Comprehensive Study Notes: Grade 10 Chemistry - Second Semester

Unit 5: Halogens and Trends in Group Properties

• Trends in Periodic Table Groups:

  • Elements are classified into groups based on physical and chemical characteristics.

  • Group VIII (Noble Gases): All elements are colorless gases and generally chemically inactive.

  • Group I (Alkali Metals): Soft metals that react with water to form alkaline solutions.

  • General Rule: There is a gradual change in characteristics as one moves down a group.

• Trends in Group VIII (Noble Gases):

  • Boiling Points: Increase as you move down the group (Heillyum to Krypton).

    • Helium (He): $-269\,^{\circ}\text{C}$.

    • Neon (Ne): $-246\,^{\circ}\text{C}$.

    • Argon (Ar): $-186\,^{\circ}\text{C}$.

    • Krypton (Kr): $-153\,^{\circ}\text{C}$.

  • Predictions: Scientists can predict values for lower elements. For Xenon (Xe), the predicted boiling point based on an average increase of $39\,^{\circ}\text{C}$ was $-114\,^{\circ}\text{C}$. The actual value is $-108\,^{\circ}\text{C}$.

  • Density: Increases down the group. For Xenon, the predicted density was $0.0045\,g/mL$, while the actual value is $0.0054\,g/mL$.

• Trends in Group I (Alkali Metals):

  • Melting Points: Decrease as you move down the group.

    • Lithium (Li): $181\,^{\circ}\text{C}$.

    • Sodium (Na): $98\,^{\circ}\text{C}$.

    • Potassium (K): $63\,^{\circ}\text{C}$.

    • Rubidium (Rb): $39\,^{\circ}\text{C}$.

    • Caesium (Cs) Prediction: Predicted between $20\,^{\circ}\text{C}$ and $30\,^{\circ}\text{C}$. Actual value is $29\,^{\circ}\text{C}$.

• Group VII: The Halogens:

  • The name "Halogen" means "Salt-former" as they react with metals to produce ionic compounds (salts).

  • Members of the group: Fluorine (F), Chlorine (Cl), Bromine (Br), Iodine (I), Astatine (At), and Tennessine (Ts).

  • Fluorine and Astatine are rarely used in school labs: Fluorine is highly reactive, and Astatine is radioactive.

  • Tennessine (Ts) was discovered recently; it is highly radioactive and decays in milliseconds.

• Common Characteristics of Halogens:

  • Seven electrons in the outer energy level.

  • Non-metallic and toxic elements with pungent odors.

  • Exist as diatomic molecules (e.g., $Cl_{2}$, $Br_{2}$, $I_{2}$).

  • Valency of 1; they form negative ions with a single charge ($Cl^{-}$, $Br^{-}$, $I^{-}$).

  • React with metals to form metallic halides and with non-metals to form covalent halides (e.g., $HCl$, $HBr$, $HI$).

• Physical Properties of Halogens at Room Temperature:

  • Chlorine ($Cl_{2}$): Pale green dense gas; toxic with a pungent smell.

  • Bromine ($Br_{2}$): Dark red liquid with a reddish-brown vapor; toxic and pungent.

  • Iodine ($I_{2}$): Gray solid with a purple vapor; toxic and pungent.

  • Trend Down the Group: State changes from gas to liquid to solid. Color becomes darker (pale green to red to gray/black).

• Reactivity and Displacement Reactions:

  • The most reactive halogen is at the top (Fluorine).

  • Reactivity decreases down the group (opposite of Group I).

  • Displacement: A more reactive halogen will displace a less reactive halogen from its halide solution.

  • Examples:

    • $Cl_{2}\,(aq) + 2KBr\,(aq) \rightarrow 2KCl\,(aq) + Br_{2}\,(aq)$ (Solution turns orange/darker).

    • $Br_{2}\,(aq) + 2KI\,(aq) \rightarrow 2KBr\,(aq) + I_{2}\,(aq)$ (Solution turns brown/darker).

