Lecture 4. Oxides and Oxoacids

Oxides of fluorine, chlorine, bromine and iodine [1]

The shape of the OF₂ molecule has been compared with HOF and H₂O and the three give bent structures where the bond angles range from 97 to 104.5 °.

Cl₂O is best prepared by treating fresh yellow mercuric oxide with Cl₂ gas:
2 Cl₂ + 2 HgO → Cl₂O + HgCl₂·HgO

It can also be prepared by reaction of Cl₂ gas with moist sodium carbonate, Na₂CO₃.
2Cl₂ + 2Na₂CO₃ + H₂O → Cl₂O + 2NaCl + 2NaHCO₃

Dichlorine monoxide is very water soluble (a saturated solution at -9.4 °C contains 143.6 g Cl₂O per 100 g H₂O) and it hydrolyses to hypochlorous acid.
Cl₂O + H₂O → 2HOCl

Cl₂O

Chlorine dioxide (ClO₂) was mentioned in the first lecture as being an important industrial chemical for bleaching flour and wood pulp and for water treatment. Over 95% of the chlorine dioxide produced in the world today is made from sodium chlorate:
2NaClO₃ + 2NaCl + 2H₂SO₄ → 2ClO₂ + Cl₂ + 2Na₂SO₄ + 2H₂O

ClO₂ was first discovered by Humphry Davy (1778-1829) in 1811 and was the first of the chlorine oxides to have been isolated. The yellow paramagnetic gas explosively decomposes into chlorine and oxygen and the decomposition is initiated by light. Thus, it is never handled in concentrated form, but is almost always used as a dissolved gas in water in a concentration range of 0.5 to 10 grams per liter. Its solubility increases at lower temperatures: it is thus common to use chilled water (5 °C) when storing at concentrations above 3 grams per liter. In many countries, such as the USA, chlorine dioxide gas may not be transported at any concentration and is almost always produced at the application site using a chlorine dioxide generator. Chlorine dioxide is less corrosive than chlorine and superior for the control of legionella bacteria. It is more effective as a disinfectant than chlorine in most circumstances against water borne pathogenic microbes such as viruses, bacteria and protozoa.

The oxidation number of the Cl is considered to be +3 and Pauling originally proposed that resonance structures be used to account for this. An alternative explanation by Brockway extended this to resonance between 3 electron 2 centre bonds. For many years it was claimed that a dimer was not formed and considering that ClO₂ is a radical this might have been one mechanism to provide some stability by using up the odd electron in a 2 electron 2 centre bond. In 1992, Rehr and Jansen conducted magnetic susceptibility and single crystal studies at various temperatures and it was found that below about 170 K the sample became diamagnetic and did indeed form a dimer. They noted that to avoid explosions all equipment had to be scrupulously cleaned with aqua regia prior to use and temperatures kept low.

ClO₂

The oxides of bromine are less numerous than for chlorine and they are generally unstable at room temperature decomposing to the elements.

The oxides of iodine are the most stable of the halogens and the first of these to be prepared was I₂O₅ which was independently discovered by Joseph Louis Gay-Lussac (1778-1850) and H. Davy in 1813 although the structure was not determined until 1970. The white hygroscopic crystals are very soluble in water and commercial I₂O₅ has been found to consist of HI₃O₈ or I₂O₅.HIO₃. Likewise crystals of commercial iodic(V) acid, HIO₃, were investigated using single-crystal and powder X-ray diffraction and the crystals turned out again to be HI₃O₈ or HIO₃·I₂O₅.

partially staggered conformation of I₂O₅

I₂O₅ is one of only a few chemicals that is capable of oxidising carbon monoxide rapidly and completely at room temperature:
I₂O₅ + 5 CO → I₂ + 5 CO₂
This reaction has been used as the basis of an analytical method for sampling CO in the atmosphere. I₂O₅ can be used for the oxidation of a number of other small molecules including: NO, C₂H₄, SO₃, H₂S.

