Samples Containing Silver

General Information

Occurrence — Silver (Ag) is the second most abundant of the Group 11 elements (Cu, Ag, Au), which are also known as the "coinage metals." With an abundance in the earth's crust of 0.08 ppm, Ag is about 20x more abundant than Au and several orders of magnitude less abundant than Cu. Silver is found in nature as the free metal, as well as bound to insoluble counter ions, with sulfide being the most common. In Peru, lumps of the metal as large as 800 lbs. have been found. Its occurrence as the free metal is the chief reason it has been known and used by humans since prehistoric times. Ag has been used as a standard of value, where its earliest use predates all historical records. In modern times, copper is added to silver coins to increase their durability (10% Cu in the USA and 7.5% in Britain). In addition to the metal, Ag is also found in minerals as a bromide, chloride, iodide, arsenide, antimonide, and sulfide (plus other thio-salts). The more important minerals are argentite, hessite, proustite, pyrargyrite, cerargyrite, and native silver. In addition, Ag is commonly found as a trace component in copper and lead ores, from which significant Ag production occurs.

Uses — Pure Ag⁰ is the whitest of the metals, giving a beautiful luster when polished. Like the other Group 11 elements, Ag⁰ exceeds most elements in malleability, ductility, and heat conductivity, and exceeds all elements in electrical conductivity. It is harder than Au but softer than Cu. It alloys easily with Cu and Au and many other elements to form a wide variety of alloys. Silver's Group 11 cousins (Cu, Au) are similarly known for their usefulness as coin metals. Ag has a broad range of applications, but is best known for its use in tableware (sterling silver), electronics, coinage, medicine, photography, dentistry, and jewelry.

Chemical Properties — Silver has +1, +2, and +3 oxidation states. The most stable and common oxidation state for Ag is +1. The trace analyst is unlikely to encounter anything other than silver in the +1 oxidation state (Ag⁺), since this is the only stable oxidation state produced when acid digestions and other common sample preparation chemical operations are performed. Ag²⁺ is very unstable, requiring very strong oxidizing agents; i.e., the reaction Ag²⁺ + e^- = Ag⁺ has a reduction potential of +1.97 v vs. N.H.E. in acidic media. Ag²⁺ in nitric acid is prepared from the reaction of Ag⁺ with ozone (O₃). The formation of the +2 state in alkali is possible at much lower energies, where the reduction potential is only +0.57 v vs. N.H.E., but preparations such as caustic fusions that remain at a high pH for measurement are not common or recommended for trace Ag determinations (although if Ag²⁺ is present, it is likely to present fewer problems than the +1 state, as you are about to see).

Silver has several chemical properties that are noteworthy:

• Ag⁺ in acidic media reacts with Cl^- at comparable concentrations to form insoluble AgCl, and with excess Cl^- can form soluble AgCl₂^-. All silver halogen salts are reduced to the metal when exposed to light (photoreduction). The metal tends to deposit on the container surface. Systematic errors due to loss and memory/contamination tend to be most problematic for the trace analyst.

• The principal hydrolysis product is brown Ag₂O, which is formed when alkali is added to a silver Ag⁺ solution in dilute HNO₃ and water. Surprisingly, the solid hydroxide, AgOH, is unstable, existing only a short time as a transient compound.

2Ag⁺ + 2OH^- → 2AgOH → Ag₂O↓ + H₂O

• Silver carbonate does not precipitate in air-saturated solutions.

• Silver does not form polynuclear silver-hydroxide compounds.

• Silver is easily reduced to the metal, and electroplating is a common practice on both an analytical and commercial scale.

• Ag⁺, which is the only cationic form of silver in dilute nitric acid, forms more insoluble salts than any other metal.

The following is a summary of solubilities for common silver salts:

Solubility of Common Silver Salts at Room Temp. (~22 °C)

| Salt | Solubility in g/100g H₂O | |---|---| | Acetate | 1.04 | | Arsenate | 0.085 | | Arsenite | 0.00115 | | Borate | 0.905 | | Bromate | 0.196 | | Bromide | 0.014 | | Carbonate | 0.105 | | Chloride (most common) | 0.0154 (~1.4 ppm Ag⁺ and 0.5 ppm Cl^-) | | Chromate | 0.00256 | | Cyanide and thiocyanate | 0.022 | | Ferricyanide | 0.066 | | Fluoride | 172 | | Iodate | 0.00503 | | Iodide | 0.028 | | Nitrate | 216 | | Oxalate | 0.00378 | | Oxide | 0.00248 | | Phosphate | 0.064 | | Sulfate | 0.83 | | Sulfide | 0.0174 | | Tartrate | 0.0201 |

The Big Problem Is Chloride — The use of nitric acid and/or HF is preferred for preparation of samples for Ag analysis. Solutions of Ag in either acid are stable for extended periods. However, trace levels of HCl or Cl^- must be eliminated; otherwise, a fixed error due to AgCl precipitation will result. With ≥1.4 ppm Ag concentrations, even 0.5 ppm levels of chloride will be a problem. Avoiding these levels of chloride in a trace metals laboratory where HCl is a common and necessary reagent is very difficult, if not impossible, from a practical point of view. Many analysts experience low Ag recoveries when working in HNO₃, due to trace chloride contamination. Although analysts are aware of the problems with chloride, they are puzzled because no AgCl is observed. However, these low AgCl levels are difficult to spot and have likely already photoreduced onto the container walls, i.e.:

AgCl (or AgCl₂^-) + hv = Ag⁰↓ + Cl^-

The most common solution to this problem is to work in HCl, where the following reaction occurs:

AgCl↓ + Cl^- = AgCl₂^- (soluble)

You still have the problem with photoreduction to the metal, but elimination of light exposure is considered more achievable than elimination of trace chloride.

