Questions on periodicity: trends across Period 3 and the blocks of the table

Question 1

Describe and explain the trend in the melting points of the elements across Period 3 from sodium to argon.

Accepted Answer

Indicative content

the melting point rises from sodium to aluminium, rises steeply to a maximum at silicon, falls sharply at phosphorus, rises at sulfur and falls to a minimum at argon
sodium, magnesium and aluminium have metallic bonding, and the ion charge and the number of delocalised electrons per atom increase while the ions get smaller
so the metallic bonding gets stronger from sodium to aluminium and more energy is needed to overcome it
silicon is giant covalent, and melting breaks many strong covalent bonds, which needs the most energy of any element in the period
phosphorus, sulfur, chlorine and argon are simple molecular, P₄, S₈, Cl₂ and Ar, and melting overcomes only weak van der Waals forces between molecules
the van der Waals forces are stronger for larger molecules with more electrons, giving the order S₈ > P₄ > Cl₂ > Ar

Levels

Level 3 (5–6 marks) the whole trend is described and every part of it is explained from structure and bonding: the strengthening metallic bonding from sodium to aluminium, the giant covalent silicon at the maximum, and the molecular elements ordered by the size of their molecules; covalent bonds are broken on melting only in silicon
Level 2 (3–4 marks) the metallic, giant covalent and molecular regions are identified and two are explained; the order among the molecular elements, or why metallic bonding strengthens, is not covered
Level 1 (1–2 marks) silicon is said to have the highest melting point, or the molecular elements low ones; phosphorus and sulfur are commonly said to melt by breaking covalent bonds

Question 2

Identify the block of the periodic table to which calcium, [Ar]4s², and bromine, [Ar]3d¹⁰4s²4p⁵, belong, giving a reason for each. Then give the full electron configuration of silicon and identify its block.

Accepted Answer

B1 calcium is in the s block, because its last electron, in filling order, enters the 4s subshell B1 bromine is in the p block, because its last electrons enter the 4p subshell; reject d block because 3d is full B1 silicon is 1s² 2s² 2p⁶ 3s² 3p²; reject [Ne] shorthand where the full configuration is asked for B1 silicon is a p-block element, because its last two electrons enter the 3p subshell

Question 3

Explain why, although the first ionisation energy rises across Period 3, it falls from magnesium to aluminium and from phosphorus to sulfur.

Accepted Answer

B1 aluminium's outer electron is removed from a 3p orbital, whereas magnesium's comes from 3s B1 the 3p electron is higher in energy and is partly shielded by the 3s electrons, so less energy is needed to remove it B1 in sulfur the electron removed is from a 3p orbital that holds two electrons, whereas each 3p electron in phosphorus is unpaired B1 repulsion between the paired electrons means less energy is needed to remove one of them; ignore any reference to a half-filled subshell being stable

Question 4

Explain, using the melting points in the table, why phosphorus, chlorine and argon all melt below 400 K, and why they melt in the order phosphorus, then chlorine, then argon.

Accepted Answer

B1 phosphorus and chlorine are simple molecular and argon exists as single atoms B1 melting overcomes only weak van der Waals forces between the particles; no covalent bonds are broken, so little energy is needed B1 P₄ has 60 electrons, Cl₂ 34 and Ar 18, so the van der Waals forces are strongest between P₄ molecules and weakest between argon atoms B1 stronger forces need more energy to overcome, giving the order P₄ > Cl₂ > Ar; reject any answer based on the strength of the P–P or Cl–Cl bond

Question 5

Calculate, from the atomic radii in the table, the percentage decrease in atomic radius from sodium to silicon, and from silicon to chlorine.

Accepted Answer

M1 sodium to silicon: (0.186 − 0.117) ÷ 0.186 × 100 A1 decrease from sodium to silicon = 37.1%, to three significant figures M1 silicon to chlorine: (0.117 − 0.099) ÷ 0.117 × 100 A1 decrease from silicon to chlorine = 15.4%, to three significant figures; allow 15%

Question 6

Compare the trend in melting point across Period 2 with the trend across Period 3.

Accepted Answer

B1 in both periods the melting point rises to a maximum at group 4, carbon in Period 2 and silicon in Period 3 B1 both maxima are giant covalent structures, in which melting breaks many strong covalent bonds B1 in both periods there is a sharp fall at group 5, nitrogen and phosphorus, which are simple molecular B1 in Period 2 the group 3 element, boron, is giant covalent, whereas in Period 3 aluminium is metallic

Question 7

Explain why the atomic radius of the elements decreases across Period 3 from sodium to chlorine.

Accepted Answer

B1 the nuclear charge increases across the period, from 11 protons in sodium to 17 in chlorine B1 the outer electrons are all in the same shell, the third, and the shielding by the inner electrons is similar; reject 'shielding increases' B1 so the outer electrons are attracted more strongly to the nucleus and are held closer to it

Question 8

Explain the general increase in first ionisation energy across Period 3.

Accepted Answer

B1 the nuclear charge increases across the period, one proton at each step B1 the electron removed is in the same shell each time, with similar shielding, and the atomic radius is smaller B1 so the outer electron is more strongly attracted to the nucleus and more energy is needed to remove it

Question 9

Explain why magnesium has a higher melting point than sodium.

