Metals, Ores, and Extraction Processes Study Guide: Extraction, and Corrosion Study Notes

Composition of the Earth's Crust

• Metal Sources: Metals are primarily obtained from the Earth's crust, which is the planet's outer layer, though some are retrieved from the sea.

• State of Metals:

  • Uncombined/Native: Some metals occur in their elemental state because they are unreactive. Examples include gold, silver, platinum, copper, and elements like sulphur.

  • Combined: Most metals are found as compounds within the crust because they are reactive and tend to form compounds.

• Chemical Composition of the Crust: If the crust were broken down into its constituent elements, the breakdown would be:

  • Oxygen: Approximately $45\%$.

  • Silicon: Over $27\%$.

  • Aluminium: $8\%$.

  • Iron: $6\%$.

  • Calcium: $5\%$.

  • Magnesium: $3\%$.

  • Sodium: $2.5\%$.

  • Potassium: $1.5\%$.

  • All other metals and non-metals: $2\%$.

• Primary Components: Nearly three-quarters of the crust is composed of just two non-metals: oxygen and silicon. These are often found together in compounds like silicon dioxide ($\text{silica or sand}$). Oxygen also appears in compounds such as aluminium oxide, iron oxide, and calcium carbonate.

• Abundancy: Aluminium is the most abundant metal, followed by iron.

Scarcity and Plentiful Metals

• Plentiful Metals: Only six metals (Al, Fe, Ca, Mg, Na, K) are considered plentiful.

• Scarce Metals: A metal is defined as scarce if it makes up less than one-thousandth ($\frac{1}{1000}$) of the Earth's crust. Examples include:

  • Copper

  • Mercury

  • Zinc

  • Silver

  • Lead

  • Gold

  • Tin

  • Platinum

• Economic Impact: Scarce metals are expensive and are being consumed rapidly, leading to concerns about them running out.

Metal Ores and Mining Economics

• Definition of Ores: Rocks containing a sufficiently high concentration of a metal or metal compound to make extraction worthwhile.

  • Rock Salt: The main ore of sodium ($\text{mostly sodium chloride}$).

  • Bauxite: The main ore of aluminium ($\text{mostly aluminium oxide}$).

  • Gold: Often found as a free, almost pure element.

• Mining Decisions: Companies must determine if mining is economical by answering the following:

  1. How much ore is present?

  2. How much metal can be extracted from it?

  3. Are there special extraction problems?

  4. What are the costs ($\text{roads, buildings, equipment, extraction plant, transport, fuel, chemicals, and wages}$)?

  5. What is the current market selling price of the metal?

  6. Will the operation yield a profit?

• Fluctuating Factors: Economic viability changes yearly. Low-grade ore may become profitable if fuel costs drop or the metal's price rises.

• Social and Environmental Impact:

  • Concerns: Pollution, dust, pits, scars on the landscape, and the need for land restoration after a mine is exhausted.

  • Benefits: New job creation for the local community.

Principles of Metal Extraction

• Reactivity Correlation: The difficulty of extraction depends on the metal's reactivity.

• Native Metals: Unreactive metals (silver, gold, platinum, some copper) are obtained by simple physical separation from impurities ($\text{similar to removing stones from soil}$). No chemical reaction is required.

• Extraction via Reduction: Metals existing as oxides (or compounds convertible to oxides) are extracted by removing oxygen.

  • Reduction Definition 1: The removal of oxygen.

  • Reducing Agent: A substance capable of removing oxygen. Carbon and carbon monoxide ($CO$) are common reducing agents.

  • Iron Extraction Example: $Fe_2O_3 (s) + 3CO (g) \rightarrow 2Fe (l) + 3CO_2 (g)$. In this reaction, iron(III) oxide is reduced, and carbon monoxide is oxidized.

• Extraction via Electrolysis:

  • Definition: Splitting a compound using electricity. This is the most powerful extraction method.

  • Conditions: The compound must be molten or dissolved so ions are free to move.

  • Economic Constraint: Because electricity is expensive, it is reserved for reactive metals ($\text{e.g., aluminium}$) that are difficult to extract otherwise.

  • Aluminium Example: $2Al_2O_3 (l) \rightarrow 4Al (l) + 3O_2 (g)$.

• Electronic Definition of Redox:

  • Reduction: The gain of electrons.

  • Oxidation: The loss of electrons.

  • Mnemonic: OILRIG ($\text{Oxidation Is Loss, Reduction Is Gain}$).