Unit 6: Electrochemistry

• Definitions of Redox (Reduction-Oxidation):

  • Oxidation: Addition of oxygen, loss of electrons ($e^{-}$), or loss of hydrogen.

  • Reduction: Loss of oxygen, gain of electrons ($e^{-}$), or gain of hydrogen.

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

  • Oxidising Agent: A substance that gives oxygen to another or gains electrons.

  • Reducing Agent: A substance that removes oxygen from another or loses electrons.

• Electron Transfer in Redox:

  • Example: Reaction between Zinc and Copper(II) sulfate.

  • Ionic Equation: $Zn\,(s) + Cu^{2+}+\,(aq) \rightarrow Zn^{2+}+\,(aq) + Cu\,(s)$.

  • Oxidation Half-equation: $Zn\,(s) \rightarrow Zn^{2+}+\,(aq) + 2e^{-}$.

  • Reduction Half-equation: $Cu^{2+}+\,(aq) + 2e^{-} \rightarrow Cu\,(s)$.

• Electrical Conductivity:

  • Conductors: Metals and graphite containing free electrons. No chemical change occurs during conduction.

  • Insulators: Non-conductors like molecular covalent compounds.

  • Electrolytes: Liquid compounds (molten or aqueous) that conduct electricity via the movement of ions. Chemical change (electrolysis) occurs during conduction.

• Electrolysis Fundamentals:

  • Electrolytic Cell: Device for electrolysis.

  • Anode (Positive Electrode): Where anions move and oxidation occurs.

  • Cathode (Negative Electrode): Where cations move and reduction occurs.

  • Inert Electrodes: Graphite or Platinum; they conduct electricity but do not react.

• Electrolysis of Molten Compounds:

  • Binary Ionic Compound: Compounds of two elements. Electrolysis decomposes them into a metal (at cathode) and a non-metal (at anode).

  • Example: Molten Lead(II) bromide ($PbBr_{2}$).

    • Cathode: $Pb^{2+}+\,(l) + 2e^{-} \rightarrow Pb\,(l)$ (Gray molten metal bead).

    • Anode: $2Br^{-}+\,(l) \rightarrow Br_{2}\,(g) + 2e^{-}$ (Brown vapor).

• Electrolysis of Aqueous Solutions:

  • Water dissociates slightly: $H_{2}O\,(l) \rightleftharpoons H^{+}+\,(aq) + OH^{-}+\,(aq)$.

  • Competition at Electrodes:

    • Cathode: If the metal is higher than hydrogen in the reactivity series (e.g., $Na$), $H_{2}\,(g)$ produced. If lower (e.g., $Cu$), the metal is produced.

    • Anode: Halide ions ($Cl^{-}$, $Br^{-}$, $I^{-}$) discharge if concentrated. Otherwise, $OH^{-}$ produces $O_{2}\,(g)$.

• Industrial Applications of Electrolysis:

  • Aluminum Production: Extracted from Bauxite ($Al_{2}O_{3}$). Dissolved in molten Cryolite to lower the melting point from $2030\,^{\circ}\text{C}$ to $900-1000\,^{\circ}\text{C}$.

    • Cathode: $Al^{3+}+\,(l) + 3e^{-} \rightarrow Al\,(l)$.

    • Anode: $2O^{2-}+\,(l) \rightarrow O_{2}\,(g) + 4e^{-}$.

  • Chlor-Alkali Industry: Electrolysis of concentrated $NaCl$ (brine).

    • Products: Chlorine gas ($Cl_{2}$), Hydrogen gas ($H_{2}$), and Sodium hydroxide solution ($NaOH$).

  • Copper Purification: Impure copper as anode, pure copper as cathode.

  • Electroplating: Coating a metal object with another metal (e.g., silver plating) for protection or decoration.

Unit 7: Applications of Organic Chemistry

• Alcohols: Ethanol:

  • Functional Group: Hydroxyl group ($OH^{-}$).

  • General Formula: $C_{n}H_{2n+1}OH$.

