QCAA Chemistry Units 1 and 2 Comprehensive Study Notes

UNIT 1: Chemical Fundamentals — Structure, Properties and Reactions

Topic 1: Properties and Structure of Atoms

1.1.1 Atomic Structure

• Model of the Atom: Atoms are modeled as a nucleus surrounded by electrons that exist in distinct energy levels.

• Nuclear Components: Separation between the atomic number ($Z$), mass number ($A$), and the nature of isotopes for any given element.

• Nuclear Symbol Notation: Utilization of the notation $^A_Z\text{M}$ to determine the totals of protons, neutrons, and electrons within atoms, ions, and isotopes.

• Orbital Energy Levels: The relative energies for the s, p, and d orbitals must be clearly stated.

• Electron Configuration Principles: The writing of electron configurations for atoms and ions up to $Z = 36$ is governed by:

  • Aufbau Principle: Electrons fill lower-energy orbitals first.

  • Hund's Rule: Every orbital in a subshell is singly occupied with one electron before any one orbital is doubly occupied.

  • Pauli Exclusion Principle: No two electrons in the same atom can have the identical four quantum numbers; an orbital can hold a maximum of two electrons with opposite spins.

• Configuration Formats: Electron configurations up to $Z = 36$ can be expressed in full (e.g., $1s^2\, 2s^2\, 2p^6\, 3s^2\, 3p^5$) or condensed form (e.g., $[Ne]3s^2\, 3p^5$).

• Electron Configuration Exceptions: Chromium ($Cr$) and Copper ($Cu$) are identified as key exceptions to standard electron filling rules.

• Ionisation Energy: Data regarding successive ionisation energies is directly related to the specific electron configuration of an atom.

1.1.2 Isotopes

• Definition: Isotopes are atoms belonging to the same element that possess different numbers of neutrons.

• Isotope Representation: Representation follows the IUPAC form $^A\text{X}$ or the hyphenated form X-A.

• Properties of Isotopes: Isotopes of a single element maintain the same electron configuration and exhibit similar chemical properties, but they possess different physical properties.

• Relative Atomic Mass: Defined as the ratio of the weighted average mass per atom of the naturally occurring element to $\frac{1}{12}$ the mass of an atom of carbon-12.

• Mass Spectrometry: This analytical technique involves the ionisation of substances followed by the separation and detection of resulting ions. Understanding the operation of the instrument itself is not required.

• Data Analysis: Analysis of mass spectrometry spectra is used to determine isotopic composition, calculate relative atomic mass, and find the percentage abundances of isotopes.

1.1.3 Analytical Techniques

• Spectra Types: Discrimination between absorption line spectra and emission line spectra.

• Electron Transfer: Flame tests and atomic absorption spectroscopy (AAS) are analytical methods that rely on the transfer of electrons between discrete atomic energy levels.

• Hydrogen Emission Spectrum: Provides evidence for the Bohr model which posits discrete electron energy levels; these levels are observed to converge at higher energies.

• Practical Application: Analysis of flame tests and AAS data allows for the identification of elements and the determination of the concentration of metallic ions in a solution.

1.1.4 Periodic Table and Trends

• Representation: Elements are represented by specific symbols.

• Structural Basis: The periodic table is structured based on increasing atomic number ($Z$).

• Sub-level Blocks: The table is organized into four distinct blocks (s, p, d, and f) associated with their respective sub-levels.

• Configuration Relationship: There is a direct relationship between the structure of the periodic table and the electron configuration of the atoms within it.

• Trends Across Periods and Groups: Trends in chemical and physical properties are observed across periods and down groups, particularly for groups 1, 2, 13–18 and period 3.

• Metallic and Non-metallic Behavior: Trends include the reactivity of alkali metals (Li–Cs) and halogens (F–I).

• Period 3 Oxide Trends: Oxides across period 3 transition from basic, through amphoteric, to acidic.

• Periodic Data Analysis: Patterns are determined through analysis of data for:

  • Atomic radii.

  • Valencies.

  • Ionic radii.

  • First ionisation energy.

  • Electronegativities.

1.1.5 Introduction to Bonding

• Valence Shell Stability: Chemical bonding ability relates to electron arrangement and the stability of the valence shell.

• Electron Transfer and Sharing: Configuration determines the specific number of electrons lost, gained, or shared.