  • Redox Reaction: A reaction where oxidation and reduction occur simultaneously.

The Reactivity Series and Extraction Methods

• Reactivity Order (Highest to Lowest):

  1. Potassium ($K$)

  2. Sodium ($Na$)

  3. Calcium ($Ca$)

  4. Magnesium ($Mg$)

  5. Aluminium ($Al$)

  6. Carbon ($C$ - for comparison)

  7. Zinc ($Zn$)

  8. Iron ($Fe$)

  9. Tin ($Sn$)

  10. Lead ($Pb$)

  11. Hydrogen ($H$ - for comparison)

  12. Copper ($Cu$)

  13. Silver ($Ag$)

  14. Gold ($Au$)

• Methodology Summary:

  • Electrolysis: Used for metals more reactive than carbon (Potassium, Sodium, Calcium, Magnesium, Aluminium). These ores are difficult to decompose.

  • Heating with Carbon/CO: Used for metals less reactive than carbon (Zinc, Iron, Lead).

  • Natural Occurrence: Metals at the bottom (Silver, Gold) occur naturally and require no chemical reaction.

Electrolysis in Detail

• Molten Lead Bromide ($PbBr_2$):

  • Electrolyte: The liquid being electrolyzed.

  • Electrodes: Graphite rods carrying current.

  • Process:

    • Negative Electrode (Cathode): Positive lead ions ($Pb^{2+}$) move here and gain electrons: $Pb^{2+} + 2e \rightarrow Pb$.

    • Positive Electrode (Anode): Negative bromide ions ($Br^-$) move here, lose electrons, and pair up: $Br^- - e \rightarrow Br$ then $Br + Br \rightarrow Br_2$.

• Molten Compounds Summary:

  • Sodium Chloride ($2NaCl$): Forms Sodium at the cathode ($Na^+ + e \rightarrow Na$) and Chlorine at the anode ($Cl^- - e \rightarrow Cl$).

  • Potassium Iodide ($2KI$): Forms Potassium at the cathode ($K^+ + e \rightarrow K$) and Iodine at the anode ($I^- - e \rightarrow I$).

  • Copper(II) Bromide ($CuBr_2$): Forms Copper at the cathode ($Cu^{2+} + 2e \rightarrow Cu$) and Bromine at the anode ($Br^- - e \rightarrow Br$).

• Electrolysis of Solutions (Aqueous):

  • Water provides extra ions: $H^+$ and $OH^-$.

  • Negative Electrode Competition: Hydrogen gas is released if the metal is more reactive than hydrogen.

  • Positive Electrode Competition: Halogen gas is released if the compound is a halide (chloride, bromide, iodide). If not, oxygen is released.

  • Sodium Chloride Solution: Contains $Na^+$, $Cl^-$, $H^+$, and $OH^-$. $H^+$ ions accept electrons more readily than $Na^+$, producing hydrogen gas. $Cl^-$ ions give up electrons more readily than $OH^-$, producing chlorine gas. $Na^+$ and $OH^-$ remain to form sodium hydroxide solution.

Extraction of Aluminium

• Source: Bauxite ore ($\text{Aluminium oxide + impurities like sand and iron oxide}$).

• Step-by-Step Process:

  1. Geological testing of bauxite content.

  2. Mining ($\text{usually surface mining}$).

  3. Purification at a bauxite plant to produce white alumina ($Al_2O_3$).

  4. Transport to electrolysis plants where electricity is cheap.

• Electrolysis of Alumina:

  • Melting Point: Pure alumina melts at $2045\,^{\circ}C$ ($\text{expensive/dangerous}$).

  • Solution: Dissolved in molten cryolite ($\text{sodium aluminium fluoride}$) at approximately $950\,^{\circ}C$.

  • Electrodes: Carbon-lined steel tank (negative) and huge carbon blocks (positive).

  • Reaction at Cathode: $Al^{3+} + 3e \rightarrow Al$ ($\text{reduction}$). Molten metal drops to the bottom and is run off.

  • Reaction at Anode: $2O^{2-} - 4e \rightarrow O_2$ ($\text{oxidation}$). Oxygen gas reacts with carbon electrodes to form $CO_2$ ($C (s) + O_2 (g) \rightarrow CO_2 (g)$), necessitating periodic electrode replacement.

Purification and Electroplating

• Copper Purification: Necessary for high electrical conductivity.

