Module 2: Lab

• Democritus - Theorized the existence tiny indivisible particles called atomos
• Law of Conservation of Mass by Antoine Lavoisier (1789)

  • Mass is neither created nor destroyed in chemical reactions.
• Law of Definite Proportions by Joseph Proust (1797)

  • A given chemical compound always contains its component elements in a fixed ratio (by mass) and does not depend on its source and method of preparation.
• Law of Multiple Proportion by John Dalton (1800s)

  • If two elements form more than one compound, then the ratios of the masses of the second element which combine with a fixed mass of the first element will always be ratios of small whole numbers
• Modified Atomic Theory

  1. not all atoms have the same mass: isotopes
  2. atoms can transform from one element to another: radioactivity
  3. composition of solid compounds is variable: crystal structures
  4. atoms are divisible in some cases: nuclear fission
• Joseph John Thomson - proved that atoms of elements can be made to emit tiny negative particles. plum pudding model (electrons)
• Ernest Rutherford's gold foil experiment showed that the atom is mostly empty space with a tiny, dense, positively-charged nucleus.
• In May 1932 James Chadwick announced that the core also contained new neutral, massive particles in the nucleus of atoms. This conclusion arose from the disparity between an element's atomic number (protons = electrons) and its atomic mass (usually in excess of the mass of the known protons present).
• Niels Bohr - line spectra. Electrons orbit the nucleus that represents specific quantities of energy. The energies of the electrons in the atoms are quantized. Only certain electron orbits (energy levels) are allowed.
• Modern atom - Composed of protons, electrons, and neutrons. Mostly empty space
• Erwin Schrodinger - Modern atomic theory describes the electronic structure of the atom as the probability of finding electrons within certain regions of space (orbitals).
• 𝑛 – principal quantum number (distance from the nucleus) 𝑛=1,2,3,4,5,6,or7
• – angular momentum quantum number (shape of the volume of space occupied) s=0, p=1 d= 2 f=4
• ml - magnetic quantum number (orientation in space)

  z = negative

  x = zero

  y = positive
• ms - spin quantum number (orientation of electron spin)
• • Shell – electrons with the same value of n

  • Subshell – electrons with the same values of n and l

  • Orbital – electrons with the same values of n, l, and ml
• Pauli’s Exclusion Principle - States that no more than two (2) electrons can occupy a single orbital.
• 2. Aufbau’s Rules in distributing the electrons, the lowest energy levels ar
• Hund’s Rule specifies that within a given sub level, one electron is placed in each orbital before pairing could occur.