C12 Periodic Table I Flashcards
State and explain the variation in the highest oxidation number of the elements in oxides (for Na2O, MgO, Al2O3, SiO2, P4O10, SO3) for elements in Period 3 and chlorides (NaCl, MgCl2, AlCl3, SiCl4, PCl5) of elements in Period 3.
The oxidation numbers exhibited by the elements correspond to the number of valence electrons gained or lost by an atom of the element to achieve an electronic configuration of ns2np6 in the formation of ionic compounds OR the number of electrons shared by an atom of the element in the formation of covalent bonds.
Oxidation numbers are positive integers as across the third period because both oxidation and chloride are more electronegative than each of them.
The maximum attainable oxidation of each element corresponds to the number of valence electrons in each of the elements as all the valence electrons can be used for bonding. This is because Period 3 elements have unoccupied, low-lying 3d orbitals which can be used for bonding through the expansion of the octet structure.
Pg 7 of notes
State and explain the variation in bonding in oxides (for Na2O, MgO, Al2O3, SiO2, P4O10, SO3) and chlorides (NaCl, MgCl2, AlCl3, SiCl4, PCl5) in terms of electronegativity (with the exception of AlCl3). Explain why AlCl3 has a simple molecular structure.
Across the period, both the oxides and chlorides formed by the Period 3 elements become less ionic and more covalent due to the decreasing difference in electronegativity values between oxygen/chlorine and the element. In general, a large difference in electronegativity gives rise to an ionic bond. Small or no difference in electronegativity gives rise to a covalent bond.
AlCl3 has a simple molecular structure. AlCl3 is a covalent compound because cation (Al3+) is highly charged and able to polarize the electron cloud of the anion (Cl–) easily to the large extent, this phenomenon has drawn the electrons into the space between the positive and negative ion and here, these electrons get shared. In the solid-state, AlCl3 forms a crystalline lattice of Al2Cl6 units. In the liquid and gas phases, it exists as dimers of Al2Cl6. Only at high temperatures, the dimers dissociate to give monomers of AlCl3.
Explain how ionic oxides are able to react with water.
Explain how covalent oxides are able to react with water. Describe the reaction of the oxides with water (for Na2O, MgO, Al2O3, SiO2, P4O10, SO3)
For ionic oxides to react with water, they must first be able to dissolve. The oxide ion released is a very strong Bronsted-Lowry base and will accept a H+ from the water: O2- + H2O -> 2OH-.
For covalent oxides to react with water, the high electronegative O atom of the water molecule causes the other element in the covalent oxide to be highly electron-deficient. It is hence susceptible to a nucleophilic attack by the water molecules, resulting in the formation of oxo-acids.
1) Na2O dissolves in water. It reacts vigorously with water to form a strongly alkaline solution with pH of 13-14: Na2O (s) + H2O (l) -> 2NaOH (aq).
2) MgO is slightly soluble in water as its lattice energy is highly negative and it is difficult to break the ionic lattice. Hence, it does not react so readily with water. The number of hydroxide ions formed is lower, leading to the formation of a solution with a lower pH (9-10) compared to Na2O: MgO (s) + H2O (l) ->< Mg(OH)2 (s) ->< Mg(OH)2 (aq)
3) Al2O3 is insoluble in water since the energy needed to break down the ionic lattice is more than the energy released in hydrating the ions. The resultant pH of the water remains at 7.
4) SiO2 is insoluble in water as a large amount of energy is needed to break the strong Si-O covalent bonds in the giant molecular structure. The resultant pH of the water remains at 7.
