Key Concepts in Chemical Changes

Topic 4: Chemical Changes

Acids and Bases

  • pH Scale

    • The pH scale is a measure of how acidic or basic a solution is.

    • The lower the pH of a solution, the more acidic it is.

    • The higher the pH of a solution, the more basic it is.

    • A neutral substance has a pH of 7.

  • How to Measure pH:

    • Indicators:

    • Indicators are dyes that change color depending on whether the solution is above or below a certain pH level.

    • Some indicators contain a mixture of dyes that gradually change color over a broad range of pH levels.

    • pH Probe:

    • A pH probe can be attached to a pH meter to read pH electronically.

Acids and Bases Neutralisation

  • Definition: When an acid reacts with a base, they neutralize each other.

  • Definitions of Key Terms:

    • Acid: A substance that has a pH of less than 7 and forms H⁺ ions in an aqueous solution.

    • Base: A substance that has a pH greater than 7.

    • Alkali: A base that dissolves in water to produce a solution with a pH greater than 7.

  • Neutralisation Reaction:

    • When an acid neutralizes a base, the products are salt and water:

    • \text{Acid} + \text{Base} \rightarrow \text{Salt} + \text{Water}

Strong Acids and Weak Acids

  • Acids can be classified as strong acids or weak acids based on their ionization in water:

    1. Strong Acids: Ionize completely in water and release a higher concentration of H⁺ ions.

    2. Weak Acids: Do not fully ionize in aqueous solution; only a portion of the acid molecules dissociate to release H⁺ ions.

  • Dynamics of Acid Reactions:

    • The reaction rate is faster for strong acids than for weak acids of the same initial concentration due to the higher concentration of H⁺ ions.

  • pH Calculation:

    • For every decrease of 1 in pH, the concentration of H⁺ ions increases by a factor of 10.

    • Thus, a strong acid will have a significantly lower pH than a weak acid if both have the same concentration.

  • Clarification of Terms:

    • Acid Strength: Refers to the proportion of acid that ionizes in water.

    • Concentration: Refers to how much acid is present in a certain volume of water.

Reactions of Acids

  • Acids react with metals, metal oxides, and metal hydroxides, typically to form salts and other products:

    • General Reactions:

    1. \text{Acid} + \text{Metal Hydroxide} \rightarrow \text{Salt} + \text{Water}

    2. \text{Acid} + \text{Metal} \rightarrow \text{Salt} + \text{Hydrogen gas}

    3. \text{Acid} + \text{Metal Carbonate} \rightarrow \text{Salt} + \text{Water} + \text{Carbon Dioxide}

  • Making Soluble Salts from Insoluble Bases:

    1. Measure 20 cm³ of acid in a measuring cylinder.

    2. Place the measuring cylinder over a tripod and gauze to heat.

    3. Add half a spatula of an insoluble base and mix with a glass rod.

    4. Keep adding until no more base dissolves, and it sinks to the bottom of the vessel.

    5. Let the solution cool and filter to obtain the salt solution.

    6. Pour the filtrate back into an evaporating dish to reduce the volume by half.

    7. Allow it to cool to form crystals.

Reactivity Series

  • The reactivity series lists metals in order of their reactivity based on how easily they lose electrons:

    • Higher reactivity correlates with faster reactions, particularly with acids, producing hydrogen gas:

    • \text{Acid} + \text{Metal} \rightarrow \text{Salt} + \text{Hydrogen}

  • Examples of Metals in the Reactivity Series:

    • Very reactive: Potassium (K), Sodium (Na), Lithium (Li)

    • Reactive: Calcium (Ca), Magnesium (Mg)

    • Fairly reactive: Zinc (Zn), Iron (Fe)

    • Not very reactive: Copper (Cu)

  • Metal Reaction with Water:

    • More reactive metals will react with water; less reactive metals typically do not.

Metal Ores and Extraction

  • Extraction of metals from their ores often involves oxidation:

    • Definitions:

    • Ore: Naturally occurring substances that contain metals in a form that can be extracted.

    • Oxidation: The process of a substance reacting with oxygen.

    • Reduction Reaction:

    • A chemical reaction that separates metals from their ores usually involves reduction, particularly using carbon:

    • Only metals that are less reactive than carbon can be extracted by this method; more reactive metals require electrolysis.

    • Examples of Ores:

    • Aluminum ore: Bauxite (Al₂O₃)

    • Copper ore: Cuprite (Cu₂O)

    • Iron ore: Hematite (Fe₂O₃)

Displacement Reactions

  • In a displacement reaction, a more reactive metal displaces a less reactive metal from its compound:

    • Example Reaction:

    • \text{Fe} + \text{CuSO}4 \rightarrow \text{FeSO}4 + \text{Cu}

Introduction to Electrolysis

  • Electrolysis: The process of splitting ionic compounds into their elements using electricity.

  • Important Considerations:

    • Electrolysis cannot occur if ions are in fixed positions (as in solids); however, molten or dissolved ionic compounds allow ions to move freely.

  • Equipment for Electrolysis:

    • Electrodes:

    • Cathode (negative electrode) - Attracts positive ions.

    • Anode (positive electrode) - Attracts negative ions.

    • Chemical Changes at Electrodes:

    • At the Anode: Non-metal anions lose electrons (oxidation) to form atoms.

    • At the Cathode: Cations gain electrons (reduction) to form metal atoms.

Electrolysis of Molten Ionic Compounds

  • Electrolysis is used to extract metals that are more reactive than carbon because they cannot be reduced by carbon:

    • This process is expensive due to the energy costs involved and the current needed.

    • For example, aluminum is extracted from bauxite by:

    1. Mixing bauxite with cryolite to lower the melting point.

    2. The molten mixture allows for the movement of free ions.

    3. Positive Al³⁺ ions are attracted to the cathode where they gain electrons to form aluminum atoms, which sink to the bottom.

    4. Negative O²⁻ ions are attracted to the anode where they lose electrons, forming oxygen.

    • The overall processes are described as:

    • Reduction: Positive ions gain electrons to form neutral elements.

    • Oxidation: Negative ions lose electrons to become neutral elements.

Electrolysis of Aqueous Solutions

  • Dissolving ionic compounds in water produces H⁺ and OH⁻ ions along with the ions from the compound:

    • At the Cathode:

    • Two positive ions are present: one from water and one from the metal ion.

    • The least reactive will gain electrons to reduce to its elemental form.

    • At the Anode:

    • There will be two negative ions: hydroxide ions and halide ions (if present).

    • If halides are present, they will form non-metals (such as chlorine, bromine, or iodine). If no halides are present, OH⁻ ions will form oxygen gas.

    • Example with Copper Sulfate (CuSO₄):

    • Cathode Reactions:

    • \text{Cu}^{2+} + 2e^- \rightarrow \text{Cu}

    • Anode Reactions:

    • 4\text{OH}^- \rightarrow \text{O}2 + 2\text{H}2\text{O} + 4e^-

Half Equations and Ionic Equations

  • Half Equations: Show the individual reactions that occur at the electrodes during electrolysis, specifically the gain or loss of electrons:

    • For example:

    • \text{Mg}^{2+} + 2e^- \rightarrow \text{Mg} (Reduction)

    • \text{Br}_2 + 2e^- \rightarrow 2\text{Br}^- (Reduction)

  • These equations represent the specific electron transfers involved in redox reactions during