CHEMISTRY UNIT 1

• Sub-atomic particles

  • Electron
  • Protons
  • Neutrons
• Mass number


  The total number of protons and neutrons in the nucleus of an atom
• Isotopes


  Atoms of the same element that have the same number of protons but different numbers of neutrons
• Relative atomic mass/Isotopic mass

  The average mass of the isotopes of an element, taking into account their natural abundances
• The nucleus contains protons and neutrons
• There is a strong force operating between the protons and the neutrons, transferring of back and forth
• As long as there is roughly a balance between energy the nucleus stays together
• The neutron to proton ratio does not depart too far from 1, the nucleus is likely to be stable
• The proton to neutron ratio is close to 1 up to about Z = 20, then increases slowly up to about 1.6:1
• Stable nuclei are shown in the green area on the graph
• Outside the stable limits, the nuclei is unstable as it contains too much inherent energy
• All hyperenergetic systems change towards a more stable state by losing energy, this is called decay
• This shows the smallest whole number ratio of elements OR ions in a compound
• Energy quanta


  Energy can be gained or lost only in whole-number multiples of the quantity h
• Energy levels


  Discrete energy levels within the atom
• Max Planck hypothesized that energy was quantized in 1900
• Einstein envisioned light as small discrete particles of energy which he called photons
• Wavelength


  Distance between two consecutive peaks
• Frequency


  Number of waves per second that pass a given point in space
• The hydrogen line spectrum contains only a few discrete wavelengths, with only four wavelengths in the visible region
• Bohr developed a quantum model for the hydrogen atom in 1913, proposing the solar system model of the atom
• The principal quantum shells apart from the first shell are split into sub-shells
• Each sub-shell contains one or more atomic orbitals
• An orbital is a region of space where there is a high probability of finding an electron
• Each orbital can hold a maximum of two electrons
• The number of each type of orbital is s=1, p=3, d=5
• The Bohr Atom


  A small nucleus of protons and neutrons surrounded by circulating electrons, where each shell or energy level could hold a maximum number of electrons and the energy of levels became greater as they got further from the nucleus
• The Bohr theory


  • It couldn't explain certain aspects of chemistry
• Principal energy levels


  The energy levels were not equally spaced, with the energy gap between successive levels getting increasingly smaller as the levels got further from the nucleus
• Sub levels


  The main energy levels were split into sub levels, with level 1 split into 1 sub level, level 2 split into 2 sub levels, level 3 split into 3 sub levels, and level 4 split into 4 sub levels
• Heisenberg's Uncertainty Principle


  You cannot determine the position and momentum of an electron at the same time
• Aufbau Principle


  Electrons enter the lowest available energy level
• Pauli's Exclusion Principle


  No two electrons can have the same four quantum numbers, but two electrons can go in each orbital if they have opposite spin
• Hund's Rule of Maximum Multiplicity

  When in orbitals of equal energy, electrons will try to remain unpaired to reduce electrostatic repulsion
• Orbital


  A region in space where one is likely to find an electron, which can hold up to two electrons with opposite spin
• Orbitals


  • Have different shapes: s (spherical), p (dumb-bell), d (various), f (various)
• Orbitals are not filled in numerical order because the principal energy levels get closer together as you get further from the nucleus, resulting in overlap of sub levels
• Order of filling orbitals

  • 1s
  • 2s 2p
  • 3s 3p 3d
  • 4s 4p 4d 4f
  • 5s 5p 5d 5f
  • 6s 6p 6d
  • 7s 7p
• The filling order shows the 'building up' of the electronic structures of the first 36 elements in the periodic table
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