G. C. E. (Advanced Level) Chemistry - Grade 12 Resource Book Study Notes

G. C. E. (Advanced Level) Chemistry - Grade 12 Resource Book

Unit Overview

  • Unit 1: Atomic Structure

  • Unit 2: Structure and Bonding

  • Unit 3: Chemical Calculations

  • Unit 6: Chemistry of s, p, and d Block Elements

  • Publisher: Department of Science, National Institute of Education, Sri Lanka. 2021

Messages

Message from the Director General

  • The National Institute of Education is taking steps to improve educational quality.

  • Supplementary resource books have been developed by curriculum experts to align with the GCE A/L syllabus introduced in 2017.

  • The resource books aim to enhance teachers' ability to plan effective learning and teaching activities, and to broaden students' understanding of chemistry.

Message from the Director

  • A refined curriculum is in effect since 2017, which includes revisions in subject content, delivery methods, and curricular materials.

  • A new Teachers' Guide was introduced, presenting learning outcomes and guidelines for effective teaching.

  • The resource book is accessible in Sinhala, Tamil, and English, providing reliable information consistent with the local curriculum.

Key Contributors

  • Supervision: Dr. (Mrs.) T. A. R. J. Gunasekara, Dr. A. D. A. De Silva

  • Authors:

    • Dr. Russel C. L. de Silva - Unit 1

    • Dr. M.A.B. Prasantha - Unit 2

    • Dr. M.N. Kaumal - Unit 3 and Unit 6

  • Editing and Support Staff were involved in the creation and review process, ensuring high standards.

Unit 1: Atomic Structure

1.1 The Atomic Theory of Matter

  • Matter: Anything that has mass and occupies space.

  • Atoms: The basic units of matter, originally theorized by Democritus as indivisible particles.

  • John Dalton’s Atomic Theory (1808):

    1. Atoms are indivisible.

    2. All atoms of a particular element are identical.

    3. Atoms are neither created nor destroyed in reactions.

    4. Compounds are formed from atoms of different elements in simple ratios.

  • Golf Ball Model: Dalton's visualization of atoms.

1.1.1 Properties of Cathode Rays

  • Cathode rays travel in straight lines, exhibit mass, and can be deflected by magnetic and electric fields.

  • Carries a negative charge, with experiments confirming the existence of electrons.

1.1.2 The Nucleus of the Atom

  • Positive Rays: Named by Eugen Goldstein from his canal ray experiments.

  • Rutherford’s Gold Foil Experiment (1909): Showed that the nucleus is dense and positively charged, proposing the nuclear model of the atom.

1.1.3 Properties of Positive Rays

  • Positive rays exhibit charge, can impact matter, and are affected by electric fields.

  • Identified protons through their different masses in various gases.

1.1.4 Atomic Number, Isotopes, and Mass Number

  • Atomic Number (Z): Number of protons in an atom.

  • Isotopes: Atoms of the same element with different neutron counts and mass numbers.

  • Mass Number (A): Total number of protons and neutrons.

  • Example: Representation of carbon isotopes – 12C, 13C:

    • 12C: 6 protons, 6 neutrons

    • 13C: 6 protons, 7 neutrons

1.1.5 The Atomic Mass Scale

  • The uniform atomic mass unit (u) represents $ rac{1}{12}$ of the mass of a carbon-12 atom.

  • Average atomic mass accounts for isotopic abundances.

1.1.6 Average Atomic Mass and Relative Atomic Mass

  • Average Atomic Mass Formula: ( ext{Average Atomic Mass} = rac{ ext{Σ(isotope mass)}}{ ext{Σ(fractional isotope abundance)}})

1.1.7 Ions

  • Cations: Positively charged ions formed by electron loss.

  • Anions: Negatively charged ions formed by electron gain.

1.2 Electromagnetic Radiation and Wave-Like Properties of Matter

  • Electromagnetic Radiation (EMR): Characterized by speed (c), wavelength ($ ext{λ}$), frequency ($ ext{ν}$), and energy ($E$).

  • Key Relationships:

    • c =
      u imes ext{λ}

    • E = h
      u (with $h = 6.626 imes 10^{-34} J s$)

1.2.1 Quantization of Energy

  • Energy emitted or absorbed occurs in discrete packets, termed quanta.

1.3 Electronic Energy Levels of Atoms

  • Principal Quantum Number (n): Defines energy levels; higher $n$ indicates further distance from the nucleus.

  • Quantum Numbers: Define every electron in an atom, with $l, ml, ms$.

1.3.1 The Hydrogen Spectrum

  • Hydrogen emits certain wavelengths due to energy transitions of electrons between levels.

1.4 Electron Configuration

  • Aufbau Principle: Electrons fill first to lowest energy orbitals.

  • Pauli Exclusion Principle: No two electrons can have the same set of quantum numbers (each orbital holds two electrons).

  • Hund's Rule: Every orbital in a subshell must be filled with one electron before any orbital is filled a second time.

1.5 Building of Periodic Table

  • Periodicity of elements is based on the electron configurations of outer shells.

1.6 Periodic Trends Shown by s and p Block Elements
1.6.1 Sizes of Atoms and Ions

  • Decreasing size from left to right due to increasing effective nuclear charge (Zeff).

  • Trends in atomic radius are influenced by number of electron shells and nuclear charge.

1.6.2 Ionization Energy

  • Energy required to remove an electron; decreases down a group and increases across a period.

1.6.3 Electron Gain Energy

  • Energy change when an electron is added; becomes more positive down the group.

1.6.4 Electronegativity

  • Attractiveness of an atom in a molecule to electrons; increases from left to right and decreases down the group.

Unit 2: Structure and Bonding

2.1 Covalent Bonds

  • Covalent bonds form by sharing electrons; defined by Lewis structures.

2.1.1 Lewis Dot Diagrams

  • Visual representations that show valence electrons and bonding within molecules.

2.2 Dative Covalent Bonds

  • Bonds formed when one atom provides both electrons for a bond.

2.3 VSEPR Theory

  • Describes molecular shapes based on minimizing electron pair repulsions.

    • Electron pair geometries: Linear, trigonal planar, tetrahedral, etc.

2.4 Ionic Bonds

  • Result from electrostatic attraction between cations and anions.

2.5 Metallic Bonds

  • Formed by attraction between metal ions and delocalized electrons.

2.6 Secondary Interactions

  • Types include ion-dipole, dipole-dipole interactions, etc.

Unit 3: Chemical Calculations

3.1 Oxidation Number

  • Used to track electron transfers in reactions.

3.2 Nomenclature of Inorganic Compounds

  • Systematic names based on IUPAC guidelines.

3.4 Types of Chemical Formulae

  • Empirical Formula: Simplest integer ratio.

  • Molecular Formula: Actual numbers of atoms.

3.6 Balancing Chemical Reactions

  • Methods for ensuring mass conservation in chemical equations.

Unit 6: Chemistry of s, p and d Block Elements

4.1 Group 1 Elements

  • Alkali metals, very reactive, lower melting points.

4.2 Group 2 Elements

  • Alkaline earth metals, less reactive, higher density than Group 1.

4.4 Group 14 Elements

  • Exhibit allotropy (different structural forms of the same element).

4.10 Transition Elements

  • Defined by partially filled d orbitals; exhibit variable oxidation states.