Oxoacids and their salts

Numerous oxoacids of the halogens exist and chlorine is able to form a complete series of such acids ranging from HOCl to HClO₄ where the formal oxidation state is +1,+3,+5 and +7. [1]

HOF has been isolated by reaction of difluorine gas with ice at 230 K but it does not ionise in water or give stable salts and on warming to room temperature it decomposes to HF.
2 HOF → 2 HF + O₂

The hypohalous salts formed from the heavier halogens are all weak acids with the hypochlorites being important industrial reagents.
Bleaching powder is a mixture of CaCl₂, Ca(OH)₂ and Ca(OCl)₂ and is manufactured by the reaction of dichlorine on Ca(OH)₂. The sodium salt is a well known disinfectant and bleaching agent

Chlorous acid, chloric acid and bromic acid cannot be isolated as pure compounds but exist in aqueous solution and are prepared from the barium salts.
Ba(ClO₂)₂ + H₂SO₄(aq) → 2 HClO₂(aq) + BaSO₄(s)
and Ba(XO₃)₂ + H₂SO₄ → 2 HXO₃ + BaSO₄ (X= Cl, Br)

The structure below shows trigonal pyramidal HIO₃ units and these are connected by extensive H-bonding. In aqueous solution it is a reasonably strong acid with pKa ~ 0.77.

• as a pesticide. In Finland it was sold under the trade name Fegabit.
• as an oxidizing agent,
• to prepare oxygen,
• as a disinfectant,
• in safety matches, and
• in explosives and fireworks.

The materials used need to be of high purity since even small amounts of impurities can lead to explosions and it is considered good practice to test a new batch of potassium chlorate with a small reaction first (less than 1 g!) Even commercial oxygen generators have been known to explode and cause problems such as at least one plane crash and a fire on the space station MIR.[10]

Both potassium bromate and iodate are used in volumetric analysis and in particular the iodate is a primary standard being available in high purity and with excellent stablity. In the standardisation of thiosulfate for example potassium iodate is used a source of diiodine.
[IO₃]^- + 5 I^- + 6 H⁺ → 2 I₂ + 3 H₂O
Iodate is used in Andrew's titrations for determination of hydrazine and in the determination of I^- in SnI₄ in the CHEM3101 laboratory programme.

Several different I(VII) oxyacids are known including periodic acid (HIO₄) and orthoperiodic acid (H₅IO₆). In the latter case there is extensive H-bonding present in the solid state resulting in a 3D network. Wherease perchloric acid exists as discrete monomeric units HIO₄ is again extensively H-bonded through cis- edge-sharing octahedra.

The reactions involving the various oxyacids in aqueous solution is complicated, as seen by the following equilibria:
[H₃IO₆]^2- + H⁺ → [IO₄]^- + 2H₂O
2[H₃IO₆]^2- → 2[HIO₅]^2- + 2H₂O
2[HIO₅]^2- → [H₂I₂O₁₀]^4- → [I₂O₉]^4- + H₂O

Count Friederich von Stadion is credited with the preparation of the first perchlorate compound in 1816 which he obtained by mixing potassium chlorate with concentrated sulfuric acid and allowing it to stand for 24 hours with frequent agitation. The water-insoluble residue was potassium perchlorate. Perchloric acid is the only oxoacid of chlorine that can be isolated and in the vapour phase the Cl-O bonds are not equivalent unless deprotonated when they become the same length.

HClO₄

Perchlorates in groundwater

Texas chemists have shown evidence for a natural atmospheric origin of perchlorate. For many, perchlorate is almost synonymous with the rocket fuel that has found its way into U.S. groundwater and surface water and subsequently into foods such as lettuce and milk. Perchlorate taints drinking water in 35 US states at levels of at least 4 ppb. When groundwaters of a 60,000-sq-mile region in Texas and New Mexico were sampled, > 80% of the wells returned positive and unexpectedly high levels of perchlorate and it was found even in samples that were unlikely ever to have been exposed to synthetic contamination. The perchlorate levels in many of the Texas samples were more than 20 ppb, and some were as high as 60 ppb. However, there were regions such as the southern high plains (Texas Panhandle) where there was no clear historical or current evidence of the extensive presence of rocket fuel or Chilean fertilizer sources. The occurrence of easily measurable concentrations of perchlorate in such places was difficult to understand. In the southern high plains groundwater, perchlorate was better correlated with iodate, known to be of atmospheric origin, compared to any other species. They showed that perchlorate was readily formed by a variety of simulated atmospheric processes. For example, it could be formed from chloride aerosol by electrical discharge and by exposing aqueous chloride to high concentrations of ozone. They further reported that perchlorate was present in many rain and snow samples, strongly suggesting that some perchlorate was formed in the atmosphere and a natural perchlorate background of atmospheric origin existed.[4][5][6]

Although the scientists were fairly certain that the perchlorate in the Texas wells was of natural origin, they still needed more evidence to prove it. Researchers at Louisiana State University (LSU) and Oak Ridge National Laboratory (ORNL) developed a new stable-isotope ratio technique that could provide the missing piece of information.