5) P4O10 reacts readily with water to form an acidic solution with a pH of 2 with the oxidation number of P remaining at +5: P4O10 (s) + H2O (l) -> 4H3PO4 (aq) (phosphoric (V) acid)
6) SO3 reacts readily with water in a highly exothermic reaction, to form an acidic solution with a pH of 2 and acidic mist, with the oxidation number of S remaining at +6: SO3 (g) + H2O (l) -> H2SO4 (aq)
Pg 9 of notes
Describe and explain the acid/base behaviour of oxides (for Na2O, MgO, Al2O3, SiO2, P4O10, SO3) and hydroxides (NaOH, Mg(OH)2, Al(OH)3), including where relevant, amphoteric behaviour in reaction with sodium hydroxide (only) and acids.
Refer to Pg 10 of notes:
Na2O (NaOH) and MgO (Mg(OH)2) are basic. They react with acids to form salt and water only.
Al2O3 (Al(OH)3) is amphoteric. It reacts with an acid to form salt and water only. It reacts with NaOH to form an aluminate complexion:
Al2O3 (s) + 2NaOH (aq) + 3H2O (l) -> 2Na+[Al(OH)4]- (aq)
Al(OH)3 (s) + NaOH (aq) -> Na+[Al(OH)4]- (aq)
SiO2, P4O10 and SO3 are acidic an rect with NaOH to form a salt (silicate, phosphate and sulfate) and water only.
SiO2 (s) + 2NaOH -> Na2SiO3 (aq) + H2O (l)
P4O10 (s) + 12NaOH (aq) -> 4Na3PO4 (aq) + 6H2O (l)
SO3 (g) + 2NaOH (aq) -> Na2SO4 (aq) + H2O (l)
Describe and explain the reactions of the chlorides with water (for NaCl, MgCl2, AlCl3, SiCl4, PCl5).
Refer to Pg 12 of notes:
Note that [Mg(H2O)6]2+ refers to an aqua complex
1) NaCl dissolves readily in water to form a resultant mixture with a pH of 7 - Hydration: NaCl (s) + aq -> Na+ (aq) + Cl- (aq)
2) MgCl2 is dissolves readily with slight hydrolysis to form a resultant mixture with a pH of 6.5
- Hydration: MgCl2 (s) + 6H2O (l) -> [Mg(H2O)6]2+ (aq) + 2Cl- (aq)
- Hydrolysis: [Mg(H2O)6]2+ (aq) ->< [Mg(H2O)5(OH)]+ (aq) + H+ (aq)
3) AlCl3 dissolves readily with appreciable hydrolysis to form a resultant mixture with a pH of 3.
- Hydration: AlCl3 (s) + 6H2O (l) -> [Al(H2O)6]3+ (aq) + 3Cl- (aq)
- Hydrolysis: [Al(H2O)6]3+ (aq) ->< [Al(H2O)5(OH)]2+ (aq) + H+ (aq)
4) SiCl4 hydrolyses to produce white fumes of HCl (g) which dissolves in excess water to give HCl (aq) to form a resultant mixture with a pH of 2.
- With limited amount of water: SiCl4 (l) + 2H2O (l) -> SiO2 (s) + 4HCl (g)
- When excess water is added: SiCl4 (l) + 2H2O (l) -> SiO2 (s) + 4HCl (aq)
5) PCl5 hydrolyses to produce white fumes of HCl (g) which dissolves in excess water to give HCl (aq) to form a resultant mixture with a pH of 2.
- When excess water is added and water is not cold:
PCl5 (l) + 4H2O (l) -> H3PO4 (aq) + 5HCl (aq)
- With limited amount of water or when the water is cold: PCl5 (l) + H2O (l) -> POCl3 (aq) + 2HCl (g)
- When more water is added: POCl3 (aq) + 3H2O (l) -> H3PO4 (aq) + 3HCl (aq)
Suggest the types of structure and bonding present in the oxides from observations of their chemical and physical properties.
By observing the physical properties of oxides, firstly, all three compounds are solids at r.t.p., indicating their high melting points. Their meting point increases from Na2O to MgO but falls when it reaches Al2O3. This could be due to the difference in their packing in the solid-state and the slight covalent character of Al2O3 (due to the high charge density and hence the high polarising power of Al3+).
