Atomic Structure, Isotopy & Periodic Classification
Subatomic particles (protons, neutrons, electrons), atomic number, mass number, isotopes, electron configuration (first 20 elements), groups, periods, and periodic trends (atomic radius, electronegativity, ionization energy).
In Ghana, nuclear medicine at the National Centre for Radiotherapy and Nuclear Medicine (Korle-Bu Teaching Hospital) uses radioactive isotopes such as Iodine-131 to treat thyroid cancer and Cobalt-60 for tumor irradiation.
Key Syllabus Learning Objectives
- 1Describe the properties and locations of protons, neutrons, and electrons within the atom
- 2Calculate mass numbers, atomic numbers, neutron counts, and relative atomic mass from isotopic abundances
- 3Write electronic configurations for the first 20 elements using the 2,8,8 notation
- 4Explain periodic trends across periods and down groups (atomic radius, ionization energy, electronegativity)
Periodic Trends Down Groups and Across Periods
- Across a Period (left to right): Nuclear charge (protons) increases while electrons are added to the same shell. Stronger electrostatic attraction pulls electrons inward: Atomic radius decreases; Electronegativity increases; First ionization energy increases.
- Down a Group (top to bottom): An additional electron shell is added at each step, increasing electron shielding of outer electrons: Atomic radius increases; Ionization energy decreases; Reactivity of metals increases, while reactivity of halogens decreases.
- Noble Gases: Have stable duplet (He: 2) or octet (Ne, Ar: 2,8; 2,8,8) configurations, requiring immense energy to gain or lose electrons, rendering them chemically inert.
Sodium metal stored under kerosene must be handled with forceps in high school labs because it reacts violently with moisture on contact.
Atomic radius decreases across a period due to increasing nuclear charge pulling electrons closer, and increases down a group due to additional electron shells.
Isotopes & Relative Atomic Mass Calculation
- Chlorine exists as 75% ³⁵Cl and 25% ³⁷Cl. Its RAM is 35.5, which is why atomic masses are not always whole numbers.
- Carbon-12 is the international standard against which all atomic masses are compared (defined as exactly 12.000 atomic mass units).
- Chemical properties of isotopes are identical because they have the same number of valence electrons; physical properties (density, boiling point) differ slightly due to mass difference.
Radiocarbon dating of wooden artifacts from the ancient Ashanti Empire relies on the steady radioactive decay of Carbon-14.
Isotopes share identical chemical reactivity because chemical bonding depends solely on valence electrons, not neutron mass.
- When writing electronic configurations, never write commas as dots; use commas (e.g. 2, 8, 8, 1).
- When asked why isotopes have identical chemical properties, the WAEC marking key expects: "They have the same number of valence electrons / same electronic configuration".
- Distinguish clearly between mass number (whole integer for a single atom) and relative atomic mass (weighted average, often fractional).
- Confusing atomic number (protons) with mass number (protons + neutrons).
- Writing that isotopes have different numbers of protons (which would make them completely different elements).
- Placing more than 8 electrons in the second or third energy level for elements up to Calcium.
- Can I calculate the relative atomic mass given isotopic percentages?
- Can I write the electronic configuration for any element from Hydrogen (1) to Calcium (20)?
- Can I deduce the Group and Period of an element directly from its electronic configuration?
- Can I explain why atomic radius decreases from Sodium to Chlorine across Period 3?
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Student Study Notes & Discussion
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