Chemistry unit 1

• Dalton's atomic theory

  • All atoms of the same element are exactly alike
  • Atoms cannot be broken down any further
  • Atoms of different elements have different masses
  • Atoms combine to form more complex structures (compounds)
• Changes to Dalton's atomic theory


  • 1897 - Thomson discovered the electron and suggested the 'plum-pudding' model
  • 1909 - Rutherford suggested the planetary model with a positively-charged nucleus
  • 1913 - Bohr suggested electrons orbit the nucleus at certain distances
  • 1932 - Chadwick discovered the neutron
  • 1926 onwards - Schrödinger, Born and Heisenberg described electrons in terms of probability and atomic orbitals
• Sub-atomic particles


  • Proton - Relative mass 1, Relative charge +1
  • Neutron - Relative mass 1, Relative charge 0
  • Electron - Relative mass 1/1836, Relative charge -1
• Effect of electric field on sub-atomic particles


  • Charged particles move towards plates with opposite charge and away from plates with same charge
  • Electrons are deflected more than protons due to their smaller mass
• Sub-atomic particles and magnetic fields
• Atomic number (Z)


  The number of protons in the nucleus of an atom
• Mass number (A)


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


  Atoms with the same atomic number but different mass numbers
• Relative atomic mass (Ar)


  The weighted average mass of naturally occurring atoms of an element on a scale where an atom of 12C has a mass of exactly 12 units
• Relative isotopic mass


  The mass of a particular isotope of an element on a scale where an atom of 12C has a mass of exactly 12 units
• Calculating accurate relative atomic masses


  1. Take into account the proportion of each isotope (relative isotopic abundance)

  2. Calculate the weighted average mass
• Types of radioactive emissions


  • Alpha (α)
  • Beta (β)
  • Gamma (γ)
• Alpha (α) decay


  Helium nuclei (positively charged particles) are emitted
• Beta (β) decay


  Electrons (produced by nuclear changes) are emitted
• Gamma (γ) decay


  Very high frequency electromagnetic radiation is emitted
• Equations can be written for each type of radioactive decay
• Uses of radioisotopes


  • Tracers for searching for faults and studying organs
  • Radiotherapy for cancer treatment
  • Dating objects using 14C
  • Smoke detectors using 241Am
  • Generating power using 235U
• Energy quanta


  The smallest fixed amount of energy required for a change
• Energy levels


  Certain fixed values of energy that electrons in atoms can have
• The atom is most stable when the electrons are in the lowest energy levels possible (ground state)
• When an electron absorbs a quantum of radiation it can move up to a higher energy level (excited state)
• When an excited electron falls to a lower energy level, a quantum of radiation is given out
• Relationship between energy difference and radiation frequency

  ΔE = hν
• Principal quantum levels/energy levels


  Numbered n = 1, n = 2, n = 3, etc. going further away from the nucleus
• Atomic emission spectra


  • Made up of separate lines
  • Lines converge (get closer to each other) as their frequency increases
• Lyman series


  Electrons fall back to the n = 1 energy level (seen in ultraviolet region)
• Balmer series


  Electrons fall back to the n = 2 energy level (seen in visible region)
• The energy associated with each line in the emission spectrum is found using ΔE = hν
• Electron


  Can jump from one energy level to another by absorbing or emitting a fixed amount (quantum) of energy
• Lines in the hydrogen emission spectrum get closer together
• Hydrogen emission spectrum


  • Each line is a result of electrons moving from a higher to a lower energy level
  • Includes Lyman series (seen in ultraviolet region) and Balmer series (seen in visible region)
• ΔE


  The energy associated with each line, found using the relationship ΔE = hν
• The greater the frequency (the smaller the wavelength), the more energy is released
• Convergence limit


  The point where the lines eventually come together, representing an electron falling from the highest possible energy level
• If the electron has more energy than the convergence limit, it becomes free from the pull of the nucleus of the atom and the atom is converted to an ion
• Bohr model of the hydrogen atom


  • The energy levels get closer together towards the outside of the atom
  • A quantum of energy moves the hydrogen electron in the n = 1 level (ground state) to the n = 2 level
  • When the electron loses energy, it will fall down to the lower energy levels, emitting radiation of characteristic frequencies
• Electrons occupy specific energy levels (quantum levels) in the atom
• When electrons gain specific quanta of energy they move from lower to higher energy levels. They become 'excited'
• When excited, electrons lose energy, they fall back to lower energy levels emitting radiation of characteristic frequency. This is the origin of the line emission spectrum
• ΔE


  The energy difference between any two energy levels, given by ΔE = hν