Chemistry

• Quantum number sets


  n = 2; l = 1; ml = -1
• Maximum number of electrons that can fill a subshell

  2l2 + 2
• Electron falls from n = 4 to n = 1

  A photon is emitted
• Electron moves from n = 2 to n = 6

  Greatest gain in energy
• Electron configuration illustrating Hund's rule


  Filling orbitals with unpaired electrons first
• Percent composition of hydrogen isotopes

  99.2% H, 0.8% D
• Electron configuration 1s22s22p63s23p64s13d5


  Represents Cr and Mn+
• Isotopes


  Atoms of the same element with varying mass numbers
• Isotopes of hydrogen


  • Protium (1 proton, 1 neutron)
  • Deuterium (1 proton, 2 neutrons)
  • Tritium (1 proton, 3 neutrons)
• Atomic number (Z)


  Number of protons
• Mass number (A)


  Number of protons + number of neutrons
• In a neutral atom, the number of protons equals the number of electrons
• Electrons are not included in mass calculations because they are much smaller
• Atomic mass


  Nearly equal to the mass number (sum of protons and neutrons)
• Atomic weight


  Weighted average of naturally occurring isotopes of an element
• There are no atoms with a mass exactly equal to the atomic weight of an element
• Mole


  A number of "things" (atoms, ions, molecules) equal to Avogadro's number (6.02 x 10^23)
• The atomic weight represents both the mass of the "average" atom of an element (in amu) and the mass of one mole of the element (in grams)
• Rutherford provided experimental evidence that an atom has a dense, positively charged nucleus
• Planck developed the first quantum theory, proposing that energy emitted as electromagnetic radiation comes in discrete bundles called quanta
• Planck's relation


  E = hf, where E is the energy of a quantum, h is Planck's constant, and f is the frequency of the radiation
• Bohr used the work of Rutherford and Planck to develop his model of the electronic structure of the hydrogen atom
• Bohr's postulate on Angular Momentum

  L = mvr = nh/2π, where L is angular momentum, m is mass, v is velocity, r is radius, n is the principal quantum number, and h is Planck's constant
• Bohr's equation for Electron Energy

  E = -RH/n^2, where E is the energy of the electron, RH is the Rydberg unit of energy, and n is the principal quantum number
• The energy of an electron increases (becomes less negative) the farther out from the nucleus it is located (increasing n)
• Ground state


  The state of lowest energy, in which all electrons are in the lowest energy levels
• Quantized energy


  Similar to the change in gravitational potential energy when ascending or descending a flight of stairs
• Staircase


  • Only allows certain discrete (quantized) changes of potential energy, unlike a ramp which allows an infinite number of steps
• Bohr's model of the hydrogen atom

  A nucleus with one proton forming a dense core, around which a single electron revolved in a defined pathway (orbit) at a discrete energy value
• Electron "jumping" from one orbit to a higher-energy one

  Requires an amount of energy exactly equal to the difference between one orbit and another
• Ground state (n = 1)


  The orbit with the smallest, lowest-energy radius
• Excited state


  When the electron is promoted to an orbit with a larger radius (higher energy)
• Bohr's Nobel Prize-winning model was reconsidered over the next two decades but remains an important conceptualization of atomic behavior
• We now know that electrons are not restricted to specific pathways, but tend to be localized in certain regions of space
• Atomic emission spectrum


  Composed of light at specified frequencies, where each line corresponds to a specific electron transition
• Each element can have its electrons excited to a different set of distinct energy levels, and thus possesses a unique atomic emission spectrum, which can be used as a fingerprint for the element
• Emissions from electrons dropping from an excited state to a ground state give rise to fluorescence, and the color of the emitted light is what we see
• Hydrogen emission line series

  • Lyman series (transitions from n ≥ 2 to n = 1)
  • Balmer series (transitions from n ≥ 3 to n = 2)
  • Paschen series (transitions from n ≥ 4 to n = 3)
• E = hc/λ

  The energy of the emitted photon corresponds to the difference in energy between the higher-energy initial state and the lower-energy final state
• The wavelengths of absorption correspond exactly to the wavelengths of emission because the difference in energy between levels remains unchanged