Oxidation-Reduction Flashcards

Introduction to Oxidation-Reduction (Redox) Reactions

The transition from internal combustion engines to electric cars signifies a shift towards utilizing chemical energy in new ways. Redox reactions are fundamental to this, converting chemical energy into electrical energy, as seen in batteries.

Key Vocabulary

• Anode: The electrode where oxidation occurs.

• Cathode: The electrode where reduction occurs.

• Electrochemical Cell: A device that generates electricity through chemical reactions or uses electricity to drive chemical reactions.

• Electrode: A conductor through which electricity enters or leaves an object, substance, or cell.

• Electrolysis: The process of using electrical energy to drive a non-spontaneous chemical reaction.

• Electrolytic Cell: A type of electrochemical cell that requires an external source of electrical energy to drive a non-spontaneous redox reaction.

• Half-Reaction: Shows either the oxidation or reduction portion of a redox reaction, including the electrons gained or lost.

• Oxidation: The loss of electrons by an atom or ion; increase in oxidation number.

• Oxidation Number (State): A positive, negative, or neutral value assigned to an atom to track electron gain or loss.

• Redox: A reaction in which both reduction and oxidation occur.

• Reduction: The gain of electrons by an atom or ion; decrease in oxidation number.

• Salt Bridge: A component of a voltaic cell that connects the two half-cells and allows for the flow of ions, completing the circuit.

• Voltaic Cell: An electrochemical cell that uses a spontaneous chemical reaction to generate electricity.

Overview of Oxidation and Reduction

Oxidation and reduction are crucial chemical reactions with both beneficial applications (e.g., batteries) and detrimental effects (e.g., corrosion). Originally, oxidation referred to the combination of a substance with oxygen, while reduction was the loss of oxygen. Now, these terms have broader meanings and are interdependent—one cannot occur without the other.

Electron Transfer

Magnesium reacting with oxygen illustrates electron transfer. Magnesium loses two electrons, and oxygen gains two electrons to achieve a stable octet configuration.
Mg + O \rightarrow Mg^{2+} + O^{2-}

Similarly, magnesium reacts with chlorine, where magnesium loses two electrons, and each chlorine atom gains one electron to also achieve a stable octet.
Mg + Cl \rightarrow Mg^{2+} + 2Cl^{-}

Definitions

• Oxidation: Loss of electrons by an atom or ion.

• Reduction: Gain of electrons by an atom or ion.

Redox reactions always occur together; neither oxidation nor reduction can happen in isolation.

Oxidation Numbers

Oxidation numbers track electron exchange during reactions. They help determine how many electrons an atom or ion gains or loses. Oxidation is linked to an increase in oxidation number, while reduction is linked to a decrease.

Mnemonic Device

• LEO says GER:

  • LEO: Loss of Electrons is Oxidation

  • GER: Gain of Electrons is Reduction

Rules for Assigning Oxidation Numbers

1. Uncombined Elements: Have an oxidation number of zero.

   • Example: In 2Na + Cl2 \rightarrow 2NaCl, Na and Cl2, have oxidation numbers of 0.

2. Monatomic Ions: Have an oxidation number equal to their ionic charge.

   • Example: In NaCl2Na + Cl_2 \rightarrow 2NaCl , Na in NaCl has an oxidation number of +1, and Cl has -1.

3. Group 1 Metals: Always have an oxidation number of +1 in compounds.

4. Group 2 Metals: Always have an oxidation number of +2 in compounds.

5. Fluorine: Always -1 in compounds.

6. Other Halogens: Typically -1 when they are the most electronegative element in the compound.

7. Hydrogen: +1 in compounds, unless combined with a metal, then -1.

   • Examples: +1 in HCl, -1 in LiH.

8. Oxygen: Usually -2 in compounds, but +2 when combined with fluorine.

   • Examples: -2 in H2O, +2 in OF2, -1 in the peroxide ion (O_2^{2-}).

9. Sum of Oxidation Numbers in Compounds: Must be zero.

10. Sum of Oxidation Numbers in Polyatomic Ions: Must equal the charge on the ion.

Sample Problems

Problem 1: Oxidation Numbers in HNO_3

• H = +1 (Rule 5)

• O = -2, total for three O atoms = -6 (Rule 6)

• Sum of oxidation numbers is zero.
  N + (+1) + (-6) = 0
  N = +5

• Oxidation numbers: H = +1, N = +5, O = -2.

Problem 2: Oxidation Number of Chromium in Dichromate Ion (Cr2O7^{2-})

• O = -2, total for seven O atoms = -14 (Rule 6)

• Sum of oxidation numbers equals the charge on the ion, which is -2.
  2(Cr) + (-14) = -2
  2(Cr) = +12
  Cr = +6

• Oxidation number of Cr = +6.

Identifying Redox Reactions

To determine if a reaction is redox, assign oxidation numbers to each atom on both sides of the equation. If there is a change in oxidation number for any atom, the reaction is redox.

• If an uncombined element appears on one side and is in a compound on the other, the reaction is redox.

• Double replacement reactions are not redox reactions.

Examples

• Redox: Zn + HCl \rightarrow ZnCl2 + H2

• Not Redox: NaCl + AgNO3 \rightarrow AgCl + NaNO3

Identifying Oxidation and Reduction

Identify the atom that has increased in oxidation number (oxidation) and decreased in oxidation number (reduction).

