Group 16 Flashcards
Chalcogen elements?
As with the other groups the metallic character of the element increases down the group oxygen and sulfur are typical non metals, selenium and tellurium have some metalloid properties and polonium is normally regarded as a metal.
Oxidation state of +6 but never observed for oxygen, maximum oxidation state of +2. Oxygen forms compounds with virtually all other elements with the exception of the group 18 elements. Oxidations states of +4 and +6 are possible for sulfur and the heavier chalcogens
Allotrope of oxygen?
Dioxygen and this exists as a colourless odourless gas at room temperature, pale blue colour. Industrially O2 prepared by fractional distillation of air or the electrolysis of water or from decomposition of either potassium chlorate KClO3 or hydrogen peroxide H2O2. Both catalysed by MnO2
2KClO3 —-> 2KCl + 3O2
2H2O2 —-> 2H2O + O2
chemistry of O2 dominated by oxidising ability, combustion corrosion and respiration all reactions where O2 is an oxidising agent. Ozone is another allotrope O3, pale blue gas formation of ozone from O2 is extremely exothermic 3O2 —> 2O3 powerful oxidising agent used in water purification
Sulfur?
Found near volcanoes, occurs in deposits where it has been formed by bacteria, many minerals are sulphides valence PbS, sphalerite ZnS or sulphate such as gypsum. Produced industrially as a byproduct from the extraction of copper form sulphide ores and from cleaning natural gas and oil to produce low sulfur fuels.
Allotropes of sulfur?
Many allotropes of sulfur, more than for any other elements, most stable at room temperature is alpha sulfur which contains cyclic S8 molecules, can be converted by heating into other allotropes. Beta sulfur also contains S8 rings, rings containing between 6 and 20 sulfur atoms are also known, one o the reasons for the diversity of the allotropes is the tendency of sulfur to catenate forming chains and rings through formation of S-S bonds. Sulfur is used industrially in vulcanisation of rubber, disulphide links form bridges between polymer chains and makes the rubber harder and more durable
Anions of group 16 elements?
Form oxides O2- ion and sulphides containing the S2- ion. Ionic compound containing the peroxide ion O22- and the superoxide ion O2- are also well known. The X2- ions become less stable going down the group 16. This is because the larger anions held their electron less tightly so they are more easily oxidised. Sulfur forms a range of polysulfides containing Sn2- ions these are anions of the catenated hydrides H2Sn. They’re are prepared by the reaction between a sulphide and sulfur
2Cs2S + S8 —> 2Cs2S5
Sulfur oxides?
Sulfur dioxide SO2 and sulfur trioxide SO3, sulfur dioxide is a colourless toxic gas that is formed when sulfur burns in air
S8 + 8O2 —> 8SO2
Sulfr dioxide is very soluble in water, the SO2 molecules react with water to give an acidic solution containing ions derived from sulphurous acid
SO2 + H2O HSO3- + H3O+
HSO3- H2O SO32- + H3O+
Why no allotropes for N4 and P4?
Triple bond for N is 3 times that for a single bond so N-N allotropes are unstable with respect to N2, P4 can also be shown to be more stable than P2 in the same way
Atom size and bond strengths in first row?
Atom size increases as we go down a group so atom-atom contact increases, as atom size increases so does the size of the orbitals and as atoms increase sin size the distance between them increases and bond strength decreases. For atom single bond strengths first row anomaly for C-C much higher since no lone pairs so no electron electron repulsions. First row anomaly for N, O and F explained by presence of lone pairs and the small size of the atoms, strong repulsions. For B and C all electron spurs and bond pairs so repulsion sis limited to bond pair bond pair repulsions, for P, S and Cl bigger atoms means larger bonds and less repulsions by diffuse orbitals, for O2 the pi component of the bond is greater than twice the sigma bond the opposite is true for S
Acidity?
Bonds are weaker as we go down a group so reduce stability and increase acidity (tendency to form H+) NH3, OH2 and HF all have H bonding. As we do down a group the energy difference between s and p orbitals increases this means there is less so mixing and bonds become more p like in character
Sulfur trioxide?
Sulfur dioxide reacts slowly with oxygen to form sulfur trioxide SO3 this reaction is important to make sulphuric acid and is catalysed by vanadium oxide
2SO2 + O2 —> 2SO3
The equilibrium constant decreases rapidly with increase temperature so the reaction is carried out at a relatively low temperature between 435 and 635 degrees. At lower temperature the reaction becomes too slow to be economic
Sulfur trioxide is a volatile white solid at room temperature the gaseous molecules are trigonal planar in shape but in the solid some of the SO3 molecules trimerise to form S3O9 rings, it reacts with water to form sulphuric acid
SO3 + H2O —> H2SO4
This reactions is strongly exothermic, oleum is created which is a picture of polysulfuric acid.
Sulfur oxacides?
Sulfuric acid H2SO4
Disulfuric aid H2S2O7
Sulfurous acid H2SO3
Sulphuric acid?
strong acid, oxidising agent and a dehydrating agent
C12H22O11 —> 12C + 11H2O
Test for sulfate ions add barium chloride and a dense white precipitate of barium sulfate is formed
Sulfurous acid?
Cannot be isolated as a pure compound but its salts, sulphites containing the SO32- ion and hydrogen sulphites containing the HSO3- ion are stable compounds sulphites are formed by reacting dissolved SO2 with a base such as sodium hydroxide
SO2 + 2OH- —> SO32- + H20
sulfur dioxyde and sulfites used a preservatives. Sulphites good reducing agents really oxidised to sulphates
Thiosulfate ion?
S2O32- prepared by boiling sulfur in an alkaline solution of sodium sulphite
8SO32- + 28 —> 8S2O32-
Thiosulfate ions are good reducing agents and are oxidised to tetrathionate ions S4O62-
Commonly used in redox titrations.
Oxygen fluorides?
Only halides possible for oxygen is fluorides only element more electronegative than itself, otherwise we get a chlorine oxide etc.
OF2 and O2F2 have similar structures to water and hydrogen peroxide. OF2 is a toxic yellow gas formed in the reaction between fluorine and hydroxide ions
2OH- + 2F2 —> OF2 + 2F- + H2O
Oxygen fluorides are only compounds of oxygen in which the oxidation state is positive which means that the O-F bonds are polarise in the opposite direction from the O-Cl bonds in Cl2O
Bonding of F to O2 molecules is by overlap of 2p on fluorine with the pi* on O2, fluorine is more electronegative than oxygen so there is little increase in pi* electron density thus preventing the double bond of O2, weirdness of structure explained by MO theory