To distinguish between natural and synthetic perchlorate it is important to recognise that perchlorate is a non-volatile oxyanion and that once formed it does not exchange oxygen atoms with those in the ambient environment. This means its oxygen ratio signature is fixed. So by looking at both the ¹⁷O/¹⁶O and ¹⁸O/¹⁶O ratios it is possible to tell whether it is from a natural source or not.

In this approach, all three stable oxygen isotopes are measured after a thermal decomposition method generates O₂ from perchlorate crystals. Measuring just ¹⁸O/¹⁶O won't give you the answer, because ¹⁸O/¹⁶O ratios for anthropogenic and atmospheric perchlorate overlap. But if you measure ¹⁷O/¹⁶O at the same time and compare the two ratios, it is possible to determine its origin. Most oxygen-containing compounds on earth have a correlation between ¹⁸O/¹⁶O and ¹⁷O/¹⁶O. Any deviation from this relationship is referred to as the ¹⁷O anomaly, or Δ¹⁷O. The results of studies on Chilean perchlorates from the Atacama desert caliche deposits found that Δ¹⁷O was between 4 and 10 whereas for man-made perchlorates the value was generally less than -0.2.[7]

In the human body, perchlorate inhibits the uptake of iodide by the thyroid and thus may lower the amount of thyroid hormone in the body. Insufficient levels of this hormone can cause permanent neurological damage in children.

The changes in perchlorate concentrations in surface water adjacent to a site of fireworks displays from 2004 to 2006 was monitored. Preceding the fireworks displays, perchlorate concentrations in surface water ranged from 0.005 to 0.081 µg/L, with a mean value of 0.043 µg/L. Within 14 h after the fireworks, perchlorate concentrations spiked to values ranging from 24 to 1028 times the mean baseline value. A maximum perchlorate concentration of 44.2 µg/L was determined following the July 4th event in 2006. After the fireworks displays, perchlorate concentrations decreased toward the background level within 20 to 80 days, with the rate of attenuation correlating to surface water temperature.[8]

The Pentagon, NASA, the US Department of Energy, and defense contractors - who could face expensive cleanups of perchlorate contamination - vigorously objected when the EPA proposed a 1-ppb limit for drinking water. The US military suggested a limit of 200 ppb.

Aqueous Solution Chemistry

The diagrams below show the variation of redox potentials with acid and base, It should be noted that the couples ½X₂/X^- are independent of pH and together with the estimated value for F₂ indicate a steady decrease in oxidising strength moving down the group.
F₂ (+2.866 V) > Cl₂ (+1.358 V) > Br₂ (+1.066 V) > I₂ (+0.536 V)

The reason the potential for difluorine must be estimated is that it is too strong an oxidant and will oxidise water to O₂ so the potential can not be easily measured!
Given that the value listed for dichlorine is also greater than what is needed to oxidise water (+1.229 V) this should happen as well however this is not the case since there is a slow kinetic step in the oxidation of water. Note as well that this is the same reason why aqueous solutions of potassium permanganate can be prepared, although they are not stable over time.

The majority of the couples are pH dependent so that an increase in pH causes a dramatic lowering of the E° value. This is expected since if we consider an example where BrO₃^- is converted to ½Br₂ by the reaction:
BrO₃^- + 6H⁺ + 5e^- → ½Br₂ + 3 H₂O (E° = +1.495 V)

the equilibrium constant includes a term with [H⁺]⁶ so that if the pH is changed from pH=1 to pH=14 then H⁺ effectively changes by 10^-14 and the estimated difference can be seen in E° as:
ΔG° = - RTln(K)
ΔG° = - nE°F

from which just looking at the effect of pH is ~(RT/nF)ln([H⁺]⁶)
or roughly (0.0592/5) * 14 * 6
that is an effect of ~+0.99 V
Experimentally the measured E° for
BrO₃^- + 3H₂O + 5e^- → ½Br₂ + 6 OH^- (E° = +0.519 V)
so estimated value is +1.495 - 0.99 = +0.51 V and found is +0.519 V showing excellent agreement.

Redox processes in acid solutions

Redox processes in basic solutions

References:

Return to Chemistry, UWI-Mona, Home Page