Topic 1

• Atomic Number (Z)


  Number of protons in the nucleus
• Mass Number (A)


  Total number of protons and neutrons in the atom
• Number of neutrons


  A - Z
• Isotopes


  Atoms with the same number of protons, but different numbers of neutrons
• Relative isotopic mass


  Mass of one atom of an isotope compared to one twelfth of the mass of one atom of carbon-12
• Relative atomic mass


  Average mass of one atom compared to one twelfth of the mass of one atom of carbon-12
• Relative molecular mass


  Average mass of a molecule compared to one twelfth of the mass of one atom of carbon-12
• Isotopes have similar chemical properties because they have the same electronic structure. They may have slightly varying physical properties because they have different masses.
• Mass spectrometer


  • Can be used to determine all the isotopes present in a sample of an element and to therefore identify elements
• Mass spectra for Cl2 and Br2


  • Cl has two isotopes Cl35 (75%) and Cl37(25%), Br has two isotopes Br79 (50%) and Br81(50%)
• Sometimes two electrons may be removed from a particle forming a 2+ ion. 24Mg2+ with a 2+ charge would have a m/z of 12
• The peak with the largest m/z in a mass spectrum will be due to the complete molecule and will be equal to the Mr of the molecule. This peak is called the parent ion or molecular ion
• Uses of Mass spectrometers


  • Identifying elements on other planets
  • Drug testing in sport
  • Quality control in pharmaceutical industry
  • Radioactive dating
• First ionisation energy


  The energy required when one mole of gaseous atoms forms one mole of gaseous ions with a single positive charge
• Second ionisation energy


  The energy required when one mole of gaseous ions with a single positive charge forms one mole of gaseous ions with a double positive charge
• Factors affecting ionisation energy

  • The attraction of the nucleus
  • The distance of the electrons from the nucleus
  • Shielding of the attraction of the nucleus
• The patterns in successive ionisation energies for an element give us important information about the electronic structure for that element.
• When the first electron is removed a positive ion is formed


  The ion increases the attraction on the remaining electrons and so the energy required to remove the next electron is larger
• The first Ionisation energy of the elements shows a repeating pattern across a period, called periodicity.
• Helium has the largest first ionisation energy because its first electron is in the first shell closest to the nucleus and has no shielding effects from inner shells.
• First ionisation energies decrease down a group because the outer electrons are found in shells further from the nucleus and are more shielded, so the attraction of the nucleus becomes smaller.
• There is a general increase in first ionisation energy across a period because the number of protons increases making the effective attraction of the nucleus greater.
• Sodium has a much lower first ionisation energy than Neon because Sodium's outer electron is in a 3s shell further from the nucleus and is more shielded.
• There is a small drop in first ionisation energy from Magnesium to Aluminium because Aluminium is starting to fill a 3p sub-shell.
• n


  In the first shell closest to the nucleus and has no shielding effects from inner shells
• n has a bigger first ionisation energy than H


  As it has one more proton
• Phosphorus 1s2 2s2 2p6 3s2 3p3, Sulphur 1s2 2s2 2p6 3s2 3p4
• Two electrons of opposite spin in the same orbital
• An early model of the atom was the Bohr model (GCSE model) with electrons in spherical orbits. Early models of atomic structure predicted that atoms and ions with noble gas electron arrangements should be stable.
• The A-level model


  • Electrons are arranged on principle energy levels numbered 1,2,3,4.. with sub energy levels labelled s, p, d and f
  • s holds up to 2 electrons, p holds up to 6 electrons, d holds up to 10 electrons, f holds up to 14 electrons
• Orbitals represent the mathematical probabilities of finding an electron at any point within certain spatial distributions around the nucleus
• Each orbital has its own approximate, three dimensional shape. It is not possible to draw the shape of orbitals precisely.
• An atom fills up the sub shells in order of increasing energy (note 3d is higher in energy than 4s and so gets filled after the 4s)
• For oxygen 1s2 2s2 2p4
• Writing electronic structure using letters and numbers

  Number of main energy level, Name of type of sub-level, Number of electrons in sub-level
• Using spin diagrams


  An arrow is one electron, The arrows going in the opposite direction represents the different spins of the electrons in the orbital, Box represents one orbital
• When filling up sub levels with several orbitals, fill each orbital singly before starting to pair up the electrons
• Periodic table blocks


  s block element is one whose outer electron is filling a s-sub shell
• Electronic structure for ions

  When a positive ion is formed electrons are lost, When a negative ion is formed electrons are gained
• Elements are classified as s, p or d block, according to which orbitals the highest energy electrons are in