This is because they are IONIC compounds where there are strong IONIC BONDS/electrostatic forces of attraction between the oppositely charged ions, giving rise to very high melting points.
Due to the presence of free-moving ions which can act as mobile charge carriers in the molten state, the three compounds are good conductors of electricity.
SiO2 is a solid at r.t.p. It has a high but relatively lower melting point than the three ionic compounds. SiO2 cannot conduct electricity due to the absence of mobile charge carriers in the molten state. Thus, this means that it has a giant molecular structure where there are strong covalent bonds between the Si and O atoms.
P4O10, SO3 have a much lower melting point compared to the previous oxides and are unable to conduct electricity. Thus, this means that mobile charge carriers are absent. Since P4O10 has a high melting point than S03 and is in the solid-state at r.t.p. while SO3 is in the liquid state at r.t.p., the two compounds have simple molecular structures where there are relatively weaker instantaneous dipole-induced dipole interactions between the molecules and the strength of the id-id interactions depends on the number of electrons/size of the electron cloud.
Suggest the types of structure and bonding present in the chlorides from observations of their chemical and physical properties.
4) NaCl, MgCl2:
Due to the large differences in electronegativity, the two compounds are ionic compounds where there are strong ionic bonds/electrostatic forces of attraction between the oppositely charged ions. Due to the presence of free-moving ions which can act as mobile charge carriers in the molten state, the three compounds are good conductors of electricity. They are solids at r.t.p. However, unlike the oxides, the melting point decreases from NaCl to MgCl2. This is because due to the high charge density and hence the high polarising power of Mg2+, MgCl2 has a slight covalent character.
5) AlCl3, SiCl4, PCl5:
These compounds have a simple molecular structure where there are relatively weaker instantaneous dipole-induced dipole interactions between the molecules and the strength of the id-id interactions depends on the number of electrons/size of the electron cloud. Since in the solid-state, AlCl3 forms a crystalline lattice of Al2Cl6 units and exists as dimers of Al2Cl6 in the liquid and gas phases (Only at high temperatures, the dimers dissociate to give monomers of AlCl3), electron cloud size decrease from Al2Cl6 to SiCl4 before increasing in PCl5 and the boiling points have the same trend.
Explain what happens when a salt is dissolved in water and undergoes hydrolysis. Describe how the process of hydrolysis and the pH of the resultant mixture will be affected by the polarising power of the cation.
When a salt is dissolved in water, the cations are naturally hydrated by the water molecules due to the formation of ion-dipole interactions. There is no hydrolysis of the ions (i.e. no reaction with water). As a result, the water molecules coordinate to the cation via the lone pair of electrons on the oxygen. During hydrolysis, 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 to readily release H+ ion. The stronger the polarising power of the cation, the greater the extent of hydrolysis
Explain what happens when a salt is dissolved in water and undergoes hydrolysis. Describe how the process of hydrolysis and the pH of the resultant mixture will be affected by the polarising power of the cation.
When a salt is dissolved in water, the cations are naturally hydrated by the water molecules due to the formation of ion-dipole interactions. There is no hydrolysis of the ions (i.e. no reaction with water). As a result, the water molecules coordinate to the cation via the lone pair of electrons on the oxygen. During hydrolysis, 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 to readily release H+ ion. The stronger the polarising power of the cation, the greater the extent of hydrolysis. Hence, more H+ ions will be produced and the pH of the solution will be lower.
Explain why SiCl4 can undergo hydrolysis when added to water but CCl4 cannot.
For covalent chlorides, during hydrolysis, the central atom needs to accept the lone pair of electrons from water. Since the silicon atom is larger in size than the carbon atom, there is less steric hindrance around the Si atom for the water molecule (nucleophile) to attack. Si also has an empty low-lying 3d orbitals to accept the lone pair of electrons from water.
On the other hand, CCl4 is unable to undergo hydrolysis when added to water as the electron-deficient carbon atom is attached to four large chlorine atoms
