Chapter 1: Atomic Structure

• Atomic number vs Mass number: Atomic number is number of protons. Mass number is number of protons + neutrons (atoms that make up the mass of an element).
• Atoms with same atomic number but different mass numbers are ISOTOPES.
• Atomic mass vs Atomic Weight: Atomic mass is nearly equal to its mass number. Atomic weight is the weighted average of the different isotopes.
• Rutherford: provided experimental evidence that an atom has a dense, positively charged nucleus that accounts for a small part of the atom's volume and the rest is basically empty space
• Planck: developed first quantum theory — proposing that energy emitted as electromagnetic radiation from matter comes in discrete bundles called quanta
  • E = hf (h = 6.626 x 10^-34 J*s)
• Bohr: took Rutherford and Planck's ideas to develop a model that shows that electrons don't move randomly, but in fixed orbits set distances away from the nucleus
  • he predicted possible values for angular momentum of an electron orbiting a H nucleus
    • L = nh / 2π → angular momentum is proportional to n (principal quantum number)
    • E = -Rh/n^2 → energy is proportional to n^2 (principal quantum number)
• Electromagnetic energy of a photon: E = hc/λ (c = 3.00 x 10^8 m/s)
• As electrons go from a lower level to a higher level, they get AHED: Absorb energy, Higher potential, Excited, and Distant (from nucleus)
• as a general statement, the principal quantum number is inversely proportional to the wavelength (λ)
• Heisenberg uncertainty principle: it is impossible to determine with perfect accuracy the momentum and position of an electron
• Pauli exclusion principle: no two electrons in a given atom can possess the same set of four quantum numbers
• Aufbau principle: electrons fill from lower to higher energy subshells and each subshell will fill completely before electrons begin to enter the next one
• Hund's rule: orbitals are filled such that there are a maximum number of half-filled orbitals with parallel spins
  • Chromium and Copper (and other elements in its groups) do not follow this rule because a full d subshell outweighs moving an electron from the p orbital
• Paramagnetic (unpaired electrons that attract in a field) and Diamagnetic (paired electrons that repel in a field)
• Quantum Numbers:
  • n = principal quantum number → describes the energy level of a shell
  • l = azimuthal quantum number → describes the subshells (s, p, d, f)
  • m(l) = magnetic quantum number → describes the orbital an electron is found in
  • m(s) = spin quantum number → always 1/2 or -1/2 (different spin orientations)

Chapter 2: The Periodic Table

• Periods are rows and groups are columns
• Metals: shiny, malleable, ductile, low effective nuclear charge, low electronegativity, large atomic radius, good conductors.
  • ex: copper, nickel, silver, gold, palladium, platinum, alkali and alkaline earth metals (Group 1 and 2)