Unit 5: Electrochemistry

Electrochemical reaction

• Transfer of electrons from one substance to another.
  2Al (s) + 3 Cu²⁺(aq) → 2Al³⁺(aq) + 3Cu (s)

• OXIDATION: loss of electrons
  Al → Al³⁺ + 3e^-

  • When a substance becomes oxidized, it becomes more positively charged because it is losing electrons (negatively charged).

• REDUCTION: gain of electrons
  Cu²⁺ + 2e^- → Cu

  • When a substance becomes reduced, it becomes more negatively charged because it is gaining electrons.

REDOX reactions

Every reduction reaction must be accompanied by an oxidation reaction since electrons must be transferred somewhere.
2Al (s) + 3 Cu²⁺(aq) → 2Al³⁺(aq) + 3Cu (s)

• OXIDIZING AGENT: reactant reduced (gains e^-) during a reaction.
  → Cu²⁺ is the OXIDIZING AGENT because it causes Al to become oxidized.

• REDUCING AGENT: reactant oxidized (loses e^-) during a reaction.
  → Al is the REDUCING AGENT because it causes Cu²⁺ to become reduced.

Oxidation number

The apparent charge that the atom would have if all of the bonds to the atom were completely ionic.

1. Atoms in elemental form = 0 (ex. N₂, Na, O₂)

2. Simple (monoatomic) ions = the charge on ion (ex. Fe²⁺ has an oxidation number of +2)

   • Li⁺, Na⁺, K⁺, and all other group 1 ions have oxidation number of +1.

   • Ca²⁺, Ba²⁺, Mg²⁺, and all of the group 2 ions have oxidation numbers of +2.

   • F^-, Cl^-, Br^-, I^- (halogens) are usually -1.
     There are exceptions, especially in covalent compounds.

3. Hydrogen = +1 (except in metallic hydrides such as NaH or BaH₂ where it is -1).

4. Oxygen = -2 (In peroxides, H₂O₂, it is -1)

5. Oxidation numbers of other atoms are assigned so that the sum of the oxidation numbers (positive and negative) equals the net charge on the molecule or ion.

Rule of OXIDATION/REDUCTION

• OXIDATION = loss of electrons = increase in oxidation number.

  • If the number of attached oxygen atoms increases → Oxidation

• REDUCTION = gain of electrons = decrease in oxidation number.

  • If the number of attached oxygen atoms decreases → Reduction

Predicting Spontaneous Reactions

1. In general, metals (exceptions Cu, Ag, Hg, and Au) are found in the bottom right half of the table (reducing agents).

2. In general, halogens and oxyanions (oxygen-containing anions) are found in the upper left half of the table (oxidizing agents).

3. Some metals, such as Fe, Sn, Cr, Hg, and Cu, have more than one common oxidation number and thus have more than one half-reaction.

4. Some ions (Cu⁺, Sn²⁺, Fe²⁺) appear on both sides of the table and can behave as oxidizing agents or reducing agents.

5. H₂O₂ can be an oxidizing agent or a reducing agent.

6. Half-reactions, including H+, must be carried out in acidic solutions.
   → Ex: MnO₄^- + 8H⁺ + 5e^- → Mn²⁺ + 2H₂O (acidic permanganate solution)

• OXIDIZING AGENTS gain electrons, tend to be CATIONS (+) OR NONMETALS.

  • Stronger oxidizing agents are located on the upper left, have a greater tendency to gain electrons (reduce).

• REDUCING AGENTS lose electrons, tend to be ANIONS (-) OR METALS.

  • Stronger reducing agents are located on the lower right, have a greater tendency to lose electrons (oxidize).

• Spontaneous reactions occur when there is:

  • an oxidizing agent (reduction) and a reducing agent (oxidation), and

  • the oxidizing agent must be above the reducing agent.

Balancing Half Reactions

1. Balance all major atoms other than oxygen and hydrogen.

2. Balance OXYGENS by adding water (H₂O) molecules.

3. Assume that solutions are acidic and balance HYDROGENS by adding H⁺.

4. Balance the CHARGE by adding electrons (e^-).

5. (If BASIC solutions) Add equal numbers of hydroxide ions (OH^-) to both sides of the equation and cancel out the H⁺ as water.

• Balanced redox reactions DO NOT SHOW ELECTRONS and the NUMBER OF ATOMS AND THE TOTAL CHARGE are balanced on both sides of the equation.

