Periodicity

• what are electrons arranged by in the periodic table ?

  • by increasing atomic number
  • in periods - showing the reactivity trends in chemical and physical properties
  • in groups - similar chemical properties
• what do the elements in the same group in the periodic table have in common ?

  all the elements have the same number of electrons in the outer shell
• what do the elements in the same period in the periodic table have in common ?

  all elements have the same number of shells
• how are elements classified in the periodic table ?


  s, p, d, f blocks

  
• what is the first ionisation energy ?

  the energy required to remove 1 mole of electrons from 1 mole of gaseous atoms
• what does a first ionisation energy equation must have ?

  Mg(g) ➔ Mg+(g) + e− = +738 kJ mol−1

  • gaseous state symbol
  • + sign ➔ showing the element is positively charged due to its loss of electrons
  • e- ➔ showing the electron lost
  • + energy ➔ process requires energy hence it being exothermic
• what are the factors affecting ionisation energy ?


  • nuclear change
  • atomic radius
  • shielding
• what is the nuclear charge ?


  the electrostatic attraction between the positively charged protons in the nucleus and the negatively charged outer electrons
• what is the atomic radius?


  the distance between the positive nucleus and the outermost shell electrons
• what is shielding ?

  the negatively charged inner shell electrons repelling the outer shells electrons
• describe the trend in first ionisation energy across a period in terms of nuclear charge

  • higher nuclear charge - higher number of positively charged protons increase the strength of the electrostatic attraction between the positive nucleus and the outermost shell electrons, increasing ionisation energy
• describe the trend in first ionisation energy across a period in terms of atomic radius ?

  • lower atomic radius - there is a higher number of positively charged protons which increase the strength of the electrostatic attraction to the extra electrons which are added in the same outermost shell, allowing them to be pack closer to the nucleus, increasing the first ionisation energy
• describe the trend in first ionisation energy across a period in terms of shielding ?

  • constant shielding - there are no further added electron shells which allows shielding to be constant across a period, this has no effect on first ionisation energy
• describe the trend in first ionisation energy down a group in terms of atomic radius ?

  • higher atomic radius - there is a higher number of electrons which means more shells are added, moving the outermost electrons further from the nucleus, decreasing the electrostatic attraction between the positively charged electrons and the negatively charged outermost electors, making it easier to be removed and so resulting in a lower first ionisation energy
• describe the trend in first ionisation energy down a group in terms of nuclear charge ?

  • higher nuclear charge - higher number of positively charged protons increase the strength of the electrostatic attraction between the positive nucleus and the outermost shell electrons, increasing ionisation energy
• describe the trend in first ionisation energy down a group in terms of shielding ?

  • higher shielding - there is a higher number of electron shells due to the increases electrons which increases the repulsion from the inner shell electrons, this allows the electrons to be removed more easily, therefore decreasing first ionisation energy
• what is the overall first ionisation energy trend down a group ?

  lower ionisation energy - atomic radius and shielding effects are greater than the nuclear charge effect, leading to an overall decrease in ionisation energies as you move down a group
• what is the overall first ionisation energy trend across a period ?

  higher ionisation energy - the increasing nuclear charge effect outweighs the similar shielding across periods, so ionisation energies generally increase as you move across a period.
• what the successive ionisation energy ?

  the energy required to remove each successive electron until only the bare nucleus remains
• what does a successive ion isatin energy equation must have ?

  Mg+(g) ➔ Mg 2+(g) + e− = +1451 kJ mol−1

  • gaseous state symbol
  • 2+ sign ➔ showing the element is positively charged due to its loss of electrons
  • e- ➔ showing the electron lost
  • + energy ➔ process requires more energy hence it being exothermic
• why are successive energies always larger ?

  as successive electrons are removed from the same shell, the remaining electrons experience greater electrostatic attraction to the increasingly positive nucleus, this increased nuclear attraction requires more energy to remove the next electron from that shell, increasing the successive energies
• why are there large jumps in successive ionisation energy between shells ?

  • when reaching a new inner electron shell, there is a big increase in the ionisation energy needed to remove the first electron in that new shell.
  • this happens because the attraction to the nucleus is much greater for inner shell electrons closer to the nucleus.
• what is the name of a repeating pattern across a period ?

  periodicty
• why has Helium the largest first ionisation energy?

  the first electron is in the first shell closest to the nucleus and has no shielding effects from inner shells. He has a bigger first ionisation energy than H as it has one more proton
• 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 a 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 is 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 electron easier to remove.
• what is metallic bonding ?

  electrostatic attrition between the positive metal cations and the negative delocalised electrons
• what affects the strength of metallic bonding ?

  1. number of electrons / strength of nuclear attraction - the higher the number of protons, the higher the electrostatic attraction to the delocalised electrons
  2. number of delocalised electrons - the higher the number of electrons the higher the electrostatic attraction to the positive charge
  3. atomic radius - electrons are closer to the nucleus increasing the attraction
  4. ion size - the larger the positive size, the higher the attraction