Types of Chemical Reactions

Types of Reactions

• Synthesis: Two or more substances combine to form a new substance. Example: A + B → AB
• Decomposition: Substance breaks down into two or more substances due to heat, light or electricity. Example: AB → A + B
• Acid-base (Neutralisation): Acid reacts with a base to form salt and water. Example: HCl(aq) + NaOH(s) → NaCl(aq) + H2O(l)
• Combustion Reactions: Substances burn in oxygen.
  • Complete Combustion: Forms carbon dioxide and water. Example: CH4(g) + 2O2(g) → CO2(g) + 2H2O(l)
  • Incomplete Combustion: Forms carbon or carbon monoxide. Examples: CH4(g) + O2(g) → C(s) + 2H2O(l) and CH4(g) + 2/3O2(g) → CO(g) + 2H2O(l)
• Precipitation Reactions (Double Displacement): Two soluble ionic compounds react to form an insoluble precipitate (ppt) or mix completely. Example: AB + CD → AD + CB
  • Precipitation Example: AgNO3(aq) + NaCl(aq) → NaNO3(aq) + AgCl(s)
• Acid-carbonate reactions: Acid reacts with carbonate to form water, salt, and carbon dioxide. Example: Na2CO3(s) + 2HCl(aq) → 2NaCl(aq) + H2O(l) + CO2(g)
• Metal - Oxygen Reactions: Metal burns in oxygen to form a metal oxide. Example: 2Mg(s) + O2(g) → 2MgO(s)
• Metal - Water Reactions: Metal reacts with water to form a metal hydroxide and hydrogen gas. Example: 2Na(s) + 2H2O(l) → 2NaOH(s) + H2(g)
• Acid-metal Reactions: Acid reacts with metal to form a metal salt and hydrogen gas. Example: Zn(s) + H2SO4(aq) → ZnSO4(aq) + H2(g)
• Metal Displacement Reactions: More reactive metal displaces a less reactive metal in a compound. Example: A + BC → AC + B
  • Displacement Example: Zn(s) + CuSO4(aq) → ZnSO4(aq) + Cu(s)
  • Metal reactivity determined by the metal activity series or reduction potentials table. Lower ionisation energy, larger atomic radius, and lower electronegativities indicate higher reactivity.

Atomic Structure

• Principal Energy Levels (1, 2, 3…): Groupings of sublevels (s, p, d, f) with similar energy.
• Sublevels (s, p, d, f): Groupings of individual orbital lobes.
• Aufbau Principle: Electrons fill lowest available energy levels first before occupying higher levels.
• Exceptions to Aufbau Principle: Occur in the d block due to reduced energy difference between s and d levels, causing Coulombic repulsion. Examples: Cu, Ag, Cr.
• Heisenberg Uncertainty Principle: The position and momentum of a particle (like an electron) cannot be measured accurately at the same time.
• Pauli Exclusion Principle & Hund's Rule: Each orbital within an energy level is filled first by electrons with spin +1/2 and then completed with electrons of -1/2 spin (identical fermions cannot occupy the same quantum state within the same system). Example: Oxygen
• Valence Electrons: Electrons in the outermost shell of an atom.
• Valency: Maximum number of direct bonds that an element can form. Variable valency can occur when an atom loses or gains more electrons than its valence shell.
• Orbital Notation: Determined by filling orbitals based on the number of electrons or using spdf blocks on the periodic table.
  • Helium and groups 1 and 2: s orbitals from energy levels 1-7
  • Transition metals: d orbitals from energy levels 3-6
  • Groups 13-18 (excluding Helium): p orbitals from energy levels 2-7
  • Lanthanides and Actinides: f orbitals from energy levels 4-5

Atomic Emission Spectra and Flame Tests

• Excitation: Electrons gain energy, move to a higher energy level.
• De-excitation: Excited electrons return to their ground state, releasing energy as different wavelengths of light.