• Production of Ethanol:

  • Hydration of Ethene: $C_{2}H_{4}\,(g) + H_{2}O\,(g) \rightarrow C_{2}H_{5}OH\,(g)$.

    • Conditions: $300\,^{\circ}\text{C}$, $60\,atm$, Phosphoric acid catalyst ($H_{3}PO_{4}$).

    • Pros: Continuous process, fast, provides high purity ethanol.

  • Fermentation: $C_{6}H_{12}O_{6}\,(aq) \rightarrow 2C_{2}H_{5}OH\,(aq) + 2CO_{2}\,(g)$.

    • Conditions: Yeast enzymes, anaerobic (no oxygen), $36\,^{\circ}\text{C}$.

    • Pros: Uses renewable resources (sugar crops), low energy cost.

    • Cons: Batch process, slow, low concentration (yeast dies at $14\%$ ethanol).

• Uses of Ethanol:

  • Solvent: Used in glues, paints, perfumes, and cleaning products.

  • Fuel: Used in rocket engines and as biofuel in cars (common in Brazil).

• Polymers:

  • Macromolecules made of repeating units called monomers connected via polymerization.

  • Addition Polymerization: Occurs in alkenes where the $C=C$ bond breaks.

    • Polyethene: From ethene monomers. Used for plastic bags and bottles.

    • PVC (Polyvinyl Chloride): From chloroethene. Used for pipes.

    • PTFE (Teflon): From tetrafluoroethene. Non-stick surfaces.

  • Condensation Polymerization: Monomers join with the loss of a small molecule (usually water).

    • Nylon: Formed from a diamine and a dicarboxylic acid.

Unit 8: Chemical Energy and Equilibrium

• Exothermic and Endothermic Reactions:

  • Exothermic: Releases heat. Energy of products < Energy of reactants ($\Delta H$ is negative). Examples: Combustion, neutralization.

  • Endothermic: Absorbs heat. Energy of products > Energy of reactants ($\Delta H$ is positive). Example: Thermal decomposition.

  • Breaking bonds: Requires energy (Endothermic).

  • Forming bonds: Releases energy (Exothermic).

• Energy Cycle Calculations:

  • $\text{Total change in energy} = (\text{Sum of reactant bond energies}) - (\text{Sum of product bond energies})$.

• Activation Energy ($E_{a}$):

  • The minimum energy required for particles to collide and react.

  • Lowered by catalysts.

• Reversible Reactions and Equilibrium:

  • Reversible reactions use the $\rightleftharpoons$ symbol.

  • Example: Hydrated Copper(II) sulfate (blue) $\rightleftharpoons$ Anhydrous Copper(II) sulfate (white) + water.

  • Dynamic Equilibrium: Occurs in closed systems when the rate of the forward reaction equals the rate of the backward reaction.

• Industrial Processes and Conditions:

  • Haber Process (Ammonia Production): $N_{2}\,(g) + 3H_{2}\,(g) \rightleftharpoons 2NH_{3}\,(g)$.

    • Conditions: $450\,^{\circ}\text{C}$, $200\,atm$, Iron catalyst.

  • Contact Process (Sulfuric Acid Production): $2SO_{2}\,(g) + O_{2}\,(g) \rightleftharpoons 2SO_{3}\,(g)$.

    • Conditions: $450\,^{\circ}\text{C}$, $1-2\,atm$, Vanadium(V) oxide ($V_{2}O_{5}$) catalyst.

  • Product Yield Factors: High pressure favors the direction with fewer gas molecules. Lower temperature favors the exothermic direction.

• Fertilizers:

  • NPK Fertilizers provide three essential elements: Nitrogen (N), Phosphorus (P), and Potassium (K).

  • Examples: Ammonium nitrate ($NH_{4}NO_{3}$), Ammonium sulfate ($(NH_{4})_{2}SO_{4}$).

  • Environmental Impact: Eutrophication occurs when excess fertilizers wash into water bodies, causing algae blooms, oxygen depletion, and death of aquatic life.