• DISPROPORTIONATION: same chemical undergoes oxidation and reduction.

Balancing Redox Equations Using Oxidation Numbers

1. Calculate the changes in oxidation numbers for the elements that undergo oxidation and reduction.

2. Balance increase and decrease in oxidation number change.

3. Balance oxygen with H₂O and hydrogen with H⁺.

Redox Titrations - Oxidizing agents

• One of the most common oxidizing agents used in redox titrations is Acidic Potassium Permanganate, KMnO₄. You can acidify the reaction flask by adding concentrated acid such as sulphuric acid (H⁺ source).
  MnO₄^- + 8H⁺ + 5e^- → Mn²⁺ + 4H₂O; E° = 1.49V

• The K⁺ in KMnO₄ is left out because it is a spectator ion.

• In permanganate titrations, the MnO₄^- acts as its own indicator.

• MnO₄^- is purple while Mn²⁺ is colourless. When MnO₄^- (purple) is added to a reducing agent, it is converted to Mn²⁺ (colourless). Once all the reducing agent has been oxidized by MnO₄^-, any additional MnO₄^- will remain purple.

Redox Titration - Reducing agents

• Commonly used reducing agent is NaI or KI (I^- is a RA). A large number of substances can oxidize I^- to I₂ according to the half-reaction:

  • Unknown OA + e^- → Reduced form of OA

  • 2I^- → I₂ + 2e^-

• Titrations involving I^- are indirect and involve 2 consecutive steps:

  1. I^- is added in excess and is oxidized to I₂ by the unknown oxidizing agent.

  2. The I₂ produced in the first step is titrated with a reducing agent of known concentration, such as thiosulphate ion S₂O₃^2- according to this reaction:
     2S₂O₃^2- + I₂ → S₄O₆^2- + 2I^-

• The concentration of the I^- can be determined from the [I₂], and this is used to determine the concentration of the unknown oxidizing agent.

• I₂ that is produced is a brown colour in the solution. Before any starch is added, enough RA (e.g., S₂O₃^2-) is added until the brown colour is faint.

• Starch is used as an indicator for the I₂ titration. The starch-I₂ complex produces a dark blue solution. When I₂ is titrated with S₂O₃^2-, the dark blue complex disappears (because it turns to I^-). The endpoint of the titration is indicated by the complete disappearance of the blue colour.

• Adding starch at the start of the titration step will interfere with the titration process because I₂ binds fairly tightly to starch.

Electrochemical cells

An ELECTROCHEMICAL CELL (galvanic or voltaic cell) is a chemical system in which a SPONTANEOUS redox reaction can produce useful electrical work.

Cu²⁺ + Zn → Cu + Zn²⁺

• Cu²⁺ is reduced to Cu, and Zn is oxidized to Zn²⁺.

• A spontaneous reaction will occur when zinc metal is placed into a solution of CuSO₄; however, no usable energy is produced because the energy is lost as heat.

• Can be used to produce electricity if the two half-reactions are made to occur in separate containers or CELLS.

Electrodes

• ANODE is the electrode where OXIDATION occurs.

• CATHODE is the electrode where REDUCTION occurs.

• AN OX CARED

  • ANode → OXidation

  • CAthode → REDuction

Cu²⁺ + Zn → Cu + Zn²⁺

• Zinc metal loses electrons and becomes oxidized to zinc ions, Zn²⁺, at the ANODE, so it decreases in mass.

• Copper ions, Cu²⁺, gain electrons and are reduced to copper metal at the CATHODE, so it increases in mass.

External Circuit

• Electrons flow through the wire from the anode to the cathode.

• “A comes before C in the alphabet”

• Electrons produced at the zinc electrode (anode) flow through the wire to the copper electrode (cathode) where they reduce copper ions. The electrons will continue flow through the wire until the redox reaction reaches equilibrium.

Connecting Cells

A SALT BRIDGE composed of a concentrated solution of strong electrolyte, usually KNO₃ or KCl, or a porous membrane connects the two cells to complete the circuit.

• The salt bridge keeps the two compartments electrically neutral by allowing ions (ex. K⁺ and NO₃ from KNO₃) to migrate between the cells.

• Positive Zn²⁺ and Cu²⁺ ions (cations) are attracted to the cathode.

• Negative NO₃^- and SO₄^2- ions (anions) are attracted to the anode.

Voltage (Electrical potential)

WORK DONE PER ELECTRON TRANSFERRED.

