Chapter 21 - Periodicity

Question 1

Why are oxidation numbers of elements across period 3 always positive (except fluorine)?

Answer 1

Oxygen is the most electronegative element other than fluorine.

Question 2

Why do elements across period 3 display increasing variation of possible oxidation numbers?

Answer 2

P, S and Cl exhibit a wider range of oxidation numbers because they can expand their octets (by utilising vacant d-orbitals).

Question 3

Why does the bonding of period 3 oxides change from ionic to covalent across the period?

Answer 3

The difference in electronegativity between each element and oxygen decreases across the period, thus bonding becomes increasingly covalent, since covalent bonds usually form between atoms of similar electronegativity.

Question 4

In period 3, why is the melting point of MgO higher than that of Na2O?

Answer 4

Mg2+ has a smaller radius and higher charge than Na+, thus lattice energy of MgO is much greater in magnitude than that of Na2O. More energy is needed to overcome the electrostatic forces of attraction between oppositely charged ions in MgO.

Question 5

In period 3, why is the melting point of Al2O3 lower than that of MgO?

Answer 5

Although Al2O3 technically should have larger LE than MgO, Al3+ has very high charge density which can polarise the O2- ion (even though O2- is small and normally not so easily polarised). Al2O3 becomes more covalent and causes its melting point to decrease.

Question 6

In period 3, why does SiO2 have a high melting point?

Answer 6

SiO2 has a giant molecular structure with strong covalent bonds between Si and O atoms. Each Si atom is bonded to 4 O atoms, and each O atom is bonded to 2 Si atoms. A large amount of energy is needed to BREAK the strong covalent bonds.

Question 7

In period 3, why does oxides of phosphorus and sulfur have lower melting points?

Answer 7

They have a simple covalent structure and have relatively weak intermolecular dispersion forces. Lesser energy is needed to overcome the weaker dispersion forces.

Question 8

Why are ionic (metal) oxides basic?

Answer 8

They contain O2- ions in their lattice, and the O2- ions readily hydrolyse in water to form OH- ions.

Question 9

Why are covalent (non-metal) oxides acidic?

Answer 9

They form acidic solutions while reacting with water (not all), producing H3O+ and an oxo-anion (like SO42-)

Question 10

How does Na2O react with water and acids? What is its pH in water?

Answer 10

It dissolves completely in water via a vigorous and exothermic reaction.
PH: arnd 13
It dissolves in acids via an exothermic reaction, forming a salt solution.

Question 11

How does MgO react with water and acids? What is its pH in water?

Answer 11

Water: very slow with limited solubility, so not all hydroxide ions formed are released into the solution (pH lower than Na2O, arnd 9)
Acids: dissolve to form a salt solution

Question 12

Why is Al2O3 amphoteric?

Answer 12

Al3+’s high charge density gives it great polarising power, allowing the O2- ion to be slightly polarised despite its small size. This confers some covalent character to Al2O3. This mixture of ionic and covalent character allows Al2O3 to display both basic and acidic properties, and is thus amphoteric.

Question 13

How does Al2O3 react with water, acids and strong bases? What is its pH in water?

Answer 13

Water: insoluble in water (pH 7)
Acids: dissolves to form a salt solution
Bases: dissolves in EXCESS base to form a salt solution, forming [Al(OH)4]- as the anion

Question 14

How does SiO2 react with water and strong bases? What is its pH in water?

Answer 14

Water: does not react with water (pH 7)
Strong bases: no reaction with hot aqueous bases; requires hot and concentrated strong bases. A silicate salt (SiO3 2-) is produced

Question 15

How does P4O10 react with water and strong bases? What is its pH in water?

Answer 15

Water: reacts violently in water to give an acidic solution of H3PO4; pH2
Strong bases: dissolves in excess to form a salt solution

Question 16

How does SO3 react with water and bases? What is its pH in water?

Answer 16

Water: violent and very exothermic reaction, producing very acidic mist of H2SO4 droplets (pH 1)
Bases: react directly with bases to form salt solution

Question 17

Why does AlCl3 sublime at ~180ºC?

Answer 17

In the vapour phase, AlCl3 is simple molecular in structure, forming gaseous Al2Cl6 dimers. There exists an equilibrium between Al2Cl6 dimers and AlCl3 monomers. As the temperature increases, the position of equilibrium shifts right to favour the monomers.

Question 18

How do ionic chlorides across period 3 react with water?

Answer 18

The ions get solvated and form favourable ion-dipole interactions with polar water molecules.

Question 19

Why is the Cl- of period 3 chlorides not responsible for deviations from neutral pH when dissolving in water?

Answer 19

It is a very poor conjugate base of the strong acid HCl, and cannot react with water.

Question 20

How does NaCl react with water & what is its pH? Why?

Answer 20

It dissolves in water to form a colourless solution of pH7.

Since Na+ has low charge density, it does not undergo hydrolysis. Thus, solution remains neutral.

Question 21

How does MgCl2 react with water & what is its pH? Why?