• Transition Element Ions: Elements in the transition block can form more than one type of ion.

• Ion Definition: Atoms or groups of atoms that carry an electrical charge because of an imbalance between protons and electrons.

• Nature of Chemical Bonds: Caused by electrostatic attractions resulting from the sharing or transfer of electrons.

• Valency: A measure indicating the number of bonds an atom is capable of forming.

• Naming and Formulas: Formulas and IUPAC names are determined for ionic and molecular compounds.

• Formula Types: Discrimination is made between empirical formulas, molecular formulas, and formula units.

• Lewis Structures: Electron dot structures for molecules and ions show all valence electrons for up to four electron pairs per atom.

• Electron Pairs: Identification of bonding pairs versus lone pairs of electrons surrounding atoms in a molecule.

Topic 2: Properties and Structure of Materials

1.2.1 Compounds and Mixtures

• Pure Substances: These may exist as either elements or compounds and possess distinct, measurable properties such as melting point, boiling point, reactivity, strength, and density.

• Mixture Properties: The properties of a mixture are dependent upon the identity and the relative amounts of its components.

• Mixture Classification: Discrimination between heterogeneous mixtures and homogeneous mixtures.

1.2.2 Bonding and Properties

• Property Categories: Comparison of melting/boiling point, thermal/electrical conductivity, strength, and hardness across ionic, covalent, and metallic compounds.

• Determining Physical Properties: Physical properties are dictated by the type of bonding present (ionic, metallic, or covalent) within a substance.

• Ionic Bonding Model: Described as ions arranged in a crystalline lattice held by strong electrostatic attractions between oppositely charged ions.

• Covalent Bonding Model: Described as a shared electron pair between two nuclei with strong electrostatic attraction.

• Covalent Variations: Discrimination between simple covalent molecules, giant covalent networks, and the various allotropes of carbon.

• Hydrocarbons: Alkanes (saturated), alkenes (unsaturated), and benzene possess different chemical properties determined by the nature of their specific bonding.

Topic 3: Chemical Reactions — Reactants, Products and Energy Change

1.3.1 Chemical Reactions

• Energy and Phase Changes: Reactions and phase changes involve energy changes, appearing as light emission or temperature changes.

• Equation Balancing: Determination of balanced chemical equations including state symbols $(s)$, $(l)$, $(g)$, and $(aq)$. Reaction types include:

  • Single displacement.

  • Double displacement.

  • Acid-base.

  • Combustion.

  • Combination.

  • Decomposition.

  • Simple redox reactions.

1.3.2 Exothermic and Endothermic Reactions

• Heat vs. Temperature: Heat is a form of energy; temperature measures the average kinetic energy of the particles.

• Bonding and Energy Conservation: Endothermic and exothermic reactions relate to the law of conservation of energy regarding the breaking and reforming of chemical bonds.

• Enthalpy Diagrams: Sketching and analysis of enthalpy level diagrams for exothermic and endothermic reactions indicate the sign of $\Delta H$ and the relative stabilities of reactants and products.

• Bond Enthalpies: Reactions are classified as exothermic or endothermic based on average bond enthalpies. Limitations exist when using average bond enthalpies to calculate enthalpy change.

• Heat Change Formula ($Q$): Calculation using mass, specific heat capacity, and temperature change:

  • $Q = mc\Delta T$

• Enthalpy Change Formula ($\Delta H$): Calculation given quantities of reactants and temperature changes:

  • $\Delta H = H_{\text{products}} - H_{\text{reactants}}$

• Specific Heat Calculations: Analysis of data for heat of combustion, heat of neutralisation, and aqueous reactions to find mass, specific heat, temperature, and enthalpy changes.

1.3.3 Mole Concept and Law of Conservation of Mass

• The Mole: A precisely defined quantity of matter equal to Avogadro's number of particles.

• Law of Conservation of Mass: Mass is neither created nor destroyed in a chemical reaction.

• Mole Calculations: The mole concept relates mass ($m$), moles ($n$), and molar mass ($M$):

  • $n = \frac{m}{M}$

• Formula Determination:

  • Percentage composition from relative atomic masses.

  • Empirical formula from percentage composition by mass.

  • Molecular formula from empirical formula and molar mass.

• Yield and Reactants: Determination of limiting reactants and discrimination between experimental and theoretical yield.

