7 - periodicity

• Periodicity


  The trends in the physical and chemical properties of elements as you go across the periodic table.
• What are the 4 blocks?

  s, p, d, f
• 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.
• What factors affect the attraction between the nucleus and the outer electrons of an atom, and therefore, the ionisation energy.


  - Atomic radius

  - Nuclear charge

  - Electron shielding
• How does atomic radius affect ionisation energy?


  The greater the distance between the nucleus and the outer electrons the less nuclear attraction. The force of attraction falls off sharply with increasing distance, so the atomic radius has a large effect.
• How does nuclear charge affect ionisation energy?


  The more protons there are in the nucleus of an atom, the greater the attraction between the nucleus and the outer electrons. Atomic radius and shielding often outweigh this factor.
• How does electron shielding affect the ionisation energy?


  Electrons are negatively charged and so inner-shell electrons repel outer shell electrons. This repulsion, called the shielding effect, reduces the attraction between the nucleus and the outer electrons.
• Second ionisation energy


  It's the energy required to remove one electron, from each ion, in one mole of gaseous 1+ ions of an element, to form one mole of gaseous 2+ ions.
• What are the three key patterns in the first ionisation energy for the first 20 elements in the periodic table?


  - A general increase in the first ionisation energy across each period (H → He, Li → Ne, Na →Ar).

  
  

  - A sharp decrease in the first ionisation energy between the end of one period and the start of the next period (He → Li, Ne → Na, Ar → K).

  
  

  - The decrease in first ionisation energy down a group in the periodic table.

  
• What's the trend in first ionisation energies down a group?


  S - Shielding increases due to more inner shells

  A - attraction decreases

  N - Nuclear charge increases (outweighed)

  D - Atomic radius increases

  I - Ionisation energy decreases
• What's the trend in first ionisation energies across a period


  S - Shielding is similar: same shell

  A - Attraction increases

  N - Nuclear charge increases

  D - Atomic radius decreases

  I - Ionisation energy increases
• Why is there a fall in ionisation energy from beryllium to boron, despite them being next to each other in the same period (ionisation energy should increase)?


  The fall in the first ionisation energy from beryllium to boron marks the start of filling the 2p subshell.

  
  

  The 2p subshell in boron has a higher energy than the 2s subshell in beryllium. Therefore, in boron the 2p electron is easier to remove than one of the 2s electrons in beryllium. The first ionisation energy of boron is less than the first ionisation energy of beryllium.

  
• Why is there a fall in ionisation energy from nitrogen to oxygen, despite them being next to each other in the same period (ionisation energy should increase)?


  The fall in the first ionisation energy from nitrogen to oxygen marks the start it electron pairing in the p orbitals of the 2p subshell.

  
  

  - In nitrogen and oxygen the higheest energy electrons are in the 2p subshell.

  
  

  - In oxygen, the paired electrons in one of the 2p orbitals repel another, making it easier to remove an electron from and oxygen atom than a nitrogen atom.

  
  

  - Therefore the first ionisation energy of oxygen is less than the first ionisation energy of nitrogen.

  
• What happens in metallic bonding?


  - In a solid metal structure. each atom has donated its negative outer shell electron to a shared pool of electrons, which are delocalised throughout the whole structure.

  - The positive ions (cations) left behind consist of the nucleus and the inner shells of electrons.
• Metallic bonding


  The strong electrostatic force of attraction between cations and delocalised electrons.
• Why do metals have a rigid structure and shape?


  The fixed position of the cations.
• Three main properties of metals


  - Strong metallic bonds

  - High electrical conductivity

  - High melting and boiling points.
• How can the physical properties of metals be explained?


  Through the giant structure of the lattice and metallic bonding.
• What states do metals conduct electricity?


  Solid and liquid.
• How do metals conduct electricity?


  When a voltage is applied across a metal, the delocalised electrons can move through the structure, carrying charge.
• What does the melting points of metals depend on?


  The strength of metallic bonds holding together atoms in the giant metallic lattice.
• Why do metals have high melting and boiling points?


  For most metals, high temperatures are necessary to provide the large amount of energy needed to overcome the strong electrostatic forces of attraction between the cations and electrons.
• Are metals soluble?


  Metals are not soluble, they do not dissolve.
• What do many non-metallic elements exist as?


  Simple covalently bonded molecules. In a solid state, these molecules for a simple molecular lattice structure held together by weak intermolecular forces.
• What type of melting and boiling points do simple molecular lattice structures of non-metals have?


  Low melting and boiling points.
• Why do simple molecular lattice structures have low melting and boiling point?


  Weak intermolecular forces holding simple covalent molecules together are easily overcome.
• What type of covalent structures do boron, carbon and silicon form?


  A network of strong covalent bonds holds atoms together in a giant covalent lattice.
• What structure does diamond have?


  Tetrahedral
• Three main properties of substances with a giant covalent lattice structure


  - High melting and boiling points.

  - Insoluble in most solvents.

  - Non-conductors of electricity (other than graphene and graphite).
• Why do substances with giant covalent structures have high melting and boiling points?


  High temperatures are necessary to provide the large amount of energy needed to break the strong covalent bonds.
• Why are substances with giant covalent structures insoluble in most solvents?


  The covalent bonds holding together the atoms in the lattice are far too strong to be broken by interaction with solvents.
• Why doesn't diamond and silicon conduct electricity?


  All four outer shell electrons are involved in covalent bonding, so none are available for conducting electricity.
• Why does graphite and graphene conduct electricity?


  Only 3 of carbons 4 outer shell electrons are used in covalent bonding of the carbon atomd. The remaining electron is released into a pool of delocalised electrons.
• What type of structure is graphene and graphite?


  Giant covalent structures.
• What is graphene?


  Graphine is a single layer of graphite, composed of hexagonally arranged carbon atoms linked by strong covalent bonds.

  
  

  Graphene has the same electrical conductivity as copper, and is the thinnest and strongest material ever made.
• What is graphite?


  Graphite is composed of parallel layers of hexagonally arranged carbon atoms, like a stack of graphene layers.
• What's the trend in melting points across period 2?


  Melting point increases across the period from group 1 to group 4 (carbon). Then there is a sharp decrease in melting point between group 4 (carbon) and group 5 (nitrogen). The melting points are comparitively low from group 5 (nitrogen) to group 0 (neon).
• What's the trend in melting points across period 3?


  Melting point increases across the period from group 1 to group 4 (silicon). Then there is a sharp decrease in melting point between group 4 (silicon) and group 5 (phosphorus). The melting points are comparitively low from group 5 (phosphorus) to group 0 (argon).
• Why is there a sharp decrease in melting points at group 4 across periods 2 and 3?


  The sharp decrease in melting point marks a change from giant structures to simple molecular structures.