CHEMISTRY M3

• What is the period table arranged by?

  Atomic number

  
  

  periods = across

  
  

  groups = down
• periodic trends in electron configurations across periods 2 and 3


  each period starts with an electron in a new subshell

  
  

  *however across period 3 the 4s subshell fills first

  
• What is the first ionisation energy?


  The energy required to remove one electron from each atom in 1 mole of gaseous atoms of an element to form 1 mole of gaseous 1+ ions.
• What is the equation for the first ionisation energy of sodium?

  Na(g) → Na+(g) + e-
• What 3 factors affect 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 the nuclear attraction to the electron
• How does nuclear charge affect ionisation energy?


  The more protons in the nucleus the greater the attraction to the electrons.
• How does electron shielding affect ionisation energy?


  Electrons are negatively charged.

  
  

  So inner shell electrons repel outer shell electrons.

  
  

  This reduces the attraction between the nucleus and outer electrons.
• What is the equation for the 2nd ionisation of iron?

  Fe+(g) → Fe2+(g) + e-
• Why is second ionisation energy always more that the first?

  Once 1 electrons goes, the remaining electrons are pulled closer to the nucleus. This reduces the atomic radius and increases nuclear attraction to the remaining electrons.
• What does a large increase in ionisation energy mean?

  The electron removed where the big jump was, is from a different shell. A shell that is closer to the nucleus with less shielding

  
  

  *table doesn't show every single IE for each element, but it shows the big jumps that reflect on changing shell

  
• What predictions can be made from successive ionisation energies?

  Number of electrons in outer shell.

  Group of the element.

  Identity of the element.
• What happens to 1st ionisation energy down a group?

  Decreases

  
• What happens to 1st ionisation energy generally across a period?

  Increases

  
• Why does 1st ionisation energy and nuclear attraction decreases down a group?

  - Atomic radius increases.

  - More inner shells = more shielding.

  - nuclear attraction to outer electrons decreases

  - This outweighs increase in atomic charge.
• Why does 1st ionisation energy and nuclear attraction increase across a period?

  - Nuclear charge increases

  - Same shell = similar shielding.

  - Decrease in atomic radius

  - Nuclear attraction increases
• Why is there a drop in first ionisation energy between Be and B?


  1st electron to be removed in boron is in a p subshell which is at a higher energy level than s subshell in beryllium. it is easier to remove one of the P electrons that one of the S electrons.

  
• Why is there a drop in ionisation energy between N and O?

  In nitrogen and oxygen the highest energy electron is in a P sub shell

  
  

  In O, 1st electron removed is from a paired orbital in which the electrons repel each other which makes it easier to remove an electron

  
• What is metalling bonding?


  A strong electrostatic attraction between cations and delocalised electrons

  
• How are electrons delocalised in metallic bonding?

  Each atom has donated its outer electrons to a sea of shared electrons. Which is spread across the whole structure.
• What maintains the structure of a metal?

  The cations which are fixed in position
• Why can metals conduct electricity?

  The delocalised electrons can flow and carry charge to form a current.

  
  

  This happens in all states
• Why are pure metals more malleable than alloys?

  
• Why do metals typically have high boiling points?

  Need a lot of energy to overcome the strong electrostatic attraction between cations and delocalised electrons.
• What does the melting point of a metal depend on?

  Strength of metallic bonding
• Why don't metals dissolve?

  Any interaction between lattice and solution would lead to a reaction instead of dissolving.
• What is a giant covalent lattice?

  A 3D structure of billions of atoms all held together by strong covalent bonds.

  
  

  (diff to simple molecular which only have a few atoms covalently bonded to form molecules which are then held together by weak intermolecular forces)
• Why do giant covalent lattices have such high melting points?

  Covalent bonds are strong and they need high temperatures to provide energy to overcome them
• Why are giant covalent lattices insoluble?

  Covalent bonds are too strong.
• What are the 3 lattice structures of carbon?

  Diamond, graphite and graphene

  
• What is the bond angle in diamond?

  109.5

  
  

  tetrahedral shape
• What are the two giant covalent lattices which can conduct electricity?

  
  

  why?

  Graphite and graphene

  
  

  C only uses 3 outer electrons to bond, 4th one is delocalised
• Why don't diamond or silicon conducts electricity?

  All 4 outer electrons are used in covalent bonding
• What is the structure graphite?

  A hexagonal carbon structure where 1 electron is donated to pool of delocalised electrons.

  
  

  the layers are weakly bonded by London forces.

  
• What is graphene?


  A single layer of graphite
• Why is there such variation in melting points in periods 2 and 3?

  Li, Be, Na, Mg, Al all form metallic structures.

  B, C, Si all form giant covalent structures.

  The rest form simple molecular.

  
• What are the strengths of of the bonding in periods 2 and 3?

  Covalent is strongest.

  Then metallic.

  Then simple molecular.

  
• why is there a gradual increase in boiling point from Na → Mg → Al?

  - ionic radius decreases and the charge density on the nucleus increases

  
  

  - more energy needed to overcome the electrostatic forces as the electrons are more attracted due to the higher charge density

  
• why are carbons and silicons melting point relatively high compared to other elements in their periods?

  - they both have a giant covalent structure in which each atom to covalently bonded
• why are N, O, F, NE (period 2)

  and P, S, Cl, Ar (period 3) melting points relatively so low?

  intermolecular forces (LDF) between molecules are very weak and not a lot of energy is needed to overcome them

  
  

  no covalent bonds broken

  
  

  decreases from group 5 to 8 the smaller the molecule, the less LDF between them