Sample Preparation Guides
Iron is the fourth most abundant element in the earth's crust after oxygen, silicon and aluminum. Iron has been known by man back to prehistoric times and was used for making weapons and tools. During the middle ages high grade steel was manufactured in Arabia and Spain. Although pure Fe finds little commercial use, the addition of a few elements makes it arguably the most important material used in the industry. Fe is also one of if not the most common elements requiring analysis by analytical laboratories that have a variety of sample preparation and measurement techniques available.
Iron is found in the +2 and +3 oxidations states and is commonly found as iron (II) oxide (FeO); iron (III) oxide (Fe2O3); magnetite (Fe3O4); limonite (Fe(OH)6·Fe2O3); siderite (FeCO3); pyrrhotite (FenSn) and a large number of combinations with many other elements such as chromite (FeO·Cr2O3). The chemical resistance to bringing about a solution varies widely with the freshly precipitated oxides/hydroxides dissolving with little effort and ignited oxides and combinations like the chromite requiring a much greater effort. The techniques proposed here are applicable to this wide spectrum of sample types.
In general, Fe compounds are not easy to dissolve and their chemical behavior is not that predictable. Fe exists in oxidation states ranging from +6, +5, +4, +3, +2, +1 and 0 with the +3, +2 and 0 states being by far the most common. Consequently, the chemistry of iron addressed in this chapter will be limited to these three oxidation states. The literature addressing the analytical chemistry of Fe is extensive. This discussion will present methods that have been performed for intended analysis by ICP-OES, ICP-MS and AAS but that should also be applicable to Fe analysis using colorimetric, titrimetric, gravimetric or electrochemical methods of analysis. Questions concerning the applicability of a given preparation technique to an intended measurement methods are welcome.
In acidic media Fe+3 is thermodynamically favored and is the state that results from most sample preparations intended for ICP measurement. In some preparations Fe will go into solution as the Fe+2 but will slowly be oxidized by dissolved oxygen to Fe+3 unless the solution is purged and blanketed with high purity nitrogen. For example, this is the case when Fe metal is dissolved with HCl. The following reduction potentials are very helpful in explaining the chemistry of Fe and its oxidation states (taken from 'Electroanalytical Chemistry, 2nd edition, James J. Lingane, Interscience Publishers, NY; 1970):
|1. Fe+3 + e = Fe +2||(0.771 volts vs NHE)|
|2. Fe+2 + 2e = Fe||(-0.440 volts vs NHE)|
|3. Fe (CN)6-3 + e = Fe (CN)6-4||(0.36 volts vs NHE)|
|4. FeO2- +e = FeO2-2 (10M NaOH)||(-0.68 volts vs NHE)|
The +2 oxidation state of Fe is much more stable in acidic than in basic media to oxidation to the +3 state. Standard solutions of Fe+2can be prepared but are more stable if air is excluded. When air exclusion is not possible then they should be restandarized each analytical day. Fe+2 hydrolyses only slightly before precipitation of the Fe(OH)2 and this precipitation does not begin until a pH of ³ 6.5. Fe+2 can be formed with non-oxidative sample preparations but with slow air oxidation to Fe+3. Fe+3 begins precipitation as the hydroxide at a pH of ~ 0.5. Unlike the +2 state, hydrolysis of Fe+3 ions involves the formation of polynuclear complexes such as Fe3(OH)4+5 . The form in which the Fe+3 hydroxide precipitates depends on conditions such as the pH, temperature, and anions present in solution and is evidenced by a wide variety of shades of brown to orange. The formation of complexes with both the +2 and +3 oxidation states results in an ability to maintain stable solutions with increasing pH and also results in a lowering of the reduction potentials. For example, a 0.1M Fe/EDTA complex is stable at pH values of 11.5 and that of 0.1M cyanide is stable at a pH >12 which is rather remarkable considering the insolubility of Fe(OH)3. When Fe(OH)2 is precipitated in the presence of air it turns black immediately due to the formation of the +3 oxidation state and a magnetite like mixed oxidation state material.
Fe(II) salts, in crystals and in solution, have a pale green color. Solutions of the salts are slightly acidic. Fe (III) salts form solutions with a brownish-yellow color. The Fe (II) sulfate hepta hydrate (48), chloride (62), bromide (117) and nitrate hexahydrate (77) are readily water soluble (solubility in g/100 g water shown in parenthesis) with the Fe (III) salts of these same anions showing similar solubility. The carbonate (Fe+2), oxides, and sulfides of both oxidation states are insoluble in water and the fluorides are only slightly water soluble. Solutions of Fe+3 when treated with alkali hydroxides all form Fe(OH)3 which is a reddish brown solid that does not redissolve in an excess of the hydroxide unlike its neighbor Cr or Co, Ni and Zn which will dissolve in excess NH4OH. Fe(OH)3 is used in trace analysis as a gathering agent because its gelatinous property and structure bring about co-precipitation of trace metals in the same solution.
Sampling and Handling
The measurement of Fe is required for a wide range of samples including ores and minerals, soils, sludges, water (both drinking and waste), biological, agricultural, metallurgical, and industrial samples. There is a great risk of contamination when any sample handling apparatus/device is made of or has come in contact with stainless steel. In addition, since Fe is so prevalent in the earth's crust the air (particulate dust) must be considered as a contamination route. The following considerations apply:
- Many tools that pulverize, mix, cut, etc. contain Fe. Attempt to use devices made of ceramics, silica/quartz, and polymers where possible.
