Nuclear Chemistry Study Notes

NUCLEAR CHEMISTRY

SYLLABUS DOT POINTS (IQ2)

  • Investigation of Isotopes
    • Basic structure of stable and unstable isotopes
    • Examination of their position in the periodic table
    • Distribution of electrons, protons, and neutrons in the atom
    • Representation of the symbol, atomic number, and mass number (nucleon number)
  • Modeling of Atomic Energy Levels
    • Discrete energy levels of the atom
    • Electronic configuration and spdf notation (ACSCH017, ACSCH018, ACSCH020, ACSCH022)
  • Calculation of Relative Atomic Mass
    • Relative atomic mass derived from isotopic composition (ACSCH024)
  • Investigation of Energy Levels in Atoms and Ions
    • Collecting primary data from flame tests using ionic solutions of metals (ACSCH019)
    • Examination of spectral evidence for both Bohr model and introduction of Schrödinger model
  • Investigation of Properties of Unstable Isotopes
    • Exploration of natural and human-made radioisotopes, including:
    • Types of radiation
    • Types of balanced nuclear reactions

LEARNING OUTCOMES

  1. Identify the three subatomic particles in the nucleus of an atom.
  2. Define an isotope.
  3. Calculate the relative atomic mass for elements using the natural abundance of its isotopes.

STRUCTURE OF AN ATOM

  • Composition:
    • Atoms are constituted by:
    • Protons: Positively charged particles located in the nucleus.
    • Neutrons: Neutral particles located in the nucleus.
    • Electrons: Negatively charged particles orbiting the nucleus.

Subatomic Particle Properties


  • Charge and Mass of Subatomic Particles

Subatomic ParticleLocation in AtomChargeMass (amu)
ProtonNucleus+11
NeutronNucleus01
ElectronOrbiting-10.00055

ATOMIC CHARACTERISTICS

  • Effects of Subatomic Particles:
    • Various features such as atomic number (Z), atomic mass (A), and charge are influenced by the number of each subatomic particle.
    • Definitions:
    • Atomic Number (Z): Number of protons in the nucleus.
    • Atomic Mass (A): Total number of protons and neutrons.
    • Charge: Determined by the comparison of proton and electron counts:
      • If proton count > electron count, atom is positively charged (+).
      • If proton count < electron count, atom is negatively charged (-).
      • If proton count = electron count, atom is neutral (0).
    • Relation: A = p + n, where p = protons, n = neutrons.

ISOTOPES

  • Definition:
    • Isotopes are variants of the same chemical element that have identical proton numbers but different neutron counts, which results in different atomic masses.
  • Examples of Carbon Isotopes:
    • Carbon contains at least four isotopes: C-11, C-12, C-13, and C-14.
    • Isotope Details:
      • C-11: 6 protons, 5 neutrons, 5 electrons, atomic mass = 11.
      • C-12: 6 protons, 6 neutrons, 6 electrons, atomic mass = 12.
      • C-13: 6 protons, 7 neutrons, 6 electrons, atomic mass = 13.
      • C-14: 6 protons, 8 neutrons, 6 electrons, atomic mass = 14.

RELATIVE ATOMIC MASS

  • Concept:
    • The relative atomic mass is the weighted average of an element's isotopic composition in nature, accounting for the different abundances of isotopes.
  • Calculation Formula:
    ext{Relative Atomic Mass} = rac{M1 imes x1 + M2 imes x2 + …}{100}
    where M = mass of isotope and x = percentage of that isotope.
  • Example Calculation:
    • For Carbon with stable isotopes 12C (98.9% abundance) and 13C (1.1% abundance):
      ext{Relative Atomic Mass} = rac{(98.9 imes 12) + (1.1 imes 13)}{100} = 12.01 ext{ amu}.
  • Sample Problems:
    1. For element X with isotopes 50X (78.2%), 51X (17.3%), and 53X (4.5%), calculate its relative atomic mass.
    2. For Cl with isotopes Cl-35 and Cl-37, if the average atomic mass is 35.45 amu, determine each isotope's percentage in the sample.

