Rucete ✏ AP Chemistry In a Nutshell
1. Structure of the Atom
In this chapter, you'll learn about historical discoveries related to atomic structure, including protons, neutrons, electrons, atomic models, and electron configurations.
Important Discoveries About the Atom
• Antoine Lavoisier (1774): Proposed the Law of Conservation of Matter, stating that matter cannot be created or destroyed during chemical reactions.
• Joseph Proust (1799): Proposed the Law of Constant Composition, stating each pure chemical compound always contains the same elements combined in a fixed ratio by mass.
• John Dalton (1803–1808): Formulated the Atomic Theory.
Dalton's Atomic Theory
• All matter is composed of tiny, indivisible particles called atoms.
• Atoms of the same element are identical in their properties.
• Chemical reactions involve rearrangements of atoms into new combinations in simple whole-number ratios.
Dalton also proposed the Law of Multiple Proportions.
Discovery of Atomic Particles
• Michael Faraday (1834): Demonstrated the electrical nature of chemical reactions.
• Sir William Crookes (1870s): Developed the cathode ray tube and initially believed cathode rays were negatively charged molecules.
• J. J. Thomson (1897): Determined cathode rays were electrons, measured the electron's charge-to-mass ratio (e/m = −1.76 × 10⁸ C/g), and proposed the "plum pudding" atomic model.
• Robert Millikan (1909): Calculated the electron's charge (−1.60 × 10⁻¹⁹ C) and mass (9.11 × 10⁻²⁸ g) using his oil drop experiment.
• Ernest Rutherford (1910): Performed the gold foil experiment, demonstrating that atoms have a small, dense, positively charged nucleus surrounded mostly by empty space.
• Rutherford (1919): Discovered the proton, a particle equal in magnitude but opposite in charge to electrons, with a mass of 1.67 × 10⁻²⁴ g.
• James Chadwick (1932): Discovered neutrons, particles with no charge and a mass nearly equal to protons.
Atomic Models
• Solid Particle Model (400 B.C.): Early concept viewing atoms as indivisible solid particles.
• Plum Pudding Model (Thomson, 1909): Atom as electrons dispersed throughout a sphere of positive charge.
• Nuclear Model (Rutherford, 1910): Small, dense nucleus with electrons orbiting in vast empty space.
• Solar System Model (Bohr, 1913): Electrons orbit nucleus at specific, quantized energy levels.
• Wave-Mechanical Model (Schrödinger, 1927): Electrons occupy probability regions (orbitals) rather than fixed orbits.
Subatomic Particles
• Electron (e⁻): Negatively charged particle, mass = 9.11 × 10⁻²⁸ g, located in electron clouds (orbitals).
• Proton (p⁺): Positively charged particle, mass = 1.67 × 10⁻²⁴ g, located in nucleus.
• Neutron (n⁰): Neutral particle, mass approximately equal to proton, located in nucleus.
Atomic Number, Mass Number, and Isotopes
• Atomic number (Z): Number of protons in an atom; identifies the element.
• Mass number (A): Total number of protons and neutrons in the atom’s nucleus (A = protons + neutrons).
• Isotopes: Atoms of the same element (same atomic number) with different numbers of neutrons (different mass numbers).
Example: Carbon isotopes—12C (6 protons, 6 neutrons), 13C (6 protons, 7 neutrons), 14C (6 protons, 8 neutrons).
Atomic Mass and the Atomic Mass Unit (amu)
• Atomic mass units are defined by carbon-12 atom: 1 amu = exactly 1/12 the mass of one carbon-12 atom.
• Atomic mass shown on periodic table is the weighted average of all naturally occurring isotopes.
Calculation of average atomic mass:
(% abundance of isotope 1 × mass of isotope 1) + (% abundance of isotope 2 × mass of isotope 2) + ...
Quantum Mechanical Model
• Electrons exist in regions of probability called orbitals.
• Orbitals are defined by quantum numbers (n, l, ml, ms).
• Electrons do not orbit the nucleus in fixed paths, but rather exist in "clouds" around the nucleus.
Quantum Numbers
• Principal quantum number (n): Energy level of electron (n = 1, 2, 3...).
• Angular momentum quantum number (l): Defines the shape of the orbital (l = 0 to n-1).
• Magnetic quantum number (ml): Orientation of orbital in space (-l to +l).
• Spin quantum number (ms): Spin orientation of electron (+½ or -½).
Electron Configuration
• Electron configurations describe the arrangement of electrons in an atom’s orbitals.
• Aufbau principle: electrons occupy lowest available energy levels first.
• Pauli exclusion principle: no two electrons can have the same four quantum numbers in an atom.
• Hund's rule: orbitals of equal energy are occupied singly by electrons before pairing.
Example: Oxygen (O, Z=8) → Electron configuration: 1s² 2s² 2p⁴