Atomic radius
Dedicated post explaining the trend in atomic radius is right here!
Melting points and structure
The melting point of the metal elements increases as we move from sodium to aluminium which reflects an increased strength in metallic bonding. The atomic radius decreases so the ions can pack together more closely and there is an increase in cationic charge (Na+, Mg2+, Al3+) meaning there are more delocalised valence electrons and stronger electrostatic attractions in the metallic lattice of aluminium compared with sodium.
Silicon is a semi-metal with a crystalline giant covalent structure similar to diamond. Each silicon atom forms four sp3 hybrid orbitals that overlap to form four strong covalent bonds with neighbouring atoms, hence the very high melting point.
The non-metal elements each have a simple molecular structure, hence the very low melting points. It doesn’t take much energy to break the intermolecular bonds between the molecules, and the trend seen above can be explained in terms of the size of the molecules that make up each element (bigger molecules = stronger intermolecular bonding).
Phosphorus exists as a number of allotropes (same element, different arrangement of atoms) but the two most common are white phosphorus and red phosphorus. White phosphorus is a waxy solid at room temperature formed from tetrahedral P4 molecules – the angles of 60° place each molecule under strain which explains its instability and highly reactive nature.
Over time (especially in the presence of heat and light) white phosphorus gradually changes into red phosphorus which has a polymer-like structure with molecules of P4 bonding with neighbours to form short chains. Red phosphorus is slightly more stable (it ignites at ∼ 300°C) as there is less strain within the molecules.
Sulphur also exists as a number of allotropes but the most common is rhombic sulphur which is a soft, yellow, crystalline solid at room temperature formed from cyclic octatomic S8 molecules.
Chlorine is a yellow-green diatomic gas, Cl2, at room temperature and argon is a monatomic Noble gas.
Ionisation energies
The first ionisation energy is the energy change when an electron is removed from an atom in the gas phase.
Na(g) ⇾ Na+(g) + e– 1st IE = + 496 kJ mol-1
The overall trend is that the 1st ionisation energy becomes increasingly endothermic as we move across period 3.
We are removing an electron from a progressively more positive nucleus and all the electrons experience this greater effective nuclear charge pulling them in more tightly, hence more energy is needed to remove an electron.
The blips can be explained by considering the detailed electronic configuration of each element (indeed the trend in ionisation energy is evidence for this model of atomic structure).
- in aluminium, the outermost electron being removed is in a 3p orbital which is higher in energy than a 3s orbital so it takes a little less energy to remove than might be expected.
- phosphorus has the electron configuration [Ne] 3s2 3p3 with each of the electrons in the 3p orbital occupying their own sub-orbital (px, py, pz). In sulphur, the electron being removed is spin-paired in one of the 3p sub-orbitals and the resulting electron-electron repulsions means that once again, less energy is needed to remove it than might be expected.
Reactions of the metals
- the metals of Period 3 all burn in air to form metal oxides
Sodium burns in oxygen to form sodium peroxide – the peroxide ion, O22-, is stabilised by the relatively large size of the Na+ ion.
2Na(s) + O2(g) ⇾ Na2O2(s)
At high temperatures sodium peroxide breaks down to form sodium oxide.
Na2O2(s) + O2(g) ⇾ Na2O(s)
Magnesium burns in oxygen to produce magnesium oxide, although if burning in air magnesium nitride, Mg3N2, is also formed.
2Mg(s) + O2(g) ⇾ 2MgO(s)
Although aluminium is a reactive metal the aluminium oxide coating is strongly bonded to the surface of the metal and this prevents further reactions. However, if aluminium powder is heated and lowered into a jar of oxygen it will burn to form the oxide.
4Al(s) + 3O2(g) ⇾ 2Al2O3(s)
- sodium and magnesium will react with cold water to give the metal hydroxide and hydrogen.
Sodium hydroxide is a strongly alkaline solution as the hydroxide is very soluble.
2Na(s) + 2H2O(l) ⇾ 2NaOH(aq) + H2(g)
Magnesium hydroxide is a weakly alkaline solution because the hydroxide is only sparingly soluble.
Mg(s) + 2H2O(l) ⇾ Mg(OH)2(aq) + H2(g)
- sodium and magnesium react with acids to give the metal salt and hydrogen
2Na(s) + 2H3O+(aq) ⇾ 2Na+(aq) + H2(g) + 2H2O(l)
Mg(s) + 2H3O+(aq) ⇾ Mg2+(aq) + H2(g) + 2H2O(l)
Aluminium is amphoteric, dissolving in both acids and alkalis to give a complex ion and hydrogen.