  • Setup: Positive electrode ($\text{impure copper}$) and negative electrode ($\text{pure copper}$) in copper(II) sulphate solution.

  • Process: Copper atoms from the impure electrode become ions and dissolve ($\text{reduction of the anode size}$), then move to the pure electrode to become copper atoms ($\text{increasing pure cathode size}$). Impurities fall to the bottom as sludge, often containing gold, silver, and platinum.

• Electroplating: Coating a metal object with a different metal via electrolysis.

  • Rules: Negative electrode = object to be plated; Positive electrode = plating metal (X); Solution = compound of metal X.

  • Uses: Chromium plating car bumpers for rust protection and aesthetics; Tin plating steel for food cans.

Properties and Uses of Metals and Alloys

• Aluminium: Low density, good conductor, malleable, non-toxic. Used in milk bottle tops, cooking foil, electricity cables, beer cans, aircraft, and space rockets.

• Copper: Excellent conductor, ductile. Used for electrical wiring.

• Lead: Soft, resists corrosion. Used to seal brickwork around chimneys.

• Alloys: Mixtures of a metal with other substances ($\text{metals or non-metals}$) to change properties.

  • Mild Steel: $99.5\% \text{ iron} + 0.5\% \text{ carbon}$. Hard and strong; used for bridges and car bodies.

  • Stainless Steel: $70\% \text{ iron}, 20\% \text{ chromium}, 10\% \text{ nickel}$. Rustproof; used for cutlery and kitchen sinks.

  • Cupronickel: $75\% \text{ copper}, 25\% \text{ nickel}$. Used for 'silver' coins.

  • Manganese Steel (Hadfield Steel): $85\% \text{ iron}, 13.8\% \text{ manganese}, 1.2\% \text{ carbon}$. Extremely hard; used for springs.

  • Titanium Alloy: $92.5\% \text{ titanium}, 5\% \text{ aluminium}, 2.5\% \text{ tin}$. High strength at high temperatures; used in jet engines.

  • Brass: $70\% \text{ copper}, 30\% \text{ zinc}$. Does not corrode; used for musical instruments.

  • Bronze: $95\% \text{ copper}, 5\% \text{ tin}$. Used for statues and church bells.

  • Solder: $70\% \text{ tin}, 30\% \text{ lead}$. Low melting point; used for joining wires.

Corrosion and Rusting

• Corrosion Definition: The attack of a metal by air, water, or other substances in the environment. Reactive metals corrode more easily.

• Rusting: Specifically the corrosion of iron and steel in the presence of both air and water. It is an oxidation process.

• Rusting Prevention:

  1. Paint: Good for bridges; lead or zinc-based paints (e.g., $Pb_3O_4$) are effective.

  2. Grease/Oil: Used for tools and machine parts.

  3. Plastic: Cheap and attractive for garden furniture/dish racks.

  4. Galvanizing: Coating with zinc.

  5. Tin Plating: Used for food cans because tin is unreactive and non-toxic.

  6. Chromium Plating: Shiny protective layer for car bumpers.

  7. Sacrificial Protection: Attaching a more reactive metal ($\text{magnesium or zinc}$) to iron. The more reactive metal corrodes instead of the iron.

• Aluminium Protection: Unlike iron, which forms flaky rust that allows further corrosion, aluminium forms a tough, sealing coat of aluminium oxide that prevents further damage. This can be thickened via electrolysis ($\text{anodized aluminium}$).

Questions & Discussion

• Q: Why is no chemical reaction needed to get gold?

  • A: Because gold is unreactive and exists naturally in its elemental state.

• Q: Lead is extracted by heating its oxide with carbon… What is reduced? What is the reducing agent?

  • A: Lead oxide is reduced. Carbon is the reducing agent and is itself oxidized.

• Q: Why is the reaction between lead oxide and carbon a redox reaction?

  • A: Because reduction (loss of oxygen by lead oxide) and oxidation (gain of oxygen by carbon) occur at the same time.

• Q: Sodium extraction from rock salt requires electrolysis. Why?

  • A: Sodium is a very reactive metal, making its compounds very stable and difficult to decompose through simple heating with carbon.

• Q: Which ions are oxidized and which are reduced in the electrolysis of sodium chloride?

  • A: Sodium ions ($Na^+$) are reduced to sodium atoms at the negative electrode. Chloride ions ($Cl^-$) are oxidized to chlorine atoms at the positive electrode.