Chemistry Notes

Metal Oxides and Hydroxides

• Many metal oxides and hydroxides behave as bases and neutralize acids.
• Bases are substances that accept hydrogen ions (H^+, also known as protons, hence they're called proton acceptors).

Metal Oxides

• The oxide ion in metal oxides accepts protons and neutralizes acids.
• Example: Calcium oxide (lime) neutralizes sulfuric acid to form calcium sulfate:
  CaO(s) + H2SO4(aq) \rightarrow CaSO4(s) + H2O(l)
  O^{2-} + 2H^+ \rightarrow H_2O
• Lime is used in agriculture to neutralize acidic soil, preventing plants from absorbing toxic ions that affect root growth.

Amphoteric Oxides

• Amphoteric oxides behave as either acids or bases.
• Elements forming these oxides are typically in the middle of a period (e.g., aluminum).
• Aluminum oxide doesn't dissolve in water but reacts with acids to form salt and water:
  Al2O3(s) + 6HCl(aq) \rightarrow 2AlCl3(aq) + 3H2O(l)
• Aluminum oxide also reacts with bases to form aluminates; example:
  Al2O3 + 2NaOH + 3H2O \rightarrow 2Na[Al(OH)4]

Metal Hydroxides

• Metal oxides react with water to form hydroxides:
  MgO(s) + H2O(l) \rightarrow Mg(OH)2(aq)
• Hydroxide ions accept protons and neutralize acids.
• Example: Magnesium hydroxide is an antacid that neutralizes hydrochloric acid in the esophagus:
  Mg(OH)2(s) + 2HCl(aq) \rightarrow MgCl2(aq) + 2H2O(l) OH^-(aq) + H^+(aq) \rightarrow H2O(l)
• Calcium hydroxide neutralizes acidic effluent, such as dilute sulfuric acid:
  Ca(OH)2(aq) + H2SO4(aq) \rightarrow CaSO4(s) + 2H_2O(l)

Alumina

• Alumina is a form of aluminum oxide found in bauxite.
• It has a high melting point due to strong ionic bonds.
• It does not conduct electricity unless molten.
• Its thermal conductivity is about 30 times less than aluminum metal.
• Extracted and purified using the Bayer process:
  • Crush the bauxite.
  • React with NaOH(aq) at 170 °C to form sodium tetrahydroxoaluminate.
  • Filter out solid impurities.
  • Allow to crystallize to form Al(OH)_3.
  • Heat in a rotary kiln to form Al2O3.
• Most alumina is used in the Hall-Héroult process to produce aluminum by electrolysis.
• Some alumina is used as a refractory material in kilns (retains strength and chemical stability at high temperatures).

Aluminium and Titanium

• Both aluminium and titanium are low-density, corrosion-resistant metals that form strong alloys.

Extraction of Aluminium

• Aluminium ore (bauxite) is processed to form alumina (Al2O3).
• Molten alumina is electrolyzed using the Hall-Héroult process.
• Cryolite is added to lower the melting point and save energy.
• The steel tank lining acts as the negative electrode (cathode).
• Aluminium ions are reduced to form molten aluminium:
  Al^{3+} + 3e^- \rightarrow Al
• Molten aluminium is drained off and cast into ingots.
• Positive electrodes are made from carbon; oxide ions are oxidized to form oxygen gas:
  2O^{2-} \rightarrow O_2 + 4e^-
• Carbon electrodes are replaced regularly because they react with oxygen.

Extraction of Titanium

• The main titanium ore is rutile (TiO_2).
• Titanium is not extracted by electrolysis due to tree-like crystal formation and side reactions causing impurities.
• Most titanium is extracted using the Kroll process:
  • Titanium(IV) dioxide, coke, and chlorine are heated at about 900°C to form titanium(IV) chloride:
    TiO2(s) + 2C(s) + 2Cl2(g) \rightarrow TiCl_4(g) + 2CO(g)
  • Magnesium is used as a reducing agent to form titanium:
    TiCl4(g) + 2Mg(l) \rightarrow Ti(s) + 2MgCl2(l)
• The Kroll process is expensive due to high energy requirements and the batch process nature.

Choosing between Aluminium and Titanium

• Material properties must be evaluated in relation to the desired use (e.g., bicycle frames).

• Aluminium frames are durable, strong, lightweight, rigid, and cheaper than titanium.
• Titanium is about 1.6 times stronger and nearly twice as stiff as aluminium but is costlier (4 times more expensive per tonne).
• Aluminium would likely be chosen due to cost-effectiveness.

Useful Products from Electrolysis of Brine

• The chlor-alkali industry produces useful materials by electrolyzing NaCl(aq) (brine).

Reactions

• Electrolysis of brine produces chlorine, hydrogen, and sodium hydroxide.
• At the positive electrode (anode), chloride ions form chlorine gas (oxidation):
  2Cl^-(aq) → Cl_2(g) + 2e^-
• At the negative electrode (cathode), hydrogen ions form hydrogen gas (reduction):
  2H^+(aq) + 2e^- → H_2(g)
• Remaining sodium and hydroxide ions form sodium hydroxide (NaOH), or caustic soda.

Diaphragm Cell

• This cell is cheaper but requires frequent diaphragm replacement.
• A porous diaphragm allows brine to pass but prevents chlorine and hydrogen gas mixing.
• Sodium hydroxide formed is mixed with brine and requires recrystallization for purification.

