Periodicity

• Elements are arranged in order of increasing atomic mass.
• Why are potassium and argon not arranged in order of increasing atomic number?

  Because they have similar properties to the elements in the groups they are placed in.
• Definition of periodicity.

  the trend in properties that is repeated across a period
• Describe the trend across a period using NASA.

  • The nuclear charge increases because the number of protons increases

  • Atomic radius decreases

  • There is no increase of shielding because electrons are added to the same shell or in the distance from the nucleus

  • The nuclear attraction between the nucleus and the outer electrons therefore increases
• Describe the trend down a group using NASA.


  • Nuclear charge increases because the number of protons increases
  • Atomic radius increases because the number of shells increases
  • Shielding increases because there is an increase in the number of inner shell electrons
  • These two factors outweigh the increase in nuclear charge so the nuclear attraction between the nucleus and the outer electrons decreases.
• Definition of first ionisation energy.

  The energy needed to remove one electron from each atom in one mole of gaseous atoms to form one mole of gaseous 1+ ions.
• Write the I.E for potassium.

  K(g) ——> K+(g) + e-
• Definition of successive ionisation.

  A measure of the energy needed to remove each electron in turn.
• An element has the following first four ionisation energies: 590, 1150, 5940,6480 kJ mol-1. Suggest which Group this element belongs to and explain why.

  Group 2, as there is a large increase in ionisation energy from the 2nd to 3rd ionisation energy. Therefore there are 2 electrons in the outer shell and the 3rd is being removed from a new shell closer to the nucleus.
• How does an increase in atomic radius affect I.E.

  Increasing the atomic radius increases the distance between the nucleus and the outermost electrons. Nuclear attraction decreases, therefore I.E decreases.
• How does an increase in nuclear charge affect I.E.

  Increasing the number of protons in the nucleus increases the nuclear atraction, therefore I.E. increases.
• How does an increase in shielding affect I.E.

  Electrons in filled inner shells repel electrons in the outer shell and decrease nuclear charge. If the number of inner shell electrons increases then shielding increase. Nuclear attraction decreases and therefore I.E. decreases.
• What is the trend in atomic radius across a period?

  • Across a period the nuclear charge increases as the number of protons increases
  • The shielding is constant because electrons are added to the same shell
  • Nuclear attraction between the nucleus and the outer electrons increases
  • Therefore the atomic radius decreases across a period
• What is the trend in atomic radius down a group?

  • The number of shells increases so the outer electrons are further away from the nucleus
  • Shielding increases as the number of inner shell electrons increases
  • These two factors outweigh the increase in nuclear charge
  • Nuclear attraction decreases
  • Therefore the atomic radius increases down a group
• What is the trend of 1st I.E across a period. NASA

  Ionisation energy increases.

  Nuclear Charge= increases as the number of protons increases

  Atomic Radius= decreases as the number of shells decreases

  Shielding= is constant because electrons are added to the same shell

  Nuclear attraction=increases
• What is the trend in 1st I.E down a group?

  Ionisation energy decreases.

  Nuclear Charge= increases as the number of protons increases

  Atomic Radius= increases as the number of shells increases

  Shielding= increases because there is an increase of inner shell electrons

  Nuclear attraction= decreases as these two factors outweigh the increase in nuclear charge
• What are the 4 combinations of structure and bonding that can occur in elements.

  1. Giant covalent
  2. Giant ionic
  3. Giant metallic
  4. Simple covalent
• Between Na and P, which element would have a higher melting point. Explain using structure and bonding.

  Na – giant metallic, strong metallic bonds = higher melting point because more energy needed to break the strong metallic bonds

  P4 – simple covalent, weak intermolecular (London) forces between the molecules
• Between Na and Al, which element has a higher point. Explain in terms of structure and bonding.

  Na – giant metallic, +1 charge on ion

  Al – giant metallic, +3 charge on ion

  stronger attraction between +3 ion and delocalised electrons therefore more energy needed to break the metallic bonds in Al
• Why do metals have a high melting point ?

  They have metallic bonding and form giant metallic lattices. The electrostatic forces of attraction between positive metals and negative delocalised electrons are strong, therefore the melting points are high.
• Describe and explain the patterns in successive ionisation energies.

  Ionisation energy increases. As each electron is removed from an atom, the remaining ion becomes more positively charged. The force attracting the outer electrons is increased. It therefore becomes increasingly difficult and more energy is required to remove the next electron each time.
• What do the large rises in the successive ionisation energies indicate ?

  They indicate the emptying of a shell and that the next electron is removed from a new shell closer to the nucleus.
• Compare the electronic configuration of Be and B and suggest a reason for the slight dip in first ionisation energy.

  Be: 1s2, 2s2

  B: 1s2, 2s2, 2p1

  The first electron removed from B is in 2p and for Be it is in 2s. 2p is higher in energy than 2s. Therefore it requires less energy to remove.
• Compare the electronic configurations of N and O and suggest a reason for the first dip in ionisation energy.

  The first electron removed from O is from a filled orbital and the one removed from N is from a partially filled orbital. Oxygen's electron experiences more repulsion and therefore it is easier to remove.
• What are the steps to explain the difference in melting points of two elements ?

  1. Identify the structure and bonding
  2. Identify the forces that break
  3. Compare the strengths
  4. Explain the effect on the melting point