Overview
This lecture covers the structure of atoms, the arrangement and roles of subatomic particles, the concept of isotopes, and calculations involving atomic mass and isotopic abundance.
Atomic Structure
- Atoms consist of a nucleus (protons and neutrons) surrounded by electrons in energy levels.
- Protons have a positive charge, neutrons are neutral, and electrons have a negative charge.
- Electrons orbit the nucleus in shells; the first shell holds up to 2 electrons, the second up to 8.
- Core electrons are in inner shells; valence electrons are in the outermost shell.
- The atomic number equals the number of protons and identifies the element.
- In a neutral atom, the number of electrons equals the number of protons.
Electron Configuration and the Periodic Table
- The number of valence electrons can be determined by the element's group number on the periodic table.
- Group 1A elements have 1 valence electron, Group 2A have 2, Group 4A (14) have 4, and Group 7A (17) have 7.
Atomic Symbols and Calculations
- Atomic symbols display the atomic number (protons) and atomic mass (protons + neutrons).
- Number of neutrons = mass number – atomic number.
- Number of electrons = atomic number – charge (for ions).
Ions
- Cations are positively charged ions (more protons than electrons).
- Anions are negatively charged ions (more electrons than protons).
- Example: Aluminum ion (Al³⁺) has 13 protons, 14 neutrons, and 10 electrons.
- Example: Phosphide ion (P³⁻) has 15 protons, 16 neutrons, and 18 electrons.
Forces in the Atom
- Opposite charges attract; like charges repel due to electric force.
- The strong nuclear force holds protons together in the nucleus despite repulsion.
Isotopes
- Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons.
- Isotopes have identical chemical properties but different nuclear properties.
- Example: Carbon-12 (6 protons, 6 neutrons); Carbon-13 (6 protons, 7 neutrons).
Average Atomic Mass and Isotopic Abundance
- Average atomic mass is a weighted average of all naturally occurring isotopes, based on percent abundance.
- Example: Carbon average atomic mass ≈ 12.01 based on 99% C-12 and 1% C-13.
- Example calculation for boron: (10 × 0.19) + (11 × 0.81) ≈ 10.81.
- For elements with two isotopes, if the average atomic mass is closer to one isotope's mass, that isotope is more abundant.
Key Terms & Definitions
- Atom — the smallest unit of an element, consisting of protons, neutrons, and electrons.
- Atomic number — number of protons in an atom.
- Mass number — sum of protons and neutrons in an atom.
- Isotope — atoms of the same element with different numbers of neutrons.
- Ion — an atom with a net electric charge due to loss or gain of electrons.
- Cation — positively charged ion.
- Anion — negatively charged ion.
- Valence electrons — electrons in the outermost shell.
- Strong nuclear force — force holding protons and neutrons together in the nucleus.
- Average atomic mass — weighted average mass of an element's isotopes.
Action Items / Next Steps
- Practice calculating protons, neutrons, and electrons for atoms and ions.
- Practice average atomic mass and percent abundance calculations for elements with multiple isotopes.