- Many polymers such as Teflon have been exposed to stainless steel (cutting/pulverizing) prior to molding. New plastic and Teflon containers should always be acid leached to remove possible contamination from the stainless steel elements.
- The collection of biological samples are also at risk of contamination due to the very low (ppb) levels of Fe sought. The use of steel needles and scalpels, or any metallic object that may contain Fe, should be avoided.
The risk of contamination is great for Fe when alloys, steels and grinding equipment are used in some part of the sample collection or preparation. For more on sample contamination risks see chapters 8, 9 and 10 of the Inorganic Ventures 'Trace analysis Guide':
For general information on sampling and sub-sampling see: http://www.inorganicventures.com/sampling-and-subsampling
The metal and Alloys
Metal – Iron metal dissolves in HCl and dilute H2SO4, forming Fe+2 and liberation of H2 See following equation:
°Fe + H2SO4 = FeSO4 + H2
Concentrated cold sulfuric acid had no action, but if hot, SO2 is evolved and Fe+3 formed as per the following equation:
2 °Fe + 6 H2SO4 = Fe2(SO4)3 + 3 SO2 + 6 H2O
Fe metal will not dissolve in cold concentrated nitric acid but when heated to near boiling it dissolves vigorously with the evolution of brown NO2.
Alloys – The use of nitric + HCl is most commonly used for Cr/Ni and conventional steels and the addition of water to this mixture is helpful in increasing the sample size and will allow for the preparation of the Fe/Ni high temperature alloys as well as conventional steels. When an element is present in the steel/alloy such as Si, Ti, Ta, Zr, W etc. the addition of HF is essential and the presence of HCl, although often used, becomes less critical. A typical preparation for a Zr steel (0.1 grams of alloy/steel + 20 mL conc. HNO3 + 10 mL 1:1. HCL+ 5 mL HF) is somewhat different than that needed for a ferrosilicon (0.1 gram sample + 20 mL conc. HNO3 + 5 mL 40% HF). With the understanding that just about any alloy of Fe will dissolve in a mixture of nitric/HCl/HF with the addition of water generally increasing the amount of sample that can be dissolved but with too much water quenching the entire reaction. An impressive mixture is a 1:1:1 mixture of HNO3 + HF + H2O and is used to dissolve high-temperature alloys of Fe/Co/Cr/Ti/AL/Mo/Ta/Hf (use 10 mL mixture/ 0.2 grams sample).
Oxides, Minerals and Ores
Oxides and hydroxides – Fe2O3 after ignition to temperatures at or above typical ashing temperatures (450 – 1000 °C) is rendered insoluble or at best very slightly soluble in acids. HCl is the best solvent but dissolution is very slow at best. If the oxide is heated with the alkali oxides or carbonate the structure will be 'opened out' and dilute acid will be sufficient. Samples containing other group I or II elements will result in and ash that is soluble in 1:1 HCl for example. Freshly precipitated iron oxides and hydroxides are all readily soluble in dilute acids. Although HCl is the best acid, the use of dilute nitric acid is common. HCl works better due to the ligand properties not present with the nitrate anion. If the sample preparation requires that the sample be ashed then the addition of an ashing aid such as sodium, magnesium or calcium carbonate makes the dissolution step much easier and this approach in compatible with Pt crucibles but is not recommended for alumina, quartz or porcelain crucibles. Many preparations of Fe containing samples use either lithium carbonate, lithium borate (meta or tetra), or sodium carbonate fusions and the ashing step is skipped all together. If the sample contains a lot of organic or combustible material then a programmable muffle furnace where the temperature is slowly ramped to 1000 °C (in the case of sodium carbonate as the ashing aid) is prudent.
Ores – The sample should be pulverized to pass an 80 to 100 mesh sieve. Sulfides and ores containing organic matter can be mixed with sodium carbonate and ashed in a Pt crucible at 500°C for 1 hour prior to bring to a temperature of 1000 °C as described below. If the ore has already been strongly ignited it can go directly to 1000 °C with sodium carbonate (0.25 to 1.0 gram sample + 20 grams sodium carbonate fused at 1000 C° for 10 to 30 minutes in a Pt crucible followed by dissolution with 1:1 HCl). When a sodium carbonate fusion is performed HF is not needed unless silica content is high (silicates formed from the carbonate fusion are soluble but do not heat the fuseate to dissolve since this will bring about the formation of polysilicic acid which is a gelatinous precipitate). This approach is applicable to sulfides, ores containing organic matter, oxides, red and brown hematites, magnetic iron ore, splathose iron ore, roasted pyrites, and iron ore briquettes.
Ashing of organic materials, foodstuffs, plant, and blood and sewage sludge as a preliminary decomposition step is not suggested for samples containing Fe unless an ashing aid such as sodium carbonate is used and the crucible material is Pt. If magnesium nitrate is used as the ashing aid then porcelain or quartz crucibles are acceptable. Conventional dry ashing procedures will result in the formation of a refractory form of iron oxide unless there is a significant amount of one or more elements from groups I, II, or III. If ashing is used it is suggested to keep the temperature low (400 to 450 deg C max) and to use an ashing aid such as high purity sodium carbonate or magnesium nitrate. If the sample is high in silica subsequent fusion of the ash with sodium carbonate is suggested or heating of the ash to fumes with sulfuric and hydrofluoric acids.
For more on ashing please see the following paper:
The acid digestion of organic samples containing Fe is also very common. Fe does not offer any unusual challenges over the other elements and, with the exception of a higher than average contamination risk, is very well behaved. There are numerous oxidative acid digestion procedures available that are tailored to the matrix.
For more on acid digestion please see the following paper: