7 - Periodicity

• periodicity - a repeating trend in properties of elements across each period
• 18th Century - Antoine Lavoisier listed substances he thought were elements (many of these substances were compounds)
• 1869 - Dimitri Mendeleev

  • arranged the elements in order of increasing atomic mass
  • he left gaps for elements not yet discovered
  • he put elements with similar properties into the same vertical group
  • if group properties did not fit he swapped elements around
• 1864 - John Newlands

  • ordered elements by atomic mass
  • every eighth element was similar (Law of Octaves)
  • only worked for the first 15 known elements
• 1820s - Johann Döbereiner

  • 3 elements with similar properties (triads)
  • atomic mass of the middle element was roughly halfway
  • not many elements fell neatly into this pattern
• now the elements are arranged:

  • in order of increasing atomic (proton) number.
  • in vertical columns called groups
  • in horizontal rows called periods
• repeating trends could include:

  • electron configuration
  • ionisation energy
  • structure
  • melting points
• the number of the period gives the number of the highest energy electron shell of an element's atoms
• group trend in electron configuration:

  • elements in each group have the same number of electrons in their outer shell
  • elements in each group also have the same number of electrons in each sub-shell
  • this is why elements in the same group have similar properties
• label the diagram

  A) s-block

  B) d-block

  C) p-block

  D) f-block

  
• fill in the table

  A) alkali metals

  B) alkaline earth metals

  C) transition elements

  D) pnictogens

  E) chalcogens

  F) halogens

  G) noble gases

  
• ionisation energy - the energy required to remove one electron from each atom in one mole of gaseous atoms of an element
• units of ionisation energy = kJ mol-1
• first ionisation energy

  X(g) -> X +(g) + e -
• factors affecting ionisation energy:

  • atomic radius
  • nuclear charge
  • inner shell sheilding
• atomic radius

  the greater the distance between the nucleus and outer shell electrons, the weaker the nuclear attraction between them
• nuclear charge

  the more protons there are in the nucleus of an atom, the greater the nuclear attraction between the nucleus and the outer shell electrons
• inner shell shielding

  increased shielding reduces the attraction between the nucleus and the outer electrons
• what is the shielding effect?

  the inner shell electrons repelling the outer shell electrons
• group trend in first ionisation energy

  • first ionisation energy decreases down a group because
  • atomic radius increases
  • more inner shells, so increased shielding
  • nuclear attraction decreases
• elements that are solids at room temperature

  • all other elements
• elements that are liquids at room temperature

  • mercury
  • bromine
• elements that are gases at room temperature

  • hydrogen
  • nitrogen
  • oxygen
  • fluorine
  • chlorine
  • all group 0 elements
• periodic trend in first ionisation energy

  • first ionisation energy increases across a period because
  • nuclear charge increases
  • same number of inner shells therefore shielding is the same
  • atomic radius decreases
  • nuclear attraction increases
• periodic trend in first ionisation energy

  sub-shell trends

  • rise from lithium to berillium
  • fall to boron
  • rise to carbon and hydrogen
  • fall to oxygen
  • rise to fluorine and neon
• complete the graph

  type done when the line is complete

  A) first ionisation energy

  B) atomic number

  C) done

  
• the fall in first IE from Beryllium to Boron

  • the 2p sub-shell has a higher energy that the 2s sub-shell
  • its is easier to remove the outer 2p electron in Boron, than the 2s electron in Beryllium
  • the first ionisation energy of Boron is less than the first ionisation energy of Beryllium
• fall in the first IE from Nitrogen to Oxygen

  • in Nitrogen there is one electron in each 2p orbital
  • this means there is equal repulsion between the electrons
  • in Oxygen there are two electrons in one 2p orbital
  • these paired electrons repel
  • its is easier to remove a spin-paired 2p electron from Oxygen than an electron from Nitrogen
  • the first IE of Oxygen is lower than Nitrogen
• successive ionisation energies

  • outer shell electrons are removed one at a time
  • first ionisation energy: He (g) -> He+ (g) + e-
  • second ionisation energy: He+ (g) -> He2+ (g) + e-
  • the second ionisation energy of helium is greater than the first
  • this is because the second electron is removed from a positive ion and the remaining electrons will have been pulled closer to the nucleus
  • nuclear attraction on remaining electron increases
  • more energy is needed to remove this second electron
• successive ionisation energies provide evidence for the existence of shells
• the value of successive ionisation energy increases with ionisation number

  • as each electron is lost, there is the same number of protons attracting fewer electrons
  • the electrons are drawn slightly closer to the nucleus, increasing the attraction
• the number of ionisation energy is the same as the charge on the ion produced
• the second ionisation of magnesium is the energy required to remove one electron from each ion in 1 mole of gaseous 1+ ions to form 1 mole of gaseous 2+ ions
• predictions

  fluorine atom has 7 electrons in its outer shell - group 7

  fluorine atom has 2 electron shells - period 2