Ionisation energies

• What is the first ionisation energy

  The first ionisation energy is the enthalpy change when one mole of gaseous atoms forms one mole of gaseous ions with a single positive charge
• Give an equation that represents the first ionisation energy

  H(g) -> H+(g) + e-

  Always gaseous

  The equation for 1st ionisation energy Always follows the same pattern

  It does not matter if the atom does not normally form a +1 ion or is not gaseous
• What is the second ionisation energy

  The second ionisation energy is the enthalpy change when one mole of gaseous ions with a single positive charge forms one mole of gaseous ions with a double positive charge
• Give an equation to represent the second ionisation energy

  Ti+(g) -> Ti²+ (g) + e-
• What are the factors that effect ionisation energy

  There are three main factors

  • The attraction of the attraction -nuclear charge
  • The distance of the electrons from the nucleus - atom size
  • Shielding of the attraction of the nucleus
• Explain the attraction of the nucleus (nucleus charge)

  The more protons in the nucleus the greater the attraction
• Explain the distance of the electrons from the nucleus (atom size)

  The bigger the atom the further the outer electrons are from the nucleus and the weaker the attraction to the nucleus
• Explain shielding of the attraction of the nucleus

  An electrons in an outer shell is repelled by electrons in complete inner shells weakening the attraction of the nucleus
• Describe successive ionisation energies

  The patterns in successive ionisation energies for an element give us important information about the electronic structure for that element
• Why are successive ionisation energies always larger

  The second ionisation energy of an element is always bigger than the first ionisation energy

  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
• How are ionisation energies linked to electronic structure

  The fith electron is in the inner shell closer to the nucleus and therfore attracted much more strongly by the nucleus than the fourth electron. It also does not have any shielding by inner complete shells of an electron here is an example. Here there is a big jump between the 2nd and 3rd ionisation energies which means that this element must be in group 2 of the periodic table as the 3rd electron is removed from an electron shell closer to the nucleus with less shielding and so has a larger ionisation energy

  
• The shape of the graph for periods two and three is similar. A repeating pattern across a period is called periodicity.

  The pattern in the first ionisation energy gives us useful information about electronic structure

  
• Why has helium the largest first ionisation energy

  It's electron is in the first shell closest to the nucleus and had no shielding effects from inner shells. He has a bigger first ionisation energy than H as it has one more proton
• Why do first ionisation energies decrease down a group

  As one goes down a group the outer electrons are found in shells further from the nucleus and are more shielded so the attraction of the nucleus becomes smaller
• Why is there a general increase in first ionisation energy across a period

  As one goes across a period the electrons are being added to the same shell which has the same distance from the nucleus and the same shielding effect. The number of protons increases However making the effective attraction of the nucleus greater
• Why has Na a much lower first ionisation energy than Neon

  This is because Na will have its outer electron in a 3s shell further from the nucleus and is more shielded. So Na's outer electron is easier to remove and has a lower ionisation energy
• Why is there a small drop from Mg to Al

  Al is starting to fill 3p sub shell whereas Mg has its outer electrons in the 3s sub shell. The electrons in the 3p subshell are slightly easier to remove because the 3p electrons are higher in energy and are also slightly shielded by the 3s electrons
• Why I there a small drop from P to S

  With sulphur there are 4 electrons in the 3p sub shell and the 4th is starting to doubly fill the first 3p orbital.

  When the second electron is added to a 3p orbital there is a slight repulsion between the two negatively charged electrons which makes the second election easier to remove

  Phosphorus - 1s²,2s²,2p⁶,3s²,3p³

  Sulphur- 1s²,2s²,2p⁶,3s²,3p⁴