Answer 21

It dissolves in water to form a colourless solution of pH 6.5.
Due to the high charge density of Mg2+ (compared to Na+), the hydrated magnesium ion undergoes slight hydrolysis to form a very weak acidic solution.

Question 22

How does AlCl3 react with water & what is its pH? Why?

Answer 22

Limited water: steamy white fumes of HCl evolves and a white solid Al(OH)3 remains, which is insoluble in water.
Excess water: AlCl3 dissolves to form a colourless solution of pH3. Al3+ forms a complex ion with 6 water molecules, and upon further hydrolysis in water, H3O+ is formed.

Question 23

Describe the acidic behaviour of Al3+ (aq). (3)

Answer 23

1) adding a controlled amount of OH- ions to an aqueous solution of Al3+ cam further deprotonate water ligands, until a white ppt of Al(OH)3 is formed.
2) adding excess OH- ions will cause the white ppt to dissolve and form a colourless solution with aluminate ion [Al(OH)4]-
3) A solution of Al3+ is acidic enough to react with sodium carbonate to produce Al(OH)3 and CO2 gas.

Question 24

How do SiCl4 and PCl5 react with water & what are their pH? Why?

Answer 24

Violent reaction with water, producing fumes of HCl gas. They undergo complete hydrolysis and form a strong acidic solution of pH 1 (containing hydrochloric acid)

Question 25

How does reactivity of Group 2 metals differ down the group? Why?

Answer 25

Reactivity of Group 2 metals increases down the group. As atomic radii increase, the metal atoms lose their electrons more readily (1st and 2nd ionisation energies decrease) going down the group. Thus, they form cations more easily. (Thus their reducing powers increase)

Question 26

How does the thermal stability of Group 2 carbonates change down the group?

Answer 26

Group 2 carbonates become more thermally stable down the group. The ones lower in the group need to be heated more strongly before they decompose, hence they will have a higher decomposition temperature.

Question 27

What are 2 factors affecting thermal stability of Group 2 carbonates?

Answer 27

1) Polarising power of cation: cationic radius increases down the group while charge remains the same, thus charge density decreases down the group. Polarising power of cations decrease down the group and are less able to distort the electron cloud of the carbonate, weakening the C–O bonds within the carbonate anion to a smaller extent. Covalent bonds within the carbonate anion are less likely to be broken down by the group. Thus, the ease of decomposition decreases and thermal stability increases.
2) Polarisability of anions: Larger anions are more susceptible to polarisation of their electron cloud. The greater the polarisability, the lower the thermal stability. Only polyatomic anions are susceptible to decomposition as monoatomic ions like Cl- cannot be broken down further.

Question 28

Describe the thermal stability of Group 1 carbonates and explain.

Answer 28

Group 1 carbonates are resistant to decomposition as they have low charge density and are unable to sufficiently distort the carbonate’s electron cloud, except Lithium which has high enough charge density.

Question 29

How does the volatility and melting & boiling points of Group 17 elements (halogens) differ down the group? Why?

Answer 29

Down the group, the size of electron cloud increases and hence, the polarisability of the halogen molecule increases. More energy is needed to overcome the strong dispersion forces between the molecules. Hence, volatility decreases down the group while melting & boiling points increase down the group.

Question 30

How does the oxidising power of halogens change down the group and why?

Answer 30

The outermost electron shell of halogens contains 7 electrons (ns2np5). Their chemistry is dominated by a tendency to gain a completely filled valence electron shell. Hence, they tend to be reduced in a redox reaction. Since their Eº values become less positive down the group, the oxidising power of halogens decreases down the group and they become less likely to be reduced in a redox reaction.

Question 31

What are 3 kinds of redox reactions halogens can take part in?

Answer 31

1) displacement reaction: a halogen higher in the group can oxidise a halide below it; a more reactive halogen will displace a less reactive one from its compounds.
2) In the presence of a halogen oxidising agent, an aqueous solution of Fe2+ ions can be oxidised to Fe3+.
3) Chlorine and bromine can oxidise thiosulfate (S2O3 2-) to sulfate, while iodine oxidises thiosulfate to tetrathionate (S4O6 2-). Iodine is a weaker oxidising agent compared to Cl and Br, thus the oxidation number change in sulfur is smaller.

Question 32

How does the thermal stability of Group 17 hydrogen halides (hydrides) change down the group and why?

Answer 32

Down the group, as atomic radius increases from F to I, the bond length of the H–X bond increases and bond strength decreases. Hence, less energy is needed to break the H–X bond. Thus the thermal stability of hydrogen halides decreases down the group.

Question 33

What are diagonal relationships in the periodic table and why does this occur?

Answer 33

Certain pairs of diagonally adjacent elements in period 2 and 3 show similarities in their properties. As a result, the first element of some groups can be seen to have slightly different properties to the rest of the group.

This is because, as we move across the period and down the group, properties tend to have opposite trends. Thus, moving simultaneously across the period and down the group by one element results in the combined effects of the opposing trends “cancelling out”.