C4.2_Knowledge_Organiser

Page 1: Prior Knowledge Review: Reactions of Metals

  • Reactivity Series

    • Metals are arranged by their reactivity, from most reactive to least reactive.

    • Key metals include potassium, sodium, lithium, calcium, magnesium, zinc, iron, and copper.

    • Unreactive metals do not easily participate in chemical reactions.

    • Reactive metals readily engage in chemical reactions.

  • Reactions with Oxygen and Acids

    • Metals react with oxygen to form metal oxides (oxidation reactions).

    • Reactions between metals and acids produce salts and hydrogen gas:

      • Hydrochloric acid (HCl)

      • Nitric acid (HNO₃)

      • Sulphuric acid (H₂SO₄)

    • Acids can neutralize by reacting with alkalis and bases to yield salts and water.

  • Properties of Acids and Alkalis

    • Alkalis release hydroxide ions (OH⁻) in solutions.

    • Alkalis and bases can be metal oxides or metal hydroxides.

  • Extraction of Metals

    • Unreactive metals like gold can be found in nature as the metal itself.

    • Most metals exist in compounds requiring chemical reactions for extraction.

    • Metals less reactive than carbon can be extracted by reduction using carbon (loss of oxygen).

  • Reactivity Series Considerations

    • Non-metals hydrogen and carbon are included in the reactivity series.

    • A more reactive metal can displace a less reactive metal from its compound.

  • Ions and Ionic Bonding

    • Electron transfer occurs during the reaction of metals with non-metals.

    • Metals become positively charged ions by losing electrons.

    • Non-metals become negatively charged ions by gaining electrons.

    • Ions produced have electronic structures similar to noble gases (Group 0).

    • Charge on ions corresponds to group number in the periodic table.

Page 2: Ionic Equations and Displacement Reactions

  • Oxidation and Reduction

    • Oxidation involves loss of electrons; reduction involves gain of electrons.

    • Metal atoms react to form positive ions, and their reactivity correlates to the formation of positive ions.

  • Ionic Equations

    • Balanced equations reflect ions involved in reactions.

    • Spectator ions (same on both sides of the equation) can be omitted from balanced equations.

    • Ionic equations can have half-equations representing oxidation and reduction.

  • Reactions of Acids with Metals

    • These reactions are classified as redox reactions.

    • Example: Magnesium reacts with hydrochloric acid to form magnesium chloride and hydrogen.

  • Electrolysis

    • Ionic compounds in liquid or aqueous states conduct electricity (electrolytes).

    • Ions move towards electrodes when electricity is passed.

    • Positive ions migrate to the cathode, gaining electrons (reduction).

    • Negative ions move to the anode, losing electrons (oxidation).

Page 3: Extracting Metals by Electrolysis

  • Electrolysis for Metal Extraction

    • Employed for metals too reactive for extraction via carbon reduction.

    • High energy consumption is required for melting compounds and generating electrical currents.

  • Electrolysis Process

    • aluminium extraction involves electrolysis of a molten mixture of aluminium oxide and cryolite; carbon is used as the anode.

    • The ore is dissolved in cryolite to reduce energy costs of melting.

  • Discharge at Electrodes

    • Electrolysis of molten ionic compounds leads to individual element production at electrodes.

    • Example: Lead bromide leads to lead at the cathode and bromine at the anode during electrolysis.

  • Electrolysis in Solutions

    • Ion discharge at electrodes varies based on reactivity; hydrogen is produced at the cathode if metal is more reactive than hydrogen.

    • At the anode, oxygen is produced unless halide ions are present, in which case halogens are produced.

Page 4: Corrosion and its Prevention

  • Corrosion

    • Chemical reactions with environmental substances lead to material degradation, a process exemplified by rusting in iron.

    • Rusting requires both oxygen and water for iron.

  • Corrosion Prevention

    • Protective coatings (greasing, painting, electroplating) serve as barriers to prevent corrosion.

    • Aluminium’s oxide coating prevents further corrosion.

    • Sacrificial protection uses a more reactive metal, e.g., zinc to galvanize iron.

  • Raw Materials and Environmental Impact

    • Limited Earth resources necessitate alternative extraction methods (phytomining, bioleaching) for metals like copper.

    • Traditional mining harms wildlife habitats.

  • Phytomining and Bioleaching

    • Phytomining uses plants to extract metal compounds that are later processed from plant ash.

    • Bioleaching employs bacteria to create leachate containing metal compounds.

    • Example: Copper can be extracted from leachate via displacement reactions using scrap iron or through electrolysis.

Recycling Materials

  • Recycling Process

    • Metals are recycled through melting and recasting to create new products.

    • The recycling process requires varying levels of material separation depending on the type and properties required for the final product, such as using scrap steel to reduce the need for new iron extraction from ore.