⁺ and electrons which flow to the cathode where oxygen and the H⁺ react to form water

  • since this reaction requires water, a dilute solution of the methanol can be used (approx. 1 mol*dm^-3)

    • This lowers the energy density, but it is still higher than hydrogen as a source in fuel cells

    • in most cells pure methanol is continuously fed into the system while water is recirculated, so the concentration of methanol remains constant

  • This type of fuel cell is not as “clean“ because it produces CO₂