2Al(s) + 6H3O+(aq) + 6H2O(l) ⇾ 2 [Al(H2O)6]3+(aq) + 3H2(g)
2Al(s) + 2OH–(aq) + 6H2O(l) ⇾ 2 [Al(OH)4]–(aq) + 3H2(g)
Reactions of the non-metals
- reactions with oxygen
Silicon is coated in a layer of silicon oxide (like aluminium) which protects the element and renders it mostly inert but it will react with oxygen at high temperatures (>900°C) to form a number of giant covalent structures with single Si-O bonds.
Si(s) + O2(g) ⇾ SiO2(s)
Sulphur burns to give sulphur dioxide.
S8(s) + 😯2(g) ⇾ 8SO2(g)
Phosphorus burns in a limited supply of oxygen to give phosphorus(III) oxide, P4O6, and phosphorus(V) oxide, P4O10, in excess oxygen.
Chlorine doesn’t react directly with oxygen but in the presence of high energy ultraviolet light it will form ClO• radicals – you can read more about these stratospheric reactions here.
- silicon, phosphorus and sulphur do not react with water, but chlorine does forming hypochlorous acid (the reactions of chlorine are covered extensively here).
Cl2(g) + 2H2O(l) ⇌ HOCl(aq) + H3O+(aq) + Cl–(aq)
Practice questions
- (a) Explain why the melting temperature of silicon (1414°C) is so different to that of white phosphorus (44°C) despite the fact that they are both non-metal elements.
(b) Sulphur and chlorine are both elements in Period 3 with a simple molecular structure. Explain why sulphur is a soft solid at room temperature and chlorine is a gas.
- The second ionisation energies of the elements of period 3 are shown below.
(a) Define the term periodicity.
(b) Write an equation to show the 2nd ionisation energy of sodium.
(c) Predict a value for the 2nd ionisation energy of aluminium. Explain the reasoning for your prediction.
- Describe and explain the trend in the third ionisation energies for Period 3 (shown below) in terms of the detailed electronic configurations of the elements.
Answers
- (a) Silicon has a giant covalent network / macromolecular structure with strong covalent bonds between the atoms which require a lot of energy to break when silicon melts, hence the high melting temperature. White phosphorus is composed of simple molecules of P4 with London forces / instantaneous dipole – induced dipole bonds between the molecules. When white phosphorrus is melted, we only need enough energy to break these intermolecular bonds, not the the intramoleular covalent bonds, hence the much lower melting temperature.
(b) Chlorine is a diatomic molecule, Cl2 and sulphur exists as molecules of S8. Larger molecules have a greater electron density and stronger London forces / instantaneous dipole – induced dipole bonds which take more energy to break, hence the higher melting temperature for sulphur.
2. (a) Periodicity is the repeating pattern or trend in physical and chemical properties of elements across a Period in the Periodic Table.
(b) Na+(g) ⇾ Na2+(g) + e– (must include state symbols)
(c) The second ionisation energy of aluminium is approximately 1800°C, slightly higher than magnesium and silicon but following the general trend (from Mg to Ar, the second ionisation energy increases). Aluminium has the electronic configuration [Ne] 3s2 3p1. The 3p electron is the first electron to be lost, so the second ionisation energy involves the removal of a 3s electron which is in an orbital / shell closer to the nucleus with less shielding. Since this 3s electron experiences a higher effective nuclear charge it is more strongly attracted and so it takes more energy than expected to remove, hence the higher than expected value for the second ionisation energy.
3. The general trend in third ionisation energy is an increase from aluminium to argon. This is because as we move across Period 3 the electrons experience a greater effective nuclear charge so it takes more energy to remove one (there is one more proton in the nucleus of each atom but the extra electrons are all in the same shell and so are equally shielded). Sodium and magnesium have very high third ionisation energies as the electron being removed is coming from a 2p orbital / shell and the electrons are less shielded and more strongly attracted. Silicon has a slightly higher than expected third ionisation energy because the electron being removed is from the 3s orbital which experiences a greater effective nuclear charge than electrons in the 3p orbital which are lost in the case of phosphorus to argon.