Accepted Answer

B1 both metals have metallic bonding: a lattice of positive ions attracted to a sea of delocalised electrons B1 magnesium ions are Mg²⁺ and smaller than Na⁺, and each magnesium atom gives two delocalised electrons rather than one B1 so the attraction between the ions and the delocalised electrons is stronger in magnesium and more energy is needed to overcome it

Question 10

Explain why silicon has the highest melting point of the Period 3 elements.

Accepted Answer

B1 silicon has a giant covalent structure; accept macromolecular B1 each atom is covalently bonded to four others throughout the structure, so melting breaks many strong covalent bonds B1 this needs much more energy than overcoming the metallic bonding of Na, Mg and Al or the van der Waals forces between the molecules of P₄, S₈ and Cl₂

Question 11

Suggest the structure of germanium, which is directly below silicon in group 4, and suggest why germanium melts at 1211 K, lower than silicon's 1687 K.

Accepted Answer

B1 germanium has a giant covalent structure like silicon's, each atom covalently bonded to four others B1 germanium atoms are larger, so the Ge–Ge bonds are longer than Si–Si bonds B1 longer bonds are weaker, so less energy is needed to break them and the melting point is lower

Question 12

Evaluate, using the data in the table, a student's statement that aluminium should melt at a much higher temperature than magnesium because its metallic bonding is stronger.

Accepted Answer

B1 the metallic bonding is stronger in aluminium, because Al³⁺ is smaller and more highly charged than Mg²⁺ and each atom gives three delocalised electrons B1 the melting points do not support 'much higher': they differ by only 10 K, because melting needs the bonding loosened only enough for the ions to slide past one another B1 the boiling points, where the ions separate completely, support the stronger bonding, so the reasoning is sound but the prediction about melting point is not

Question 13

Outline how the electrical conductivity of the elements changes across Period 3, with a reason for each part of the pattern.

Accepted Answer

B1 sodium, magnesium and aluminium conduct, increasingly well, because the number of delocalised electrons per atom increases B1 silicon is a semiconductor, because its electrons are held in covalent bonds but a few are freed as the temperature rises B1 phosphorus, sulfur, chlorine and argon do not conduct, because they have no delocalised electrons or ions

Question 14

Explain why the ionic radius falls from Na⁺, 0.102 nm, to Mg²⁺, 0.072 nm, to Al³⁺, 0.054 nm.

Accepted Answer

B1 the three ions are isoelectronic, each with ten electrons in the configuration 1s²2s²2p⁶ B1 the nuclear charge increases from 11 to 12 to 13 protons while the number of electrons stays the same B1 so the electrons are attracted more strongly to the nucleus and held closer, and the ionic radius decreases

Question 15

Explain why the sulfide ion, S²⁻, radius 0.184 nm, is larger than the sulfur atom, radius 0.104 nm.

Accepted Answer

B1 the sulfide ion has two more electrons than the atom, 18 against 16, added to the third shell, so the number of shells is unchanged B1 the nuclear charge is the same, 16, so it attracts more electrons and each is held less strongly B1 the extra electrons increase the repulsion between the outer electrons, so they move further from the nucleus and the ion is larger

Question 16

Justify the statement that argon has the lowest melting point of the Period 3 elements, although solid chlorine and solid argon are held together by the same kind of force.

Accepted Answer

B1 both solids are held together only by induced dipole–dipole forces between their particles, single atoms in argon and Cl₂ molecules in chlorine, with no covalent bonds between the particles to break B1 each argon atom has 18 electrons, whereas each chlorine molecule has 34, so the temporary dipoles are smaller and the forces between argon atoms are weaker B1 less energy is needed to separate argon atoms, so argon melts at 84 K and chlorine at 172 K; reject any mention of breaking Cl–Cl bonds

Question 17

Calculate the energy, in kJ, needed to turn 0.540 g of gaseous aluminium atoms into Al⁺(g) ions. The first ionisation energy of aluminium is 578 kJ mol⁻¹ and its relative atomic mass is 27.0.

Accepted Answer

M1 amount of aluminium = 0.540 ÷ 27.0 = 0.0200 mol A1 energy = 0.0200 × 578 = 11.6 kJ, to three significant figures; error carried forward from the amount

Question 18

Suggest why tables of atomic radius for Period 3 often give no value for argon that can be compared with the others.

Accepted Answer

B1 atomic radius is measured as half the distance between the nuclei of two bonded atoms of the element B1 argon atoms do not form bonds, so any value is a van der Waals radius between non-bonded atoms, which is larger and not comparable

Question 19

Suggest which of K⁺ and Cl⁻ is the smaller ion, and why. Both have the electron configuration 1s²2s²2p⁶3s²3p⁶.

Accepted Answer

B1 K⁺ is smaller than Cl⁻ B1 the two ions are isoelectronic, but potassium has 19 protons against chlorine's 17, so the same 18 electrons are pulled closer to the nucleus in K⁺

Question 20

State what is meant by the term periodicity.

Accepted Answer

B1 a repeating pattern of physical or chemical properties across the periods of the periodic table; allow 'a trend that repeats in each period'

Element	Formula of particle	Melting point / K
Phosphorus	P₄	317
Chlorine	Cl₂	172
Argon	Ar	84

Element	Melting point / K	Boiling point / K
Magnesium	923	1363
Aluminium	933	2792