Periodic Table

• Atomic radius is half the distance between the centres of 2 adjacent atoms in the structure of the element
• Metallic radius is half the distance between the centres of 2 adjacent atoms in a metal
• covalent radius is half the distance between the centres of 2 adjacent atoms that are covalently bonded
• van der Waal’s radius is half the distance between the centres of 2 adjacent atoms which are not chemically bonded
• Across the period, atomic radii decreases

  • number of protons increases, nuclear charge increases
  • same number of inner shell electrons, shielding effect approximately constant
  • increase in effective nuclear charge
  • increase in electrostatic forces of attraction between nucleus and valence electrons
• ionic radii decreases

  • nuclear charge increases
  • same inter-electronic repulsion (same number of electrons)
  • stronger electrostatic forces of attraction
• Anions bigger than cations

  • anions have one more filled principal quantum shell
  • valence electrons further from nucleus
  • weaker electrostatic forces of attraction
• 1st IE increases across the period:

  • nuclear charge increases
  • shielding effect approximately constant
  • increase in effective nuclear charge
  • increase in electrostatic forces of attraction
  • more energy needed to remove valence electron
• 1st IE of Al (grp 13) lower than Mg (grp 2)

  • 3p electron is at higher energy level
  • less energy needed to remove
• 1st IE of S (grp 16) lower than P (grp 15)

  • inter-electronic repulsion between paired electron of the same orbital in S
  • less energy needed to remove
• Melting point of giant metallic lattice structure

  • high due to large amount of energy needed to overcome strong metallic bonds
  • as number of delocalised electrons increase, smaller cationic radius --> higher charge density, more energy needed to overcome strong electrostatic forces of attraction between metal cations and sea of delocalised electrons
• Electrical conductivity of giant metallic lattice structure

  • sea of delocalised electrons as mobile charge carriers
• Melting point of giant molecular structure

  • large amount of energy needed to overcome strong and extensive covalent bonds between atoms
• Melting point of simple molecular structures

  • low, small amount of energy needed to overcome weak id-id interactions
  • melting point increases with increase in number of electrons, increase in polarisability of electron cloud and thus more energy needed to overcome stronger id-id
• across the period, oxidation numbers increases

  • number of valence electrons available increases
• P and S can exhibit multiple oxidation numbers

  • presence of vacant and energetically accessible d orbitals, can be used for bonding through expansion of octet structure
• sodium oxide dissolves completely in water to form sodium hydroxide (pH14 solution)

  magnesium oxide dissolves partially in water due to high lattice energy (pH9 solution)

  • basic O2- ion reacts with H2O by accepting H+ to form hydroxide ion (forms 2 hydroxide ions)
• aluminium oxide does not dissolve in water due to extremely high lattice energy, large amount of energy needed to break strong ionic bonds (pH7 solution)
• Silicon dioxide is a giant molecular structure, does not dissolve in water due to strong and extensive covalent bonds between Si and O atoms
• P4O10 and SO3 dissolves in water to give phosphoric acid and sulfuric acid

  • S=O and P=O bond reacts with H2O to give P/S-O-H bonds, then dissociate to give H+
• Group 2 elements atomic and ionic radius increases down the group

  • both nuclear charge and shielding effect increase, valence electrons located further away from the nucleus
  • weaker electrostatic forces of attraction between nucleus and valence electrons
• group 2 elements first IE decreases down the group

  • nuclear charge and shielding effect increase, valence electrons located further away from the nucleus
  • weaker electrostatic forces of attraction betw nucleus and valence electrons, less energy needed to remove
• group 2 elements electronegativity decreases down the group

  • weaker electrostatic forces of attraction between nucleus and electron pair in covalent bond
• group 2 elements melting point decreases down the group

  • size of cations increases, charge density decreases (charge remains the same)
  • weaker electrostatic forces of attraction between cations and sea of delocalised electrons
  • less energy needed to overcome
• Group 17 elements (halogens) structure and bonding

  • simple molecular structure
  • weak Id-id between non-polar molecules
• group 17 elements boiling point increases down the group

  • number of electrons increase
  • polarisability of electron cloud increase
  • more energy needed to overcome stronger id-id
• structure and bonding of hydrogen halides (HX)

  • simple molecular structure
  • weak id-id and pd-pd between polar molecules
• Boiling point of hydrogen halides increases down the group

  • number of electrons increases
  • polarisability of electron cloud increases
  • more energy needed to overcome stronger id-id
• thermal stability of hydrogen halides decrease down the group

  • size of halogen atom increases
  • effectiveness of orbital overlap decreases
  • bond strength decreases
  • less energy required to break bond
• Basic oxides: sodium oxide, magnesium oxide
• Amphoteric oxide: aluminium oxide
• acidic oxides: phosphorus pentoxide, sulfur trioxide
• sodium chloride undergoes complete hydration, but does not undergo hydrolysis
• magnesium chloride undergoes:

  • hydration, forms hydrated metal ions (complex ion) and chloride ions
  • slight hydrolysis, metal cation has high charge density, polarises electron cloud of surrounding water molecules, break O-H bond to release H+ (results in acidic solution, pH 6.5)
  • dissolves completely in water

  
• aluminium chloride:

  • undergoes hydration, forms hydrated metal ions (complex ion) and chloride ions
  • slight hydrolysis, metal cation has high charge density (aluminium higher than magnesium), polarises electron cloud of surrounding water molecules, can break O-H bonds, release H+ (results in acidic solution, pH 3)
  • dissolves completely in excess water
• Metal ions that form complexes: magnesium, aluminium

  • water molecules form dative bonds with cation
  • metal ion uses 1 empty 3s, 3 3p, 2 3d orbitals to form 6 hybrid orbitals, accepts 6 lone pairs from 6 water molecules
• Aluminium chloride in limited water:

  • complete hydrolysis, has vacant and energetically accessible 3p orbital, can accept lone pair from water molecules to form dative bonds
  • reacts vigorously, white solid of aluminium hydroxide/aluminium oxide forms, white fumes of HCl
  • pH 3
• silicon tetrachloride:

  • does not undergo hydration
  • complete hydrolysis, has vacant and energetically accessible 3d orbitals, can accept lone pair from water molecules, form dative bond
  • reacts vigorously with water, forms white solid (silicon dioxide), white fumes of HCl
  • pH 2
• phosphorus pentachloride:

  • does not undergo hydration
  • complete hydrolysis, P has vacant and energetically accessible 3d orbitals, can accept lone pair from water molecules, forms dative bond
  • reacts vigorously with water, white fumes of HCl
  • pH 2
• down group 17:

  • atomic radius increases, lower tendency to accept electrons
  • lower tendency for atom to be reduce to anion
  • oxidising power of halogens decreases down the group, E value becomes less positive (smaller tendency to be reduced, weaker OA)