Chemistry Notes for Grade 12

THEME 1: THE PARTICULATE NATURE OF MATTER AND STOICHIOMETRY

1.1 ATOMS, MOLECULES AND STOICHIOMETRY

1.1.1 RELATIVE MASSES OF ATOMS AND MOLECULES

• Definitions of Key Terms:
    - Relative Atomic Mass: The weighted average mass of an element's isotopes relative to 2 of the mass of carbon-12.
    - Isotopic Mass: The mass of one atom of a specific isotope compared to 2 of mass of carbon-12.
    - Molecular Mass: The weighted mean mass of molecules relative to 2 of the mass of carbon-12.
    - Formula Mass: The weighted mean mass of formula units (for ionic compounds) relative to 2 of the mass of carbon-12.

• Unified Atomic Mass Unit:
    - The mass of one unified atomic mass is defined as one-twelfth of the mass of a carbon-12 isotope:
      - 1 u = $1.66 imes 10^{-27} ext{ kg}$.

• Calculating Relative Atomic Mass (Ar):
    - Formula:
  $ext{Ar of X} = ( ext{ ext{abundance}} imes ext{Ar (isotope 1)}) + ( ext{abundance} imes ext{Ar (isotope 2)}) + ( ext{abundance} imes ext{Ar (isotope 3)})$

• Example Calculations:
    - Hydrogen (H) with isotopes H-1 (99.99%); H-2 (0.009%); H-3 (0.001%):
      - $ext{Ar} = (99.99/100 imes 1) + (0.009/100 imes 2) + (0.001/100 imes 3) ext{ g/mol}$
    - Copper (Cu) with Cu-63 (69%) and Cu-65 (31%):
      - $ext{Ar} = (69/100 imes 63) + (31/100 imes 65) ext{ g/mol}$
    - Boron (B) with B-10 (19.9%) and B-11 (80.1%):
      - $ext{Ar} = (19.9/100 imes 10) + (80.1/100 imes 11) ext{ g/mol}$

• Relative Isotopic Mass:
    - For example:
      - Relative isotopic mass of $Cl^{17}<em>{35}$ is 35 and of $Cl^{17}</em>{37}$ is 37.

• Relative Molecular Mass (Mr):
    - Mr is defined only for molecular substances and can be found by summing the relative atomic masses of all the atoms in the molecule.
    - Example for Silicon Dioxide $SiO_2$:
      - Mr = $28.1 + (16 imes 2)$ = 60.1.

• Relative Formula Mass (Mr):
    - Similar to Mr but applicable to both ionic and covalent compounds.

1.1.2 THE MOLE AND THE AVOGADRO CONSTANT

• Mole and Avogadro Constant:
    - The Avogadro constant (N_A) is the number of particles (atoms, molecules, ions) in one mole, defined as $N_A = 6.02 imes 10^{23}$.
    - One mole of an element has a mass equal to its relative atomic mass in grams.

• Examples:
    - 12 g of carbon = 1 mole of carbon = $6.02 imes 10^{23}$ atoms.
    - 18 g of water (for $H_2O$):
      - $ext{Mr of } H_2O = (2 imes 1) + 16 = 18 ext{ g/mol}$.

1.1.3 THE DETERMINATION OF THE RELATIVE ATOMIC MASSES (Ar)

• Mass Spectrometry:
    - A mass spectrometer converts atoms to ions and uses a magnetic field for detection:
      - Heavier isotopes are closer to the detector.

• Mass Spectrum:
    - Displays relative abundances on y-axis and mass to charge ratio (m/z) on x-axis.
    - Example: Zirconium mass spectrum displays five peaks, indicating five isotopes.

• Calculating Relative Atomic Mass from Mass Spectrum:
    - $ext{Ar} = ($.

1.1.4 THE CALCULATION OF EMPIRICAL AND MOLECULAR FORMULAE

• Empirical and Molecular Formulas:
    - Empirical Formula shows simplest whole number ratio of atoms.
      - Example: Propene $C_3H_6$ (ratio 3:6) has empirical formula $CH_2$ (ratio 1:2).
    - Molecular Formula shows number and type of each atom in a molecule.
      - Example: Ethanoic acid $C_2H_4O_2$.

• Method for Calculation:
    - Combustion method or mass percentage of each element.
    - Steps include finding moles, determining ratios, deriving empirical formula.

1.1.5 REACTING MASSES AND VOLUMES OF SOLUTIONS AND GASES

• Writing Chemical Equations:
    1. Write a word equation.
    2. Write the chemical equation with state symbols: (s), (l), (g), (aq).
    3. Balance the equation using coefficients.

• Calculating with Masses and Gases:
    - The conservation of mass states that mass cannot be created or destroyed, so:
    1. Write a balanced chemical equation.
    2. Calculate molar masses.
    3. Use ratios to determine masses and volumes.

• Molar Volume:
    - At RTP (room temperature and pressure), 1 mole of gas occupies 24 dm³, and at STP (standard temperature and pressure), 1 mole occupies 22.4 dm³.

1.2 ATOMIC STRUCTURE

1.2.1 PARTICLES IN THE ATOM

• Subatomic Particles:
    - Electrons, protons, and neutrons are essential building blocks of atoms:
      - Electrons:
        - Charge: -1, Mass: 1/1836 (negligible).
      - Protons:
        - Charge: +1, Mass: 1.
      - Neutrons:
        - Charge: 0, Mass: 1.

• Mass Distribution:
    - Most mass is in the nucleus (protons and neutrons), with electrons contributing negligible mass in orbitals.

• Proton and Nucleon Numbers:
    - Proton number (Z): Number of protons (also equals the number of electrons in neutral atoms).
    - Nucleon number (A): Total number of protons and neutrons in the nucleus.

1.2.2 THE NUCLEUS OF THE ATOM

• Definition of Isotopes:
    - Atoms of the same element with different numbers of neutrons and, hence, different mass numbers.

• Nuclear Notation:
    - Element symbol with superscript as nucleon number and subscript as proton number:
      - Example: $^{A}_{Z}X$.

1.3 CHEMICAL BONDING

1.3.1 IONIC BONDING

• Ionic Bond Formation:
    - Occurs when metals transfer electrons to nonmetals, resulting in ions.
    - Cation forms from metal (positive ion), and anion forms from non-metal (negative ion).

• Electrostatic Force in Ionic Compounds:
    - Attraction between oppositely charged ions.
    - Leads to the formation of crystalline solids (ionic lattices).
    - Physical properties influenced include melting points, boiling points, hardness, and brittleness.

Summary Activities and Examples

• Regularly engage in activities such as creating mass spectra, calculating empirical formulas, and exploring bond properties to reinforce concepts.

• Understanding quantum mechanics, electron configuration, and bonding theories can bridge connections between topics.