Overview
This lecture covers the structure, properties, history, and identification of atoms, including their components, interactions, and significance in chemistry and physics.
Atomic Structure and Properties
- Atoms are the smallest units of chemical elements, made of a nucleus (protons and neutrons) and electrons.
- The number of protons determines the element; isotopes are atoms with the same protons but different neutrons.
- Atoms are about 100 picometers wide, much smaller than visible light wavelengths.
- Most atomic mass is in the nucleus; protons are positive, neutrons neutral, electrons negative.
- Atom overall charge is neutral if protons = electrons; ions are charged atoms with unequal numbers.
- Electrons are attracted to the nucleus by electromagnetic force; protons/neutrons are bound by the nuclear force.
Historical Development of Atomic Theory
- Ancient philosophers proposed atoms as indivisible units.
- John Dalton (1800s) provided evidence for atoms with the law of multiple proportions.
- J.J. Thomson (1897) discovered electrons; atoms are divisible.
- Rutherford found the atom’s positive charge in a central nucleus.
- Bohr’s model introduced discrete energy orbits for electrons.
- Discovery of protons and neutrons completed the modern atomic model.
- Quantum mechanics (Heisenberg, Schrödinger) describes electron behavior as probability clouds, not fixed orbits.
Subatomic Particles
- Atoms are made of electrons (negative), protons (positive), and neutrons (neutral).
- Protons and neutrons are made of quarks, held by the strong force.
- Nucleons (protons and neutrons) cluster in a tiny nucleus at the atom’s center.
Electron Cloud and Energy Levels
- Electrons occupy quantized orbitals defined by probability distributions.
- Each orbital has a specific energy level; transitions between levels involve absorbing/emitting photons.
- Binding energy is much higher for nucleons than electrons.
- Valence electrons determine chemical bonding; atoms form molecules to fill outer shells.
Atomic Properties
- Elements are defined by proton number; isotopes vary in neutron number.
- Stable isotopes do not decay; radioactive isotopes undergo alpha, beta, or gamma decay.
- Atomic mass mainly comes from protons and neutrons (mass number).
- Atoms lack rigid boundaries—atomic radius is an average measure.
Chemical Bonding and the Periodic Table
- Valency is determined by the number of valence electrons.
- Elements in the periodic table group by similar valence electron counts; noble gases have full outer shells and are inert.
States of Matter and Identification
- Atoms exist in solids, liquids, gases, or plasmas depending on temperature and pressure.
- Scanning tunneling microscopes and mass spectrometry allow visualization and identification of atoms and isotopes.
Origin and Rarity
- Most atoms formed after the Big Bang; heavier atoms arose in stars or by cosmic processes.
- Some rare and superheavy elements exist as radioactive isotopes or are artificially created.
Key Terms & Definitions
- Atom — Smallest unit of a chemical element, consisting of nucleus and electrons.
- Isotope — Atoms with identical proton number but different neutron numbers.
- Ion — Atom with unequal protons and electrons, resulting in an electric charge.
- Nucleus — Dense atomic core containing protons and neutrons.
- Electron Cloud — Probabilistic region where electrons are likely found around the nucleus.
- Valence Electron — Electron in the atom’s outermost shell, involved in bonding.
Action Items / Next Steps
- Review the periodic table and practice identifying isotopes and ions.
- Solve problems involving chemical bonding and atomic mass calculations.