Membrane Cell

• This cell is more expensive and requires minimal maintenance.
• An ion-exchange membrane allows positive sodium ions to pass but not negative chloride ions.
• Brine enters only from the anode side, resulting in purer sodium hydroxide at the cathode.

Production Methods in Industry

• Chemists consider many factors when choosing a production method:
  • Electrolyzing a solution vs. a molten compound.
  • Energy costs (e.g., electrolysis vs. metal oxide reduction).
  • Use of a catalyst (cost vs. increased reaction rate).
  • Effect of changing conditions (e.g., pressure).
  • Batch vs. continuous process.

Transition Metals

• The transition metals are in the central block of the periodic table.
• They form stable ions with an incomplete d-subshell.
• Scandium (Sc) and zinc (Zn) are d-block elements but not transition metals because their stable ions have full or empty d-subshells.

Electronic Configurations

• Chromium (Cr^{3+}): 1s^2 2s^2 2p^6 3s^2 3p^6 3d^3
• Zinc (Zn^{2+}): 1s^2 2s^2 2p^6 3s^2 3p^6 3d^{10}

Physical Properties

• High melting and boiling points.
• High densities.
• Good conductors of heat and electricity.
• Stronger and harder than Group 1 and 2 metals.

Formation of Complex Ions

• A complex ion forms when a ligand bonds to a metal ion via a dative covalent bond.
• A ligand is a species with at least one lone pair of electrons to donate to the metal ion.
• Example: Copper sulfate in water forms hexaaquacopper(II) ion ([Cu(H2O)6]^{2+}).

Chemical Properties

• Form coloured compounds.
• Exist in variable oxidation states.
• Act as catalysts.
• React with ligands to form complex ions.

Transition Metals as Catalysts

• Transition metals speed up reactions by providing alternate mechanisms with lower activation energy.

Haber Process

• The Haber process is used to make ammonia:
  N2(g) + 3H2(g) \rightleftharpoons 2NH_3(g)
• Conditions:
  • Temperature of 400 °C.
  • Heterogeneous catalyst of iron.
  • Pressure of 200 atm.
  • Recycling of nitrogen and hydrogen.

Heterogeneous Catalysis

• Catalyst is in a different physical state from reactants.
  1. Reactants adsorb to the catalyst surface.
  2. Bonds in the reactants weaken, and the reaction occurs.
  3. Products desorb from the catalyst surface.

Homogenous Catalysis

• Catalyst is in the same physical state as reactants.

• Example: Fe^{2+}(aq) ions catalyze the reaction between I^-(aq) and S2O8^{2-}(aq).

  • Uncatalyzed reaction is slow due to collision of two negative ions.
  • Catalyzed reaction is two steps involving positive and negative ions.
    2Fe^{2+}(aq) + S2O8^{2-}(aq) \rightarrow 2Fe^{3+}(aq) + 2SO4^{2-}(aq) 2Fe^{3+}(aq) + 2I^-(aq) \rightarrow I2(aq) + 2Fe^{2+}(aq)
  • Overall:
    S2O8^{2-}(aq) + 2I^-(aq) \rightarrow I2(aq) + 2SO4^{2-}(aq)

Contact Process

• The contact process is used to make sulfuric acid.
• It uses vanadium(V) oxide as a heterogeneous catalyst to form sulfur trioxide:
  SO2(g) + V2O5(s) \rightarrow SO3(g) + V2O4(s)
  O2(g) + V2O4(s) \rightarrow V2O_5(s)
• Sulfur trioxide is mixed with water to form sulfuric acid.

Formulae in Organic Chemistry

Molecular Formulae

• Tell the number and type of each atom in a molecule.
  • Butane: C4H{10} (4 carbons, 10 hydrogens).
  • Phenol: C6H6O (6 carbons, 6 hydrogens, 1 oxygen).
• Give no information on atomic arrangement.

Structural Formulae

• Show atomic arrangement without displaying all bonds.
  • Butane: CH3CH2CH2CH3
• Branched groups are in brackets to the right of the carbon atom.
  • 2-methylbutane: CH3CH(CH3)CH2CH3

Displayed Formulae

• Show all bonds present.
• Propene (C3H6):

Skeletal Formulae

• Simplified formulae with hydrogen atoms removed and the carbon chain reduced to lines.

3D Representations

• Use dashes and wedges to represent the three-dimensional nature of molecules
  • Solid lines : in plane
  • Wedges: coming out of the paper
  • Dashed lines: going into the paper

Alkanes

Bonding in Alkanes

• Sigma (\sigma) bonds form between carbons and hydrogens, and between adjacent carbons.

• Single bonds are free to rotate.

• Carbon orbitals involved in bonding are sp^3 hybrid orbitals (formed when 2s and three 2p orbitals rearrange).

Boiling Points of Alkanes

• Longer carbon chain = higher boiling point.

• Stronger London forces due to more electrons and larger electron clouds.

Types of Alkanes

• Straight chain alkanes: single chain of carbon atoms
  • CH3CH2CH2CH2CH2CH3 (hexane)
• Branched alkanes: one or more carbon atoms attached to main chain
  • (2-methylpentane)
• Cyclic alkanes: carbon atoms joined in a ring
  • (cyclohexane)

Structural Isomers

• Molecules with the same molecular formula but different structural formulae
• Hydrocarbons can have isomers due to branches on the carbon chain.
  • Pentane
  • 2-methylbutane
  • dimethylpropane
• Isomers are chemically similar but have different physical properties like boiling points.
• Branching reduces contact points, lowering London forces.