Chemical Changes and Reactivity Series

Reactivity Series

• The reactivity series lists metals from most reactive (top) to least reactive (bottom).
• Non-metals (often in red) are used for displacement reactions.

Displacement Reactions

• A more reactive element displaces a less reactive element from its compound.
• Example: Lithium displacing copper in copper chloride.
  • Lithium is more reactive than copper.
  • $Li + CuCl_2 \rightarrow LiCl + Cu$
• Example: Aluminium displacing zinc in zinc sulfate.
  • Aluminium is more reactive than zinc.
  • $2Al + 3ZnSO<em>4 \rightarrow Al</em>2(SO<em>4)</em>3 + 3Zn$
• These reactions are called displacement reactions because the more reactive element displaces the less reactive element.

Carbon for Metal Purification

• Carbon is used to purify metals.
• Carbon is more reactive than zinc, iron, and copper.
• Carbon removes oxygen from metal oxides.
  • $2Fe<em>2O</em>3 + 3C \rightarrow 4Fe + 3CO_2$
  • Carbon takes the oxide to form carbon dioxide, leaving pure iron.

Metal Extraction from Ores

• Metals are extracted from ores obtained through mining in quarries.
• Metal ores are metal oxides (rock + oxygen).
• Iron ore is an example of a metal oxide.
• To obtain the metal, oxygen must be removed from the ore.
• Heating with Carbon in a Blast Furnace:
  • Metal ore is heated to a very high temperature with carbon to remove the oxygen.

Environmental Issues of Mining

• Environmental Concerns:
  • Mining causes habitat and landscape destruction.
  • Heavy machinery leads to noise and dust pollution.
  • Process produces carbon dioxide ($CO_2$).
• Positives:
  • Metals (iron, copper, etc.) are essential for building materials.

Reduction Reactions

• The process of removing oxygen is called reduction.
  • Iron oxide is reduced to iron when it loses oxygen.

Oxidation and Reduction

• Reduction: Loss of oxygen.
  • $Fe<em>2O</em>3 \rightarrow Fe$ (Iron oxide loses oxygen to become iron).
  • $CuO \rightarrow Cu$ (Copper oxide loses oxygen to become copper).
• Oxidation: Gain of oxygen.
  • $O<em>2 + Fe \rightarrow Fe</em>2O_3$ (Oxygen reacts with iron to form iron oxide).
  • $O_2 + Cu \rightarrow CuO$ (Oxygen reacts with copper to form copper oxide).

pH Scale

• Acids:
  • Strong acids: pH 0-2 (e.g., hydrochloric acid (HCl), nitric acid ($HNO<em>3$), sulfuric acid ($H</em>2SO_4$)).
  • Weak acids: pH 3-6 (e.g., lemon juice, vinegar).
• Neutral:
  • pH 7 is neutral (e.g., water).
• Alkalis/Bases:
  • Weak alkalis: pH 9-11 (e.g., baking soda).
  • Strong alkalis: pH 12-14 (e.g., sodium hydroxide (NaOH)).
• Acids and alkalis are naturally colorless; universal indicator changes color in their presence.

Ions in Acids and Alkalis

• Acids contain hydrogen ions ($H^+$).
• Alkalis contain hydroxide ions ($OH^-$).

Making Salts

• General Formulas:
  • Acid + Metal → Salt + Hydrogen
    • Example: Hydrochloric acid + Magnesium → Magnesium chloride + Hydrogen
      • $2HCl + Mg \rightarrow MgCl<em>2 + H</em>2$
    • Example: Sulfuric acid + Calcium → Calcium sulfate + Hydrogen
      • $H<em>2SO</em>4 + Ca \rightarrow CaSO<em>4 + H</em>2$
    • Example: Nitric acid + Iron → Iron nitrate + Hydrogen
      • $2HNO<em>3 + Fe \rightarrow Fe(NO</em>3)<em>2 + H</em>2$
  • Acid + Base → Salt + Water
    • Example: Hydrochloric acid + Copper oxide → Copper chloride + Water
      • $2HCl + CuO \rightarrow CuCl<em>2 + H</em>2O$
    • Example: Sulfuric acid + Lithium hydroxide → Lithium sulfate + Water
      • $H<em>2SO</em>4 + 2LiOH \rightarrow Li<em>2SO</em>4 + 2H_2O$
  • Acid + Metal Carbonate → Salt + Water + Carbon Dioxide
    • Example: Nitric acid + Calcium carbonate → Calcium nitrate + Water + Carbon dioxide
      • $2HNO<em>3 + CaCO</em>3 \rightarrow Ca(NO<em>3)</em>2 + H<em>2O + CO</em>2$

Neutralization

• Acid + Alkali → Salt + Water (pH 7)
• Neutralization is the canceling out of $H^+$ and $OH^-$ ions to form water ($H_2O$).
  • Example: $HCl + NaOH \rightarrow NaCl + H_2O$

Practical: Forming Crystals

1. Add copper oxide to sulfuric acid.
2. Stir the mixture to ensure reaction and heat to speed up the reaction.
3. Resulting solution contains copper sulfate, water, and excess copper oxide.
4. Filter the solution to remove excess copper oxide.
5. Leave the filtered solution to slowly evaporate, forming copper sulfate crystals (crystallization).
6. Dab crystals dry with a paper towel.
7. Always wear safety glasses when handling acids.

Practical: Electrolysis

• Electrolysis is the process of using electricity to decompose a compound.
• Electrodes:
  • Conduct electricity and are inert (do not react).
• Ions are free to move when an ionic lattice is melted or dissolved in water.
• In the solution, ions (positive and negative) move to carry the charge.
• Metals form positive ions, and nonmetals form negative ions.
• During electrolysis:
  • Positive ions move to the negative electrode (cathode) and are reduced (gain electrons).
  • Negative ions move to the positive electrode (anode) and are oxidized (lose electrons).
  • Example: Electrolysis of Lead Iodide ($PbI_2$)
    • Lead (Pb) is formed at the negative electrode.
    • Iodine ($I_2$) is formed at the positive electrode.
  • Example: Electrolysis of Potassium Bromide (KBr)
    • Potassium (K) is formed at the negative electrode.
    • Bromine ($Br_2$) is formed at the positive electrode.
• Halogens (Group 7):
  • Iodide ($I^-$) becomes iodine ($I_2$).
  • Bromide ($Br^-$) becomes bromine ($Br_2$).
  • Chloride ($Cl^-$) becomes chlorine ($Cl_2$).
• Oxidation and Reduction in Electrolysis:
  • Chloride ion loses an electron: $2Cl^- \rightarrow Cl_2 + 2e^-$.
  • Potassium ion gains an electron: $K^+ + e^- \rightarrow K$.