NUCLEUS STABILITY

LEARNING OUTCOMES

  1. Outline criteria for stable nuclei.
  2. Describe types of radioactive decay an unstable nucleus may undergo.
  3. Write nuclear equations for radioactive decay and other radioactive processes.

Factors Affecting Nucleus Stability

  • Determination of nucleus stability depends on the neutron-to-proton (n:p) ratio:
    • An ideal neutron count exists for a specific number of protons to form the most stable nucleus.
    • Less stable isotopes exhibit shorter half-lives.
  • Half-life Examples:
    • C-11: 20 min
    • C-12: Stable
    • C-13: Stable
    • C-14: 5730 years
    • C-15: 2.5 s

Zone of Stability

  • The visual representation for nucleus stability:
    • An unstable nucleus can be caused by:
    1. n:p ratio outside the zone of stability:
      • For Z = 1-20, n:p = ~1
      • For Z = 21-50, n:p = ~1.3
      • For Z = 51-80, n:p = ~1.5
    2. When Z > 82, the nucleus becomes too large with excessive protons that the strong nuclear force can no longer overcome repulsions.

RADIOACTIVE DECAY

  • Definition:
    • Radioactive decay is the spontaneous emission of radiation from an unstable nucleus seeking to achieve stability.
    • Parent Isotope: Original unstable isotope
    • Daughter Isotope: Isotope resulting from decay

Types of Radioactive Decay

  • Key Decay Types:
    1. Alpha (α) Decay
    • Involves emission of a helium nucleus (alpha particle).
    • Occurs in heavy unstable nuclei (Z > 82).
    • General equation:
      \text{X} \rightarrow \text{Y} + 4_2\text{He}
    • Example:
      ^{226}{88}\text{Ra} \rightarrow ^{222}{86}\text{Rn} + ^{4}_{2}\text{He}
    • Penetrating Power: Low - stopped by paper.
    • Ionising Power: High - harmful if inhaled or ingested.
    1. Beta-negative (β-) Decay
    • Emission of electrons from a neutron decay.
    • Occurs in nuclei with excessive neutrons (high n:p ratio).
    • General equation:
      10\text{n} \rightarrow 11\text{p} + 0_{-1}\text{e}
    • Example:
      ^{14}{6}\text{C} \rightarrow ^{14}{7}\text{N} + 0_{-1}\text{e}
    • Penetrating Power: Medium - stopped by 0.5 mm lead.
    • Ionising Power: Medium.
    1. Beta-positive (β+) Decay
    • Emission of positrons from a proton decay.
    • Occurs in nuclei with excessive protons (low n:p ratio).
    • General equation:
      11\text{p} \rightarrow 10\text{n} + 0_1\text{e}^+
    • Example:
      ^{10}{6}\text{C} \rightarrow ^{10}{5}\text{B} + 0_1\text{e}^+
    • Penetrating Power: Medium - similar to β- decay.
    • Ionising Power: Medium.
    1. Gamma (γ) Decay
    • Involves emission of gamma rays (electromagnetic radiation in the form of photons).
    • Occurs from excited nuclei and often accompanies α and β decay.
    • Penetrating Power: High - requires dense materials to shield.
    • Ionising Power: Low.

OTHER NUCLEAR PROCESSES

  • Nuclear Processes:
    1. Fusion: Combining two lighter nuclei into a heavier nucleus.
    2. Fission: Splitting a large nucleus into smaller nuclei.

Balancing of Nuclear Chemical Equations

  • Nuclear equations must maintain conservation laws, conserving both atomic number and mass.
  • Example reaction for alpha decay:
    • ^{25}{12}\text{Mg} + ^{4}{2}\text{He} \rightarrow ? + ^{1}_{1}\text{p}
  • Extend balancing using conservation laws to find unknown nucleotides.

UNDERSTANDING PRACTICE

  1. Identify the type of radioactive decay for given unstable nuclei and write corresponding nuclear equations.
  2. Investigate Technetium-99m and Cobalt-60's formation, radioactive decay processes, and applications in the medical field.