Periodic table period 3 Flashcards
Why can P and S form PCl3 + PCl5, P4O6 + P4O10, SO2 + SO3
P and S can exhibit variable oxidation states. They can expand their octet structure due to the presence of empty and low-lying 3d orbitals.
Ionic oxides and structure
Na2O, MgO, Al2O3
Giant ionic lattice structure
- large amount of energy required to overcome strong attractive forces between oppositely charged ions –> high melting point
Covalent oxides and structures
SiO2, P4O10, SO3
SiO2 : Giant molecular structure: Large amount of energy required to overcome strong covalent bonds between atoms –> high melting point
P4O10 and SO3: Simple molecular structure: small amount of energy required to overcome weak instantaneous dipole-induced dipole forces of attraction between molecules –> low melting point
melting point trend for the oxides
https://docs.google.com/document/d/1Vikugw-VWMbwAsNE_mTpPyQaz6xlhnUnUZfFY_7UIBg/edit?usp=sharing
Electrical conductivity for the oxides
basically, in solid state, all of the oxides are non-conducting. But in molten state, ionic oxides are good conductors as mobile ions act as mobile charge carriers
Basic oxides/hydroxides
NaOH. Na2O and Mg(OH)2, MgO
Na2O(s) + 2HCl(aq) -> 2NaCl(aq) + H2O(l)
NaOH(aq) + HCl(aq) -> NaCl(aq) + H2O(l)
MgO(s) + 2HCl(aq) -> MgCl2(aq) + H2O(l)
Mg(OH)2(aq) + 2HCl(aq) -> MgCl2(aq) + 2H2O(l)
Amphoteric oxides/hydroxides
Al(OH)3, Al2O3
Al2O3(s) + 6HCl(aq) -> 2AlCl3(aq) + 3H2O(l)
Al(OH)3(s) + 3HCl(aq) -> AlCl3 (aq) + 3H2O(l)
Al2O3(s) + 2NaOH(aq) + 3H2O(l) -> 2Na[Al(OH)4] (aq)
Al(OH)3(s) + NaOH(aq) -> NaAl(OH)4
acidic oxides
SiO2, P4O10, SO3
SiO2(s) + 2NaOH(conc) Na2SiO3(aq) + H2O(l)
P4O10(s) + 12NaOH(aq) -> 4Na3PO4(aq) + 6H2O(l)
SO3(l) + 2NaOH(aq) -> Na2SO4(aq) + H2O(l)
oxides Reaction with water and their pH
Na2O Reacts vigorously and exothermically to form
a colourless, strongly alkaline solution.
Na2O(s) + H2O(l) -> 2NaOH(aq) pH 13
MgO Very slow reaction. MgO is only slightly
soluble in water. (Appears to be insoluble.)
Reacts less readily (due to its more
exothermic lattice energy) to form a weakly
alkaline solution.
MgO(s) + H2O(l) ⇌ Mg(OH)2(aq)
pH 9
Al2O3 Does not react with water due to its highly
exothermic lattice energy.
Al2O3 is insoluble in water.
pH 7
SiO2 Does not react with water due to the strong
covalent bonds between Si and O atoms in
the giant covalent lattice structure.
SiO2 is insoluble in water.
pH7
P4O10 Reacts vigorously to form a colourless, acidic
solution.
P4O10(s) + 6H2O(l) -> 4H3PO4(aq)
pH 2
SO3 Reacts vigorously and exothermically to form
a colourless, strongly acidic solution.
SO3(l) + H2O(l) -> H2SO4(aq)
pH 2
ionic chlorides and structure
NaCl, MgCl2
Giant ionic lattice structure
Large amount of energy required to overcome strong attractive forces between oppositely charged ions –> high melting point
covalent chlorides and structure
AlCl3, SiCl4, PCl5
Simple molecular structure
small amount of energy required to overcome weak id-id forces of attraction between molecules –> low melting point
Electrical conductivity of chlorides
in solid state: non-conducting
For ionic chlorides - good conductor in molten state due to presence of mobile ions acting as mobile charge carriers
MgCl2 reaction with water
Hydration:
MgCl2(s) + 6H2O(l) -> [Mg(H2O)6] 2+(aq) + 2Cl–(aq)
Partial hydrolysis:
[Mg(H2O)6]
2+(aq) + H2O(l) ⇌
[Mg(OH)(H2O)5]
+(aq) + H3O+
(aq)
pH 6.5
slightly acidic
- undergoes partial hydrolysis
AlCl3 in water
Large amount of water
Hydration:
AlCl3(s) + 6H2O(l) -> [Al(H2O)6] 3+(aq) + 3Cl (aq)
Partial hydrolysis:
[Al(H2O)6]
3+(aq) + H2O(l) ⇌ [Al(OH)(H2O)5]
2+(aq) + H3O+(aq)
pH 3
Acidic solution
larger extent of hydrolysis but still partial
Limited amount of water:
AlCl3(s) + 3H2O(l) -> Al(OH)3(s) + 3HCl(g)
SiCl4 in water
SiCl4(l) + 2H2O(l) -> SiO2(s) + 4HCl(aq)
pH 2
Complete hydrolysis
strong acid
PCl5 reaction with water
PCl5(s) + 4H2O(l) -> H3PO4(aq) + 5HCl(aq)
pH 2
complete hydrolysis
strongly acidic
Limited amount of water or cold water:
PCl5(s) + H2O(l) ->POCl3(l) + 2HCl(g)
When more water is added:
POCl3(l) + 3H2O(l) -> H3PO4(aq) + 3HCl(aq)
Reactivity of Group 17 elements
down the group,
- more difficult for group 17 elements to gain electrons
- number of electronic shells incerases
- distance between incoming outershell electron and nucleus increases
- valence electrons experience more shielding which outweighs the increase in nuclear charge down the group
- more difficult for X2 to be reduced
- oxidising power of X2 decreases
- reactivity of X2 decreases
-
Period 3’s 3d orbitals
___is in Period 3 and has empty and low-lying 3d orbitals which can be attacked by the lone pair of electrons on the oxygen atom of water molecules.
How acidic solutions are formed from Period 3’s reactions with water
Cations with high charge density and hence high polarising power are able to distort
the electron cloud of the water molecules, weakening the O–H bond. As a result, the
O–H bond undergoes heterolytic fission readily to release H+
Electronegativity differences and the nature of bonding
- large difference in electronegativity —> ionic bond
- small or no difference in electronegativity —> covalent bond
Across period, both oxides and chlorides formed by Period 3 elements become less ionic and more covalent due to the decreasing difference in electronegativity values between oxygen/chlorine and the element
