Chapter 7: Electrochemistry and Redox Reactions

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Chapter Overview & SLOs

What is electrochemistry? Electrochemistry is the branch of chemistry that deals with the relationship between electrical energy and chemical changes, specifically redox reactions where electron transfer occurs.

What is the fundamental difference between oxidation and reduction?

• Oxidation is defined as:
  • Loss of electrons
  • Increase in oxidation state
  • Addition of oxygen
  • Removal of hydrogen
• Reduction is defined as:
  • Gain of electrons
  • Decrease in oxidation state
  • Addition of hydrogen
  • Removal of oxygen
• Mnemonic: OIL RIG (Oxidation Is Loss, Reduction Is Gain)

How do we identify oxidizing and reducing agents?

• Oxidizing agent (oxidant): The reactant that contains the element being reduced (gains electrons). It causes oxidation in another substance.
• Reducing agent (reductant): The reactant that contains the element being oxidized (loses electrons). It causes reduction in another substance.
• Example: In $H_2S + Cl_2 \rightarrow 2HCl + S$, $Cl_2$ is the oxidizing agent (reduced to $Cl^-$), and $H_2S$ is the reducing agent (oxidized to $S$).

How do we calculate oxidation states (oxidation numbers)? Rules for assigning oxidation numbers:

• The oxidation number of an element in its free (uncombined) state is 0 (e.g., $O_2$, $H_2$, $Na$).
• The oxidation number of a monatomic ion equals its charge (e.g., $Na^+ = +1$, $Cl^- = -1$).
• Fluorine always has oxidation number -1 in compounds.
• Oxygen usually has oxidation number -2 (except in peroxides like $H_2O_2$ where it's -1, and in $OF_2$ where it's +2).
• Hydrogen usually has oxidation number +1 (except in metal hydrides like $NaH$ where it's -1).
• The sum of oxidation numbers in a neutral compound is 0.
• The sum of oxidation numbers in a polyatomic ion equals the ion's charge.

Examples of oxidation state calculations:

• $H_2CrO_4$: H = +1 (×2 = +2), O = -2 (×4 = -8), let Cr = $x$. Equation: $+2 + x - 8 = 0$ → $x = +6$.
• $NO_2$: O = -2 (×2 = -4), let N = $x$. Equation: $x - 4 = 0$ → $x = +4$.
• $H_2S$: H = +1 (×2 = +2), let S = $x$. Equation: $+2 + x = 0$ → $x = -2$.
• $K_2Cr_2O_7$: K = +1 (×2 = +2), O = -2 (×7 = -14), let Cr = $x$ (×2 = $2x$). Equation: $+2 + 2x - 14 = 0$ → $2x = 12$ → $x = +6$.

What are the redox applications in daily life?

• Photography: The reduction of silver bromide ($Ag^+$) to black metallic silver ($Ag$) forms the image on photographic film.
• Batteries: Redox reactions produce electrical energy.
• Corrosion: Rusting of iron is an oxidation process.
• Hydrogen Peroxide ($H_2O_2$) versatility:
  • As a reducing agent (with silver oxide): $$Ag_2O + H_2O_2 \rightarrow 2Ag + H_2O + O_2$$
  • As an oxidizing agent (with lead(II) sulphide): $$PbS + 4H_2O_2 \rightarrow PbSO_4 + 4H_2O$$

What is corrosion and how can it be prevented? Corrosion is the deterioration of metals due to chemical reactions with their environment (mostly oxidation).

Corrosion prevention methods:

• Tin plating (tinning): Coating iron or steel with tin. Used in food industry to prevent organic acids from reacting with iron (e.g., tin cans).
• Galvanization: Coating iron with zinc (more reactive, acts as sacrificial anode).
• Passivation: Forming a protective oxide layer on the metal surface (e.g., aluminum oxide on aluminum, stainless steel).
• Protective coatings: Painting, oiling, greasing to prevent exposure to air and moisture.
• Corrosion inhibitors: Chemicals added to the environment to slow down corrosion.
• Cathodic protection: Connecting the metal to a more reactive sacrificial anode (e.g., magnesium or zinc).
• Alloying: Mixing metals to create corrosion-resistant alloys (e.g., stainless steel contains chromium).

Key redox terms summary:

| Process | Electron Change | Oxidation State | Example | |---------|----------------|----------------|---------| | Oxidation | Loss | Increases | $Fe \rightarrow Fe^{3+} + 3e^-$ | | Reduction | Gain | Decreases | $Ag^+ + e^- \rightarrow Ag$ |

Important note: The oxidizing agent is the substance that is reduced (gains electrons), and the reducing agent is the substance that is oxidized (loses electrons).

These notes are strictly aligned with the Student Learning Outcomes (SLOs) for the FBISE 2026 annual examination.

• How do we define and contrast redox processes? Distinguish between oxidation and reduction based on electron transfer (loss vs gain), changes in oxidation states (increase vs decrease), and the movement of oxygen or hydrogen atoms (addition/removal).
• How do we calculate oxidation numbers in complex compounds? Apply algebraic rules to solve for the oxidation state of central atoms in molecules such as $H_2CrO_4$ (Cr = +6), $NO_2$ (N = +4), $H_2S$ (S = -2), and $K_2Cr_2O_7$ (Cr = +6), ensuring the sum of all oxidation numbers in a neutral molecule equals zero.
• How do we identify oxidizing and reducing agents in chemical reactions? Analyze equations like $H_2S + Cl_2 \rightarrow 2HCl + S$ to determine which substance acts as the oxidizing agent (the one reduced, $Cl_2$) and which acts as the reducing agent (the one oxidized, $H_2S$) based on their changes in oxidation state.
• How is corrosion managed in engineering and industry? Evaluate different methods for preventing the deterioration of metals, including tin plating (tinning for food industry), galvanization, passivation, protective coatings (paint, oil), corrosion inhibitors, cathodic protection, and alloying to create corrosion-resistant materials.

Frequently Asked Questions (FAQ)