• Percentage Yield: Calculated via the formula:

  • $\text{percentage yield (\%)} = \left( \frac{\text{experimental yield}}{\text{theoretical yield}} \right) \times 100$

UNIT 2: Molecular Interactions and Reactions

Topic 1: Intermolecular Forces and Gases

2.1.1 Intermolecular Forces

• VSEPR Theory: Valence Shell Electron Pair Repulsion theory determines the shape and bond angles of molecules, including linear, bent, trigonal planar, tetrahedral, and pyramidal shapes (d-orbital hybridisation is not required).

• Molecular Polarity: Determined through molecular shape, symmetry, and comparisons of electronegativity.

• Physical Property Relationships: Vapour pressure, melting point, boiling point, and solubility are related to the strength and nature of intermolecular forces:

  • Dispersion forces.

  • Dipole-dipole attractions.

  • Hydrogen bonding.

2.1.2 Chromatography Techniques

• Composition and Purity: Paper and thin layer chromatography (TLC) are used to determine substance purity.

• Phase Interaction: Separation of components occurs due to variations in interaction strength between the mobile phase and the stationary phase.

• Chromatograph Analysis: Determination of composition and purity involves calculating $R_f$ values.

2.1.3 Gases

• Molar Volume: The relationship between the volume of a gas, moles, and molar volume at STP.

• Kinetic Theory of Gases: Explains relationships between pressure, temperature, and volume as they apply specifically to ideal gases.

• Ideal Gas Equation: Applied to calculate mass or volume of gas at STP:

  • $PV = nRT$

Topic 2: Aqueous Solutions and Acidity

2.2.1 Aqueous Solutions and Molarity

• Water Properties: The unique properties of water are related to its molecular shape and the presence of hydrogen bonding.

• Terminology: Discrimination between solute, solvent, and solution; and between strength and concentration in acids and bases.

• Concentration Notation: Square brackets $[ ]$ denote concentration.

• Solubility States: Discrimination between unsaturated, saturated, and supersaturated solutions.

• Molarity Formula: Application of the mole concept to concentration ($c$), moles ($n$), and volume ($V$):

  • $c = \frac{n}{V}$

2.2.2 Identifying Ions in Solution

• Precipitation Reactions: Use of ionic and chemical formulas to construct balanced equations (including states).

• Solubility Rules: Applied to predict whether a precipitate will form.

• Ion Identification: Analysis of precipitation and acid-carbonate reactions to determine the presence of specific ions.

2.2.3 Solubility

• Substance Solubility: Comparison of ionic and molecular substance solubility in water and the intermolecular forces at play.

• Temperature Effects: Solvent temperature changes affect the solubility of gaseous and solid solutes.

• Solubility Curves: Data analysis to determine concentration of ions in aqueous solutions and the solubility of compounds.

2.2.4 pH

• Hydrogen Ion Concentration: pH is dependent on $[H^+]$ in solution.

• Logarithmic Scale: The pH scale operates as a logarithmic scale.

• Arrhenius Model: Explains the behaviour of strong and weak acids and bases within aqueous solutions.

2.2.5 Reaction of Acids

• Reaction Types: Balanced chemical and ionic equations (including state symbols) for reactions of acids with:

  • Bases.

  • Metals.

  • Carbonates.

Topic 3: Rates of Chemical Reactions

2.3.1 Rates of Reactions

• Rate Factors: Reaction rates are affected by temperature, surface area, pressure (in gaseous systems), concentration, and the presence of a catalyst.

• Collision Theory: Explains how concentration, temperature, pressure, and surface area affect the frequency and energy of collisions.

• Maxwell-Boltzmann Distribution: Curves sketched for reactions to show energy distributions with and without catalysts.

• Activation Energy ($E_a$): Description of the energy barrier for a reaction; the relationship between reactant bond strength/number and the magnitude of $E_a$.

• Energy Profile Diagrams: Sketching and analysis of diagrams to determine enthalpy change and activation energy for catalyzed and uncatalyzed reactions.

• Rate Calculation: Reaction rates are calculated by measuring the rate of product formation or reactant depletion:

  • $\text{rate} = \frac{\Delta[P]}{\text{time}}$

  • $\text{rate} = - \frac{\Delta[R]}{\text{time}}$

• Graphical Analysis: Analysis of concentration, volume, and mass changes against time to determine the rate (order of reaction is not required).