Chem Exam: Unit 1

• • John Dalton: Matter is made of tiny definite particles called atoms.
  • Atoms cannot be created, destroyed, or divided.
  • Atoms of different elements combine to form compounds.
  • Atoms must combine in whole number ratios.
  • Circles model
• • J.J. Thompson: proved existence of electron.
  • Used a cathode ray and observed a negative pole repelled the ray.
  • Hypothesized that the ray from the cathode was made of negatively charged particles.
  • Plum Pudding Model.
• • James Chadwick: Discovered the neutron.
  • Worked with Rutherford, found electrons orbit the nucleus.
  • Electrons do not contribute to nuclei weight.
  • Weight of atom was not solely made of protons, therefore, a third subatomic particle with no charge must exist.
  • Model similar to Bohr, nucleus in middle, electrons orbiting.
• • Ernest Rutherford: Discovered nucleus of atom.
  • Fired alpha particles through gold foil and saw deflection, reflection and passing.
  • The atom was mostly empty space, with a positively charged center that contains mass.
  • Model included electrons orbiting a positively charged nucleus.
• Niels Bohr: Developed quantum model of hydrogen atom.

  Used an emission spectrum and proposed that electrons only move in specific orbits around the nucleus.

  Energy orbit levels increase as distance from the nucleus increases.

  When an electron gains energy, it moves away from the nucleus.
• Ernest Schrodinger: Electrons do not follow precise orbitals.

  Developed clouds showing probability of finding an electron in a certain part of the shell.

  Related to quantum, mechanical, and wave properties.
• Max Planck: Energy emitted by an object is not continuous/indefinite. Energy can only be emitted in definite amounts/quanta.
• • Louis de Broglie: Dual nature of electrons; particle and wave.
  • Electrons have wave-like motion and are restricted to circular orbits.
  • Electrons are only allowed certain wavelengths, frequencies, and energies.
  • Planck's quantum theory + Einstein's theory of relativity
• Heisenberg's Uncertainty Principle: Impossible to know exact position and speed of an electron at a given time.

  Can use photon to measure an instant, but once a particle is hit by a photon, it is know travelling at a different speed.
• James Maxwell: Light is an electromagnetic wave that can also act as a particle. Continuous wavelengths form a spectrum.

  E=hf

  h= Planck's constant

  f= frequency
• Photon: Unit of light energy
• • Isotope: Atoms of the same element with different numbers of neutrons.
  • Can be stable or radioactive.
  • Must contain same number of protons.
• Standard Atomic Notation

  A) Chemical Symbol

  B) Atomic Number (number of protons.

  C) Mass Number

  
• Quantum Mechanics: Branch of physics using math equations to describe wave properties of sub-microscopic particles like atoms.
• Orbitals: Regions of space around the nucleus, where there is high probability of finding an electron at any given energy.
• Quantum Numbers: Each electron has a unique set of four numbers, describing quantum mechanical properties of orbitals.
• • Principal (n): Energy level and relative size of the orbital.
  • Must be a positive integer
  • As n increases, electron energy in the orbital increases.
• • Secondary (l): Energy sublevels within each principal energy level.
  • Describes shape of atomic orbital
  • Whole number sublevels from 0-n-1
  • Letters used: s,p,d,f
• • Magnetic (Ml): Orientation of orbital in space relative to other orbitals.
  • Integers from -l to +l including zero
  • X,Y,Z, pairs
• • Spin (Ms): Direction of the spin of the electron (2 electrons per orbital.)
  • Value of 1/2 or -1/2 (Spins up or down)
• Emission Spectrum: Band of colours, as seen in a rainbow, produced by separation of components of light by different degrees of refraction according to wavelength.
• Continuous Spectrum: Consists of light at all wavelengths without interruption. Seen with hot/dense objects.
• Emission Spectrum: Emits light at specific wavelengths, forming a pattern (Hot gases).

  Spectra occurring from transition of excited state to ground/relaxed state.
• Absorption Spectrum: Looks like rainbow with black lines at specific wavelengths. Caused by cloud of gas blocking continuous light.

  Spectra occurring in the excited atom state, potential energy greater than ground state.
• • Energy Level Diagrams: Show an atoms orbital's in increasing energy. Visual representation of 4 quantum numbers for each electron.
  • Electrons closer to the nucleus have lower energy.
  • Orbitals with same n do not have equivalent energies.
  • Higher energy levels are closer together within bigger atoms.
• Pauli Exclusion Principle: No two electrons in an atom can have the same four quantum numbers.
• Aufbau Principle: Each electron is added to the lowest energy orbital available in an atom or ion.

  Exceptions to Aufbau: Chromium: 4s1, 3d5, Copper 4s1, 3d10

  Half filled/completely filled orbitals are more stable than unfilled subshells.
• Hund's Rule: One electron occupies each orbital at the same energy before a second electron can occupy the same orbital.
• Orbital Diagram: Same as energy level diagrams but easier to draw.
• Electron Configurations: Show number and arrangement of electrons in their orbitals.
• Condensed Electron Configurations: Core electrons will be same as last noble gas, but only valence electrons will be different.
• • Electronegativity: The ability of an atom to attract shared electrons in a covalent bond.
  • 0.0-0.4: Non polar covalent bond
  • 0.41-1.8: Polar covalent bond
  • 1.8+: Ionic bond
• • Metallic Bond: Metal atoms have less than half filled valence shells therefore, cannot combine to form a filled shell, or contribute electrons to form filled shells.
  • Valence electrons move freely among the atoms in an electron sea. The positively charged metallic ions are simultaneously attracted to multiple electrons.
  • Metals release electrons to achieve a stable octet, leading to free electrons.
• • Ionic Bond:
  • Oppositely charged ions attract one another, and form a crystalline structure/lattice.
  • Metal atoms lose electrons and become positive ions.
  • Non-Metal ions gain electrons and become negative ions.
  • Attracted to many oppositely charged ions.
• • Covalent Bond: Non-metal have more than half-filled valence shells and share electrons.
  • Covalent bonds are shown as a solid line.
  • Coordinate Covalent Bonds: One atom contributes both shared electrons, (donor + acceptor).
  • Expanded Octet: Valence energy levels of a central atom that has more than eight electrons.
  • Incomplete Octet: Halides of Boron and Beryllium have incomplete valence shells but are still stable molecules.
• • Valence Bond Theory: Covalent bond forms when orbitals of two atoms overlap.
  • A pair of electrons occupies the region of overlap.
  • Not all single bonds are derived directly from S and P orbitals, (hybridized orbitals as seen in methane, allowing for four identical bonds.)
• Hybrid Orbitals: Appropriate geometry to allow for bonding. Often happens when an atom is involved in more than one bond.

  Hybridization depneds on how many bonds need to be formed by the atom.
• Molecular Orbital Theory: When atomic orbitals overlap, they combine to form new orbitals called molecular orbitals.

  Molecular orbitals have new shapes and energy levels. Electrons are delocalized throughout the new orbital.
• Sigma Bonds: End to end overlap of two orbitals.

  Pi Bonds: Side to side overlap of two orbitals.
• • Single Bonds: One sigma bond.
  • Double Bonds: One sigma bond and one pi bond.
  • Triple Bonds: One sigma bond and two pi bonds.