Periodicity

• Periodicity

  The repeating trend in properties of elements across the periodic table
• The periodic table is the most important organisational tool in chemistry
• The periodic table is the first point of reference for chemists everywhere and most chemistry laboratories have a periodic table prominently placed on the wall
• You do not need to memorise the periodic table (although some people have done so)
• It is nevertheless helpful to know where the common elements are positioned, and most students know the atomic numbers and relative atomic masses of common elements such as hydrogen, carbon, nitrogen, and oxygen
• It is essential that you are able to use the periodic table
• Periodic table

  • Arrangement, pattern, and shape reveal trends among the elements
  • Positions of the elements are linked to their physical and chemical properties
  • Essential for predicting the properties of elements and their compounds
• Atomic number

  Each successive element has atoms with one extra proton
• Groups

  Vertical columns where each element has atoms with the same number of outer-shell electrons and similar properties
• Periods

  Horizontal rows where the number of the period gives the number of the highest energy electron shell in an element's atoms
• Across each period, there is a repeating trend in properties of the elements, called periodicity
• The most obvious periodicity in properties is the trend from metals to non-metals
• Periodic trend in electron configuration

  1. Starts with an electron in a new highest energy shell

  2. Across Period 2, the 2s sub-shell fills with two electrons, followed by the 2p sub-shell with six electrons

  3. Across Period 3, the same pattern of filling is repeated for the 3s and 3p sub-shells

  4. Across Period 4, the 3d sub-shell is involved, but only the 4s and 4p sub-shells are occupied
• Elements in each group have atoms with the same number of electrons in their outer shell and the same number of electrons in each sub-shell
• Blocks

  Elements in the periodic table can be divided into blocks corresponding to their highest energy sub-shell (s, p, d, and f)
• Group names and numbers

  • Alkali metals (Group 1)
  • Alkaline earth metals (Group 2)
  • Transition elements (Groups 3-12)
  • Pnictogens (Group 15)
  • Chalcogens (Group 16)
  • Halogens (Group 17)
  • Noble gases (Group 18)
• Ionisation energy

  Measures how easily an atom loses electrons to form positive ions
• First ionisation energy

  The energy required to remove one electron from each atom in one mole of gaseous atoms of an element to form one mole of gaseous 1+ ions
• Factors affecting ionisation energy

  • Atomic radius (greater distance between nucleus and outer electrons means less nuclear attraction)
  • Nuclear charge (more protons in nucleus means greater attraction to outer electrons)
  • Electron shielding (inner-shell electrons repel outer-shell electrons, reducing nuclear attraction)
• Successive ionisation energies

  The energy required to remove an electron from each ion in one mole of gaseous (n-1)+ ions of an element to form one mole of gaseous n+ ions
• Successive ionisation energies provide important evidence for the different electron energy levels in an atom
• Atom

  Two protons attracting two electrons in the 1s sub-shell
• After the first electron is lost

  1. The single electron is pulled closer to the helium nucleus

  2. The nuclear attraction on the remaining electron increases

  3. More ionisation energy will be needed to remove this second electron
• Second ionisation energy

  The energy required to remove an electron from each ion in one mole of gaseous + ions to form one mole of gaseous 2+ ions
• Successive ionisation energies

  • They show large jumps when an electron is removed from a different shell, closer to the nucleus and with less shielding
• The large increase between the seventh and eighth ionisation energies of fluorine suggests that the eighth electron must be removed from a different shell, closer to the nucleus and with less shielding
• First shell

  Contains two electrons (n=1, closer to the nucleus)
• Second shell

  Contains seven electrons (n=2, the outer shell)
• Successive ionisation energies allow predictions to be made about the number of electrons in the outer shell, the group of the element in the periodic table, and the identity of an element
• The ionisation energies shown in Table 1 steadily increase but then there is a large increase between the third and fourth ionisation energies, showing that the fourth electron is being removed from an inner shell
• There are three electrons in the outer shell and the element must be in Group 13 (3), and since it is in Period 3, the element must be aluminium
• Trends in first ionisation energies

  • They provide important evidence for the existence of shells and sub-shells
• There is a general increase in first ionisation energy as each period increases (H-He, Li-Ne, Na-Ar)
• There is a sharp drop in first ionisation energy between the end of one period and the start of the next period (He-Li, Ne-Na, Ar-K)
• Trend in first ionisation energy down a group

  First ionisation energy decreases due to increased atomic radius and increased shielding
• The first ionisation energy graph for Period 2 shows three rises and two falls, linked to the filling of the s- and p-sub-shells
• The fall in first ionisation energy from beryllium to boron marks the start of filling the 2p sub-shell, as the 2p electron in boron is easier to remove than one of the 2s electrons in beryllium
• The fall in first ionisation energy from nitrogen to oxygen marks the start of electron pairing in the p-orbitals of the 2p sub-shell, as the paired electrons in oxygen repel each other making it easier to remove an electron
• The successive ionisation energies for an element in Period 3 are 578, 1817, 2745, 11577, 14842, 18379 kJ mol-1
• The element has six ionisation energies