Example: MnO2 + 4HCl \rightarrow MnCl2 + Cl2 + 2H2O

• Chlorine in HCl has an oxidation number of -1, while some chlorine atoms in Cl_2 have an oxidation number of 0. Thus, chlorine is oxidized.

• Manganese changes from +4 in MnO2 to +2 in MnCl2, so manganese is reduced.

Oxidizing Agents and Reducing Agents

• The substance oxidized is the reducing agent.

• The substance reduced is the oxidizing agent.

In the example above:

• Mn^{+4} is reduced by gaining electrons from Cl^{-}. Therefore, Mn^{+4} is the oxidizing agent, causing the Cl^{-} to be oxidized.

• Cl^{-} is oxidized by losing electrons to Mn^{+4}. Therefore, Cl^{-} is the reducing agent, causing the Mn^{+4} to be reduced.

Half-Reactions

Half-reactions show the oxidation or reduction portion of a redox reaction, including the electrons gained or lost.

• Reduction Half-Reaction: Shows an atom or ion gaining electrons, with a decrease in oxidation number.

  • Example: Fe^{3+}(aq) + 3e^- \rightarrow Fe(s)

• Oxidation Half-Reaction: Shows an atom or ion losing electrons, with an increase in oxidation number.

  • Example: Fe(s) \rightarrow Fe^{3+}(aq) + 3e^-

Half-reactions must follow the law of conservation of matter and charge.

Balancing Half-Reactions

1. Assign oxidation numbers to each element.

2. Write a partial half-reaction showing the change in oxidation state.

3. Add electrons to explain the change in oxidation number and conserve charge.

Example: Cu + AgNO3 \rightarrow Cu(NO3)_2 + Ag

• Partial Half-Reactions:

  • Oxidation: Cu \rightarrow Cu^{2+}

  • Reduction: Ag^+ \rightarrow Ag

• Balanced Half-Reactions:

  • Oxidation: Cu \rightarrow Cu^{2+} + 2e^-

  • Reduction: Ag^+ + e^- \rightarrow Ag

• Balance the number of electrons lost and gained:

  • Multiply the reduction equation by two to balance electrons:
    2(Ag^+ + e^- \rightarrow Ag) \Rightarrow 2Ag^+ + 2e^- \rightarrow 2Ag

• Combine the half-reactions:
  2Ag^+ + Cu \rightarrow Cu^{2+} + 2Ag

Spontaneous Reactions - Voltaic Cells

In a voltaic cell, chemical energy is spontaneously converted to electrical energy. A salt bridge connects the two containers, providing a path for ion flow between the beakers, completing the circuit.

Electrochemical Cells

• Voltaic Cell: A spontaneous chemical reaction produces a flow of electrons.

• Electrolytic Cell: Requires an electric current to force a non-spontaneous chemical reaction.

• Electrodes: Conduct electricity; oxidation or reduction occurs here.

  • Anode: Site of oxidation.

  • Cathode: Site of reduction.

Example

Zn(s) + Cu^{2+}(aq) \rightarrow Cu(s) + Zn^{2+}(aq)

Zn(s) \rightarrow Zn^{2+}(aq) + 2e^-

Cu^{2+}(aq) + 2e^- \rightarrow Cu(s)

Electrons lost during oxidation at the anode travel through a wire to the cathode, where the material is reduced.

Activity Series

The activity series (Table J) identifies the anode and cathode in a voltaic cell. The higher metal on the chart will be oxidized (anode), and the lower metal is the site of reduction (cathode).

Mnemonics: RED CAT (Reduction occurs at the Cathode), AN OX (Anode is the site of Oxidation).

Sample Problem: ZnO(s) + Pb^{2+}(aq) \rightarrow Zn^{2+}(aq) + PbO(s)

• Zinc is higher than lead on Table J, so zinc is oxidized and the anode. Pb^{2+} ions are reduced at the cathode (lead metal).

• Electrons flow from zinc to lead (anode to cathode).

Electric Potential

When a voltaic cell reacts, electrons flow from anode to cathode. A voltmeter measures the electric potential in volts. Table 9-1 shows reduction potentials compared to the standard hydrogen cell (0.00 V).

The more positive the E^0 value, the more likely the reduction. The voltage between two half-cells is calculated as:

E{cell}^0 = E{reduction}^0 - E_{oxidation}^0

Example: Zn^{2+}/Zn and Cu^{2+}/Cu

E{cell}^0 = E{Cu^{2+}/Cu}^0 - E_{Zn^{2+}/Zn}^0 = +0.34V - (-0.76V) = +1.10V

Non-spontaneous Reactions-Electrolytic Cells

Electrolytic cells use electricity to force a non-spontaneous chemical reaction. In electrolysis, an electrical generator forces electrons to flow from anode to cathode.

Applications of Electrolysis:

• Obtaining active elements (e.g., sodium and chlorine from molten salts).
  2NaCl(l) \rightarrow 2Na(s) + Cl_2(g)

• Electroplating: The material to be plated is the cathode, and the anode is the plating metal. The electrolyte contains ions of the plating metal.

Key Points:

• Positive ions migrate away from the anode (anode is positive).

• Positive ions move toward the cathode (cathode is negative).

• An external power source forces electrons from anode to cathode.

Differences Between Voltaic and Electrolytic Cells:

Similarities Between Voltaic and Electrolytic Cells:

• Both use redox reactions.

• The anode is the site of oxidation.

• The cathode is the site of reduction.

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