chem chapter 6

• Electromagnetic Radiation


  • Differences are due to different wavelength

  
• Colours of light


  • green
  • violet
  • blue
  • red
  • yellow
• Light


  Electromagnetic Radiation transmitted in the form of a wave

  
• Waves


  • Wavelength (λ, lambda): distance between 2 peaks or troughs (or corresponding points) in wave
  • Frequency (ν, nu): number of waves per second that pass a given point
  • Speed (c): speed of light is 2.9979 x 10^8 m/s (same for electromagnetic radiation in a vacuum)

  
• • 1 μm = 10^-6 m
  • 1 nm = 10^-9 m
  • 1 Ångström = 10^-10 m
• Calculating frequency from wavelength

  Use equation: c = λν

  
• Quantum Theory


  • Energy, like matter, is discontinuous and composed of tiny particles called photons
  • The energy of a quantum/photon of electromagnetic radiation is proportional to the frequency of the radiation: Energy = hν

  
• Calculating energy of a photon

  Use equation: Energy = hν

  
• Calculating energy of photons

  For photon with wavelength of 1150 nm

  2. For photon with frequency of 4.98 × 10^12 Hz

  
• Calculating energy of photons for photodissociation of O2

  Energy of one photon

  1. Energy of a mole of photons

  
• Continuous spectrum


  Contains all the wavelengths (& all energies) of light – sunlight

  
• Line (discrete) spectrum


  Contains only some of the wavelengths of light – atomic emission

  
• Ground State


  The lowest energy arrangement of electrons in an atom
• Excited State


  When an atom absorbs energy it can move an electron into a higher energy orbital

  
• Absorption


  The atom ABSORBS light energy and the electron is excited to a higher energy level

  
• Emission


  An excited atom is unstable, and the electron falls back from the higher energy level to the ground state RELEASING energy in the form of light = COLOUR

  
• Emission Spectra


  When an excited atom relaxes back to the ground state, it RELEASES energy in the form of light = COLOUR

  
• Rydberg Equation

  Used to calculate: 1. The energy of an electron in a particular energy level(n); 2. The energy change when an electron moves from one level to another

  
• Calculating energy of an electron in a particular energy level

  Use equation: En = -RH/n^2

  
• Calculating energy change when an electron moves from one level to another

  Use equation: ΔE = EFinal - EInitial = -RH(1/nf^2 - 1/ni^2)

  
• Calculating wavelength from energy change

  Use Rydberg equation to calculate ΔE

  2. Use Planck's equation: ΔE = hc/λ

  
• Atomic Structure


  Protons & Neutrons in the nucleus, Electrons around the nucleus in orbitals corresponding to specific energy levels
• Orbitals


  • s orbitals - spherical

  p orbitals - dumbbell shape

  d orbitals - 4 lobes

  f orbitals - complex
• Quantum Numbers


  4 quantum numbers that define the "address" of an electron: principal quantum number (n), angular momentum quantum number (l), magnetic quantum number (ml), and spin quantum number (ms)
• Energy Levels


  As n increases, the distance from the nucleus increases

  
• Principal Quantum Number (n)


  Describes the energy level and average distance of the electron from the nucleus

  
• Orbital


  s, p, d or f
• Schrödinger's Wave Equation


  Derived from
• First 3 quantum numbers

  • n - Principal Quantum Number (energy level)
  • ℓ - Angular Momentum Quantum Number (orbital type)
  • mℓ - Magnetic Quantum Number (orbital orientation)
• 4th quantum number - behaviour of the electron
• Energy Levels


  • Nucleus
  • As n increases, the distance from the nucleus increases

  
• Principal Quantum Number (n)


  • Values range from 1, 2, 3...
  • Describes the energy level
  • Describes the average distance of the electron from the nucleus
  • Orbitals of the same n belong to the same shell

  
• As the value of n increases

  • The size of the orbital increases
  • The distance of the electron from the nucleus increases
  • The electron is held less tightly - implications for chemical reactivity

  
• Angular Momentum Quantum Number (ℓ)


  • Also called Subsidiary or Azimuthal Quantum Number
  • Represents a subshell
  • Values range from 0, 1, 2...(n - 1) for each value of n
  • Relates to the orbital shape/type
  • Orbitals of same n & ℓ belong to same subshell
• Predicting values of ℓ from values of n
• Orbital types


  • s orbital (ℓ = 0)
  • p orbital (ℓ = 1)
  • d orbital (ℓ = 2)
  • f orbital (ℓ = 3)
• Within the subshell: s < p < d < f (increasing energy)
• Subshells in a simplified model of the atom
• Magnetic Quantum Number (mℓ)

  • Values range from -ℓ to +ℓ
  • Defines orientation of the orbital in space
  • Gives number of that type of orbital per shell: 2ℓ + 1 orbitals in subshell ℓ
• Predicting values of mℓ from values of n and ℓ