MYP Chemistry 5.13 Corrosion Notes

What Is Corrosion, and How Is It Different from Simple Oxidation?

Simple Oxidation

Corrosion

Corrosion:

1. Usually occurs at the surface but can penetrate deeper.
2. Leads to loss of strength, holes, and structural damage.
3. Involves a series of redox reactions.

Corrosion as a Redox Process

1. Redox recap:
   1. Oxidation = loss of electrons.
   2. Reduction = gain of electrons.
2. In corrosion:
   1. The metal (e.g. iron) is oxidised (loses electrons).
   2. Another substance (often oxygen or H⁺ in acidic solutions) is reduced (gains electrons).
3. Simplified steps of rusting (iron + water + oxygen):
   1. Oxidation of iron (anodic reaction): $$\mathrm{Fe}(\mathrm{~s}) \rightarrow \mathrm{Fe}^{2+}(\mathrm{aq})+2 e^{-}$$ Iron atoms lose electrons and form Fe²⁺ ions.
   2. Reduction of oxygen (cathodic reaction): $$\mathrm{O}_2(\mathrm{~g})+2 \mathrm{H}_2 \mathrm{O}(\mathrm{l})+4 e^{-} \rightarrow 4 \mathrm{OH}^{-}(\mathrm{aq})$$ Oxygen gains electrons and forms hydroxide ions.
   3. Formation of rust:
      1. Fe²⁺ reacts with OH⁻: $$\mathrm{Fe}^{2+}+2 \mathrm{OH}^{-} \rightarrow \mathrm{Fe}(\mathrm{OH})_2(\mathrm{~s})$$
      2. Further oxidation and hydration produce hydrated iron(III) oxide, the familiar rust.

What Conditions Speed Up Corrosion, and How Can We Investigate Them?

Corrosion happens fastest when the environment helps electrons move and accelerates redox reactions.

Factors That Increase Corrosion

1. Presence of water and oxygen
   1. Both are needed for rusting of iron.
   2. No water → no rust.
   3. No oxygen → no rust.
2. Electrolytes (e.g. salt)
   1. Salt (NaCl) dissolved in water makes the water a better conductor.
   2. This speeds up the electron and ion movement → faster corrosion.
   3. Explains why cars and bridges near the sea corrode faster.
3. Acidic or polluted environments
   1. Acid rain (containing H₂SO₄ or HNO₃) increases the concentration of H⁺ ions.
   2. Hydrogen ions can take part in redox reactions, speeding up corrosion.
   3. Industrial areas with acidic gases (SO₂, NOₓ) show more corrosion.
4. Higher temperature
   1. Increases the rate of chemical reactions.
   2. Metals often corrode faster in warm, humid climates.

Simple Laboratory Investigations

You can design experiments to isolate each factor.

Effect of Water and Oxygen

1. Tube A: Iron nail + distilled water + air
2. Tube B: Iron nail + boiled water (no dissolved oxygen) + layer of oil on top (keeps oxygen out)
3. Tube C: Iron nail + dry air (e.g. with drying agent) but no water
4. Observation:
   1. Rust forms in Tube A only.
   2. Tube B (no oxygen) and Tube C (no water) show no rust.

Conclusion: Both water and oxygen are required for rusting.

Effect of Salt (Electrolytes)

1. Tube D: Iron nail + distilled water
2. Tube E: Iron nail + salt solution (NaCl(aq))
3. Observation:
   1. Rust appears faster and more extensively in Tube E.

Conclusion: Salt (electrolytes) increase corrosion rate by improving conductivity.

Effect of Acids

1. Tube F: Iron nail + neutral water
2. Tube G: Iron nail + dilute acid (e.g. dilute H₂SO₄)
3. Observation:
   1. Rust and surface damage occur more quickly in Tube G.

Conclusion: Acidic conditions accelerate corrosion.

How Can We Prevent or Reduce Corrosion, and Why Does It Matter?

Because corrosion weakens metals and damages structures, we use several strategies to slow or prevent it.

Methods of Corrosion Prevention

Barrier Protection

1. Goal: Keep water and oxygen away from the metal surface.
2. Methods:
   1. Painting – creates a continuous barrier (e.g. bridges, ships, cars).
   2. Oil or grease – protects moving parts (e.g. tools, machinery).
   3. Plastic coatings – cover the metal completely (e.g. plastic-coated wire, fences).

Sacrificial Protection

1. Use a more reactive metal to protect a less reactive one.
2. The more reactive metal corrodes first, “sacrificing” itself.
3. The protected metal is forced to stay as the cathode in the redox process.

Alloying

Mixing a metal with other elements to make it more corrosion-resistant.

Chemical Treatments

1. Chemicals can create or enhance a passive (protective) layer on a metal.
2. Passivation – treating a metal with chemicals (e.g. nitric acid for stainless steel) to form a protective oxide layer.
3. Anti-rust agents – coolants and fluids in cars often contain inhibitors that slow down corrosion in radiators and engines.

Why Is It Important to Prevent Corrosion?

Corrosion has economic, environmental and safety impacts.

1. Economic Reasons
   1. Replacement costs:
      1. Corroded pipelines, bridges, vehicles and buildings often need expensive repair or replacement.
   2. Maintenance costs:
      1. Regular painting, coating, inspections and sacrificial anode replacement cost money and time.
   3. Globally, corrosion costs are estimated to be a large fraction of a country’s GDP in maintenance and replacements.
2. Environmental Reasons
   1. Resource use:
      1. Corroded materials must be replaced → more mining, refining and manufacturing, which use energy and raw materials.
   2. Pollution from leaks:
      1. Corroded oil or chemical pipelines can leak, contaminating soil, rivers, and oceans.
      2. Corroded storage tanks may release harmful substances into the environment.
3. Safety Reasons
   1. Corrosion can weaken:
      1. Bridges and buildings → risk of collapse.
      2. Aircraft and ships → structural failure.
      3. Gas and oil pipelines → leaks, fires or explosions.
   2. These failures can cause:
      1. Injury or loss of life.
      2. Damage to wildlife and ecosystems.