• The tendency of electrons to flow in an electrochemical cell.

• Since electrons cannot flow in an isolated half-cell, an individual half-cell voltage cannot be determined. However, the difference in electrical potentials between two half-cells can be measured.

• A ZERO-POINT is ABITRARILY defined on the voltage scale. The voltage for the HYDROGEN HALF-CELL is defined to be:
  2H⁺ + 2e^- → H₂ ; E° = 0.00V

STANDARD STATE

An electrochemical cell is said to be at the STANDARD STATE if

1. it is at 25°C, and

2. all gases are at 101.3 kPa (1atm), and

3. all elements are in their standard states (normal phases at 25°C), and

4. all solutions involved in the cell have a concentration of 1.0 M.

STANDARD REDUCTION POTENTIALS

• Positive reduction potentials indicate that the half-reaction has a greater tendency to be reduced than the hydrogen half-cell.

• Negative reduction potentials indicate that the half-reaction has a lower tendency to be reduced than the hydrogen half-cell. The oxidation reaction will occur when these half-reactions are connected to a hydrogen half-cell.

• The value of an oxidation potential is the negative of the reduction potential (opposite signs).

• The electrochemical cell potential is determined by adding the half-cell voltages for the reduction potential and oxidation potential.

  • If E° is POSITIVE for a redox reaction, the reaction is expected to be SPONTANEOUS.

  • If E° is NEGATIVE for a redox titration, the reaction is NON-SPONTANEOUS.

SURFACE AREA of the electrode

The SURFACE AREA of an electrode does not affect the voltage of the cell. This is because the CONCENTRATION OF THE SOLID IS CONSTANT.

• Increasing the surface area of a solid electrode has no effect on the concentration of the solid.

• No equilibrium shift and no change in half-cell potential.

• Increasing the surface area of an electrode increases the AMERAGE (Rate of electron flow), but not the voltage.

• Increasing the surface are of the electrode, also increases the length of time that the cell can operate.

Le Chatelier’s Principle

Cu²⁺ + 2e^- ⇌ Cu_(s) ; E° = +0.34 V

• If [Cu²⁺] > 1.0 M, the equilibrium shifts right and the reduction potential increases.
  → E > +0.34 V

• If [Cu²⁺] <1.0 M, the equilibrium shifts left and the reduction potential decreases.
  → E < +0.34 V

Notice that the “°” is omitted from the E° symbol since it is NOT AT STANDARD STATE.

Electrochemical cells NOT AT EQUILIBRIUM

2Ag⁺ + Cu → 2Ag + Cu²⁺ ; E° = +0.46 V

Initially, this reaction has a great tendency to form products. As the cell operates, using up reactants and forming products, two effects are found:

• REDUCTION REACTION:

  • 2Ag⁺ + 2e^- → 2Ag ; E < + 0.80 V

  • The [Ag⁺] decreases as the reaction occurs, and so the reduction potential decreases.

• OXIDATION REACTION:

  • Cu²⁺ + 2e^- ← Cu ; E° > + 0.34 V

  • The [Cu²⁺] increases as the reaction occurs, so the reduction potential Cu²⁺ + 2e^- → Cu increases. The tendency for reduction increases as the cell operates.

When an electrochemical cell reaches EQUILIBRIUM, the VOLTAGE OF THE CELL IS 0.00 V.

SELECTING PREFFERED REACTIONS

• STRONGEST OXIDIZING AGENT will become reduced.

• STRONGEST REDUCING AGENT will become oxidized.

Note: Any ion capable of being reduced will be a SPECTATOR ION if there is another ion in the same solution which has a greater tendency to be reduced. Similarly, any ion capable of being oxidized will be a spectator ion if there is another ion in the same solution which has a greater tendency.

CORROSION OF METALS

Rusting is a common electrochemical reaction resulting from the oxidation of iron metal (Fe) by oxygen (O₂) in the air (CORROSION = oxidation of a metal).

When iron is exposed to the atmosphere in the presence of water, the following spontaneous reactions occur.

