Structure and Properties of Matter

SCIENTIST CONTRIBUTIONS TO THE ATOMIC MODEL
1. Dalton
  1. Atomic Theory
  2. Billiard Ball Model:
2. Thomson
  1. Raisin Bun Model:
  2. Discovered the electron using the cathode ray tube
3. Rutherford
  1. Gold Foil Experiment
    1. He shot alpha particles at a thin sheet of gold where he expected the particles to just pass through but the results were that most particles passed through some were deflected and some bounced back.
    2. Based on the electrons passing through, he concluded that the atom is mostly empty space.
    3. Based on the electrons bouncing back, he concluded that the atom contains a small dense center and called it the nucleus.
    4. Based on the electrons deflecting, he concluded that the nucleus contains a poitively charged particle which he called a proton.
4. Maxwell
  1. Proposed that light is an electromagnetic wave.
  2. The electromagnetic spectrum:
5. Planck
  1. Studied Black Bodies
  2. Found that energy is found in packets, he called quantum of energy
  3. First indicaton of light acting as a particle instead of a wave.
6. Einstein
  1. Explained photoelectric effect, stating that "photons" (quanta) existed with descrete energy levels
  2. Only frequency affects current not brightness.
7. Bohr
  1. Studied spectral lines to determine electrons exsist in descrete energy levels.
  2. principal quantum number "n"
  3. Bohr Diagram:
8. deBroglie
  1. He found that electrons are waves.
9. Shrodinger
  1. Discovered that electrons exist in probability clouds which are where electrons probably are.
  2. Wave Equation
10. Heisenberg
  1. Heisenberg Law of Uncertainty
    1. You can never tell exactly what an electron is doing because by measuring it, we are changing it.
QUANTUM NUMBERS
1. Principal Quantum Number "n"
  1. describes which orbital or engery level
2. Secondary Quantum Number "l"
  1. Describes the shape
  2. Electron Configurations
    1. There are different values that have names : 0=sharp "s" 1=principal "p" 2=diffuse "d" 3=fundamental "f"
    3. subshells are written in order according to the Aufbau principle
    5. Trends in the Periodic Table
3. Magnetic Quantum Number "ml"
  1. describes the orientation
4. Fourth Quantum Number "ms"
  1. describes the spin
ELECTRON ORBITALS
1. Shape and Structure
  1. There is a large energy difference between individual energy levels
  2. There are smaller energy differences within energy levels
2. Orbitals
  1. Electron position is described using probabilities
  2. Orbital is the area within where the electron is most likely to be found.
  3. Within all orbitals there are different probabilities of where the electron will be.
3. Filling Orbital Diagrams
  1. Pauli's Exclusion Principle
    1. states that no two electrons may have the same set of quantum numbers
  2. Hund's Rule
    1. One electron must occupy each orbital of the same enrgy level before a second electron occupies an orbital
  3. Aufbau Principle
    1. "Building up" start at the lowest energy level and build up
    2. Exceptions: cronium, copper promote an electron from a lower orbital to fill all levels of a higher orbital
CHEMICAL BONDING
1. Lewis Theory
  1. Octet Rule
    1. Elements gain or lose electrons to attain an electron configuration of the nearest noble gas.
  2. Electron Pairing
    1. electrons are most stable when paired
  3. Sharing of Electrons
    1. Electrons can be shared between 2 non-metals in a covalent bond
  4. place electrons around atoms singly before pairing them
2. Valence Bond Theory
  1. Atomic orbitals or hybrid orbitals overlap to form new orbitals containing a pair of electrons of opposite spin.
  2. proposed by Linus Pauling
  3. Overlap arrangement results in a decrease in the energy of the atoms forming the bond
VSEPR THEORY
1. Way of visualizing molecular structure - around a central atom
2. Atoms in a molecule are bonded together by electron pairs called bonding pairs. Electrons not involved in bonding are called lone pairs.
3. Electron pairs are negatively charged and will get as far apart from each other as possible.
4. Lone pairs occupy more space than bonding electron pairs.
POLARITY OF MOLECULES
1. Depends on 2 Things:
  1. Polar Bonds
    1. Depends on electronegativity, which is the ability of an atom to attract electrons. It has polar bonds if the electronegativity difference is between 0.2 -1.7.
  2. Shape
    1. Needs to have a net dipole (if net dipole=0, it is non-polar). Assymetrical
2. To be a molecule it must have polar bonds and a net dipole .
3. Bond dipole: the electronegativity difference of two bonded atoms represented by an arrow pointing from the lower electronegativity to the higher electronegativity. The length of the arrow indicates the magnitude of the En difference.

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