Chem Exam IV

• Spectroscopy is the study of interactions between electromagnetic radiation (“light”) and matter.
• Light is also subject to particle-wave duality. It can be thought of as discrete, massless particles called photons. Photons can be thought of as “packets” of energy.
• The classifications of light from highest energy to lowest is

  1. Gamma Rays
  2. X-Rays
  3. Ultraviolet Rays
  4. Visible Light
  5. Infrared
  6. Microwaves
  7. Radio Waves
• In absorption, a photon collides with an electron in an orbital called the ground state. Because of this, the electron gains potential energy- it will move up to a higher orbital called the excited state.
• Quantized means occurs in steps.
• In emission, an electron transitions from an excited state to a ground state. (This process is also quantized)
• Besides from absorption and emission, scattering is the other main type of spectroscopy. Scattering occurs when a photon collides with an electron, but no quantum transitions occur. The direction/energies of scattered photons can be used to determine information about a chemical sample.
• How different regions of the EMS excite different parts of an atom.

  Gamma rays = nucleus

  X-rays = inner electrons

  UV, Visible and Infrared= bonding electrons

  Microwaves = spin of electrons

  Radio waves = spin of nuclei
• The three main types of spectroscopy is Absorption, Emission and Scattering.
• Zeff is the amount of nuclear charge that functionally attracts an electron. How attractive is an electron to atoms positively charged nucleus.
• All of these INCREASE as you go up and RIGHT across a periodic table.

  • Zeff
  • Ionization Energy - the amount of energy required to remove an electron from an atom, creating a cation.
  • Electronegativity - an atom’s tendency to attract electrons within a covalent bond.
• Atom size increases you go LEFT and DOWN the periodic table.

  Larger atoms remain at the bottom.
• As you move down a periodic table, electron shielding increasing which means that valence electrons are less likely to be attracted to the nucleus. (Super long configurations)
• If the electronegativity between two elements is significantly different, their bond is said to be polar. This means that there is a partial charge on each, even if the bond is not ionic. The more electronegative atom takes on a partial charge. If the difference in electronegativity is too great, atoms will not even form covalent bonds. Instead, they will form ionic substances.
• Formal charge is a model used to approximate how charge is distributed between atoms in a Lewis structure.
• The number of bonds between two atoms can be represented as bond order. For example, a triple bond has a bond order of 3.
• When comparing the same bond order, more polar bonds correspond to a stronger/shorter bond.
• hydrogen and halogens will ONLY form one bond
• For larger molecules, one must choose a central atom. It should be the least numerous and least electronegative (hydrogen and halogens are POOR choices!)
• Sometimes, a structure can have more than one valid Lewis structure. In these cases, the molecule’s structure is thought to exist as an average between these structures (resonance). In resonance, electrons are said to be delocalized. This means that electrons are perpetually moving around and not stuck (localized) at one atom. This is how averaging happens.This average is weighted, with the most stable structure contributing most to the resonance hybrid.
• Remember that hydrocarbons mostly contain hydrogen and carbon. Also remember that carbon atoms will tend to make 4 bonds while hydrogen atoms will tend to make 1 bond. Stick hydrogen atoms on each carbon atom, and you will be able to figure out which type of bond must be between each carbon atom.
• Valence-shell electron repulsion (VSEPR) theory is built on the idea that electron orbitals will repel from one another.
• Electronic geometry (i.e., how orbitals are positioned around it) in a molecule. This is based on the number of electron domains (connected atoms or lone pairs) on an atom.

  A) Linear

  B) Trigonal Planar

  C) Tetrahedral

  D) Trigonal Bipyramidal

  E) Octahedral

  
• For every lone pair in your Lewis structure, you can “cut off” an orbital. This resulting structure is the molecular geometry. It does not matter where the “cut” orbital is, except for the trigonal bipyramidal and octahedral geometries. These must come off from the sides.
• c = speed of light (3.00 x 108 meters/seconds)
• h = Planck’s Constant (6.626x 10^-34 Joules/second)
• E = energy (Joules)
• λ = wavelength (meters)
• ν = frequency (s-1 or Hertz)
• λ ( wavelength) = C/V
• ν ( frequency) = E/H