Redox Reactions

• Half-Equations: Show electron transfer.
• Neutral Species Equations: Show compounds involved.
• Complete Ionic Equation: Shows all ions in the reaction mixture.
• Net Ionic Equation: Shows only the ions undergoing the reaction, excluding spectator ions.
• Deriving Equations:
  1. Balanced Formula (Neutral Species) Equation: 2KI(aq) + Pb(NO2)3(aq) → PbI2(s) + 2KNO3(aq)
  2. Total Ionic Equation (Complete Ionic): 2K+(aq) + 2I-(aq) + Pb2+(aq) + 2NO3-(aq) → PbI2(s) + 2K+(aq) + 2NO3-(aq)
  3. Net Ionic Equation: Pb2+(aq) + 2I-(aq) → PbI2(s)

Galvanic Cells and Standard Electrode Potentials

• Galvanic Cells: Generate electricity via redox reactions.
  • Anode: Oxidation occurs, producing free electrons in the more reactive metal.
  • Cathode: Electrons flow to the cathode, resulting in reduction.
  • Salt Bridge: Completes the circuit, conserving solution neutrality.
    • Anions flow toward the anode to balance positive ions.
    • Cations flow toward the cathode to balance negative charges.
• Standard Hydrogen Half-Cell: Platinum metal in a 1 mol/L H+ solution with hydrogen gas bubbled over at 100 kPa.
• Standard Electrode Potentials: Higher values indicate a stronger tendency to form the non-ionic form; lower values suggest a tendency to form ions.
• Voltage of a Galvanic Cell under Standard Conditions: Etotal = E°reduction + E°oxidation
  • Galvanic cells are spontaneous when cell voltage is positive.
  • Oxidizing agents are reduced, and vice versa.
• Galvanic Cell Notation: X(s) | X+(aq) || Y+(aq) | Y(s) with the anode/oxidation on the left.
• Galvanic Cell Example:
  • Copper gains electrons at the cathode.
  • Zinc loses electrons at the anode.
  • Notation for Zn(s) | Zn2+(aq)(1M) || Cu2+(aq)(1M) | Cu(s)

Rates of Reactions

• Collision Theory: Reactions occur due to collisions between reactant particles.
  • Requires:
    • Collisions between particles
    • Sufficient energy to break bonds and reach activation energy
    • Correct orientation to break bonds and form products
• Activation Energy: Minimum energy required for a collision to lead to reaction.
• Measuring Reaction Rates: Use time or loss of mass.
• Factors Influencing Reaction Rates:
  • Surface Area: Higher surface area leads to more collisions.
  • Temperature: Higher temperature increases kinetic energy, increasing collision probability.
  • Concentration: Higher concentration means more particles, increasing collision probability.
  • Pressure (Gases): Higher pressure brings particles closer, increasing collision probability.
  • Presence of a Catalyst: Catalysts speed up reactions, but are not consumed. They provide an alternative pathway with lower activation energy.
  • Examples:
    • Iron (Fe) in ammonia production: N2(g) + 3H2(g) → 2NH3(g) ΔH = -92 kJ/mol
    • Platinum (Pt) in carbon monoxide to carbon dioxide: CO(g) + 1/2 O2(g) → CO2(g)
    • Potassium permanganate (KMnO4) in hydrogen peroxide decomposition: 2H2O2(aq) → 2H2O(l) + O2(g)
    • Chlorophyll in leaves for photosynthesis
  • Heterogeneous Catalysts: Different phase from reactants (e.g. gas - liquid)
  • Homogeneous Catalysts: Same phase as reactants (e.g. liquid - liquid)
  • Homogeneous Reactions: Occur in a single phase (e.g. liquid - liquid)

Drivers of Reactions

• Chemical Potential Energy: All substances hold chemical potential energy.
• Bond Breaking: Requires energy.
• Bond Making: Releases energy.
• Exothermic Reactions: Release heat into the surroundings.
  • Energy released during bond making is greater than the energy needed to break bonds.
• Endothermic Reactions: Absorb heat from the surroundings.
  • Energy required to break bonds is greater than the energy released during bond making.
• Enthalpy: Total heat content of a system, measured in joules (J).
• Change in Enthalpy (ΔH): ΔH = Hproducts - Hreactants where H is heat content.
• Heat of Combustion (ΔHc): Energy released when 1 mol of a substance combusts completely at SLC.
  • Energy released for n moles of fuel burning: Energy released = n x ΔHc (when ΔHc is known)
• Determining Heat of Combustion Experimentally: q = mcΔT where q is energy, c is specific heat capacity (often 4.18 for water), and ΔT is temperature change.
• Specific Heat Capacity: Amount of energy to raise the temperature of 1g of a substance by 1K. Water has a high specific heat capacity (4.18), requiring lots of energy for temperature increases.