1. Fe is oxidized at the anode (oxygen-poor) region of the metal.
   Fe_(s) → Fe²⁺ + 2e^-

2. Electrons from Fe oxidation migrate to the region, acting as a cathode.

3. O₂ is reduced at the cathode (oxygen-rich) region at the outer surface.
   ½ O₂ + H₂O + 2e^- ⇌ 2OH^-

4. Fe²⁺ oxidized to Fe³⁺, rust (Fe₂O₃) forms.
   Fe_(s) + ½ O_2(g) + H₂O_(l) → Fe(OH)_2(s)

   • Fe(OH)₂ is oxidized to a complex mixture of Fe₂O₃ and H₂O by the O₂ in the air. “Rust” is just Fe₂O₃.xH₂O, where “x” is the variable.

   • “x” explains the different colours rust can have (red, brown, yellow, black) since different numbers of water molecules attached to the Fe₂O₃ will affect the colour of the component.

• No salt bridge is needed because Fe²⁺ forms a precipitate with OH^-.

Rate of corrosion

Rate of corrosion changes when iron is in contact with different metals.

• When the metal is a stronger reducing agent than Fe (more readily oxidized), it is preferentially oxidized.

• If the metal is a weaker reducing agent than Fe, such as Cu, it acts as a conductor of electrons and speeds up the rate of Fe oxidation.

Isolating the metal from the environment

• Apply a protective layer to the metal, such as paint, plastic, or grease. If oxygen and water cannot reach the metal, it will not corrode.

• Apply a metal that is corrosion-resistant to the surface of the original metal. Tin can be applied to the surface of steel cans. The tin is quickly oxidized to produce a thick layer of tin oxide, which protects the underlying metal from further corrosion.

Electrochemical methods

• Cathodic protection: protecting a substance from unwanted oxidation by connecting it to a stronger reducing agent.

• Sacrificial electrodes using magnesium, for example, are used to protect underground tanks or pipes. Replacing the sacrificial electrodes is more cost-effective than replacing the iron objects they are protecting.

• Change the condition of the surroundings: we can prevent corrosion by altering the reduction reaction that causes Fe to become oxidized
  ½ O₂ + H₂O + 2e^- ⇌ 2OH^-

  • If oxygen is removed from the solution, the reduction potential decreases.

  • Increase [OH^-] ions also causes the reduction potential to decrease (ie, alkaline environment).

Electrolysis

The process of supplying electrical energy to a molten ionic compound or a solution containing ions to produce a chemical change.

• ELECTROLYTIC CELLS involve NONSPONTANEOUS redox reactions that require energy to occur (ENDOTHERMIC).

• ELECTROCHEMICAL CELLS involve SPONTANEOUS redox reactions that release energy (EXOTHERMIC).

Electrolysis of Molten Binary Salts

• Binary salts are made up of only two different elements (e.g., NaCl, KBr, MgI₂, AlF₃, etc.)

• Molten binary salts, such as NaCl, contain only two kinds of ions.

  • Molten NaCl is NaCl_(l), only Na⁺ and Cl^- ions present.

  • NaCl_(aq) is a solution that contains Na⁺, Cl^-, and H₂O.

• No need for a salt bridge to keep the reactants separated since no spontaneous reaction will occur, although a barrier can be used to prevent products from mixing.

Electrolysis of Aqueous Solutions

Electrolysis of aqueous solutions must consider the possibility that H₂O may be oxidized and/or reduced.

• Electrolysis of aqueous solutions containing Cl^- or Br^- will produce Cl₂ or Br₂ at the anode.

• Electrolysis of aqueous solutions containing Fe²⁺, Cr³⁺, or Zn²⁺ will produce Fe, Cr, or Zn at the cathode.

Determining the overall reaction

1. The STRONGEST OXIDIZING AGENT becomes reduced and the STRONGEST REDUCING AGENT become oxidized.

2. The preferred overall reaction will be the one requiring the LEAST voltage input.

OVERPOTENTIAL EFFECT

Higher potential than calculated must be supplied to cause electrolysis.

• Causes include the nature of the electrodes, temperature, current, and time.

• Difference between actual potentials required for electrolysis and the calculated potential is termed HALF-CELL OVERPOTENTIAL.

• Overpotentials vary for each half-cell but the difference in potentials is very large for oxidation and reduction of neutral water.

ELECTROPLATING

An electrolytic process in which a metal is reduced or “plated out” onto an object connected to the cathode (e.g., coat something with metal).

• The CATHODE is made out of the material that will receive the metal plating.

• The ELECTROPLATING SOLUTION contains ions of the metal that is to be “plated” onto the cathode.

• The ANODE may be made of the same metal that is to be “plated out” onto the cathode (This is normal, but an inert electrode can also be used).