Consequently, most analysts are forced to work in a matrix containing HCl levels of 1% v/v or greater. If HCl must be included in your analytical sample preparation, which is often the case, then these tips should prove helpful.

1. If the sample preparation requires the use of HCl, keep the HCl content high (2-10% v/v).
2. Keep the Ag⁺ concentration low (0.5 µg/mL).
3. If the Ag⁺ concentration must go above 10 µg/mL, the HCl concentration must increase to >10% (v/v). Add HCl generously.
4. Keep the exposure to light at a minimum.
5. LDPE bottles wrapped in tape are a better option than HDPE or amber HDPE, due to impurities from the amber pigment and the HDPE. There is sufficient chloride from the catalyst residue in the HDPE to cause a problem.
6. The solution should be kept for no more than 5 months.

When working with HCl, it cannot be overemphasized that a high HCl content (>1% v/v) is needed to keep the Ag in solution as the AgCl₂^- anionic chloride complex. In short, keep the HCl concentration high and the Ag concentration low. And remember that solutions containing suspended AgCl and/or the AgCl₂^- anionic chloride complex are photosensitive. The Ag⁺ will undergo photoreduction to the metal (Ag⁰). When intentionally working in HCl, minimize exposure to light.

At Inorganic Ventures — Inorganic Ventures uses metallic Ag⁰ shot as the starting material, with a documented purity that is confirmed to be 99.999+% using ICP-MS and ICP-OES. The pure metal is dissolved in dilute, high-purity HNO₃. Impurities for each lot (measured using ICP-MS and ICP-OES) are reported on the Certificate of Analysis. The most common matrix for Ag CRMs manufactured at IV is HNO₃.

Acids other than HNO₃ require the use of an oxidizing agent, such as H₂O₂ (trace metals grade, unstabilized), before dissolution can be effected. On special occasions, high-purity AgNO₃ is used, which avoids certain problems arising from sub-oxide impurities left from HNO₃ acid. The solubility of AgNO₃ in water is 216 g/100 g H₂O at 20 °C, making it the most soluble (yields the most Ag per gram of salt) of the common Ag salts. Complications that occur from reactions with other solution components with easily reduced metals are avoided by heating the freshly prepared Ag in HNO₃ solutions to rid them of the nitrogen sub-oxides.

Photo 1. Taken after 1 day of exposure; all is clear.

However, after 13 days it was observed that a precipitate had formed in the open beaker (Photo 2), but the solution in the sealed bottle was still clear.

Photo 2: Taken after 13 days of exposure. There is no indication of a ppt in the bottle of sealed 1000 ppm Ag after sitting in the glove box with an open bottle of 50% (v/v) HCl.

The risk of contamination is low for Ag. However, trace analysts must be made aware of the high risk of contamination from used laboratory glassware, and the problems from HCl contamination. 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 subsampling, see Sampling and Subsampling.

Detailed handling information related to Ag-containing solutions, as well as suggestions for ICP analyses of silver, may be found by clicking on the Ag element symbol at our Interactive Periodic Table.

The Metal and Alloys

Metal — HNO₃ is the best "solvent" for Ag⁰. HNO₃ reactions will vary depending upon dilution with water, and when very dilute, the HNO₃ will lose its oxidizing ability. A 2:1 dilution in water is recommended, where the following reaction occurs:

Ag⁰ + 2HNO₃ ↔ AgNO₃ + NO₂↑ (brown fumes) + 2H₂O

Ag⁰ is soluble in hot 80% (v/v) H₂SO₄ in the presence of air (O₂).

Alloys — Many alloys can be dissolved using 2:1 HNO₃/H₂O. Alloys containing the "HF" elements will dissolve in a 1:1:1 mixture of concentrated HNO₃/HF/H₂SO₄.

Oxides, Minerals and Ores

Oxides — The oxides are soluble in dilute acids. Dilute HNO₃ is most popular. If the oxides are somewhat resistant to dilute acid attack, then a fusion with sodium carbonate should be used first.

Ores — An acid dissolution can be performed in HF+HNO₃+HClO₄. Ores can also be fused with disulfate.

Organic Matrices

Ashing of organic materials, foodstuffs, plant, and blood and sewage sludge, as a preliminary decomposition step, can be expedient for samples containing Ag. If ashing is used, it is suggested to keep the temperature low (400 to 450 °C max.) and to use an ashing aid, such as high-purity sodium carbonate. The dry ashing of samples results in the formation of Ag⁰ when organics are present, and refractory metal oxides of many elements, which can be very difficult to dissolve in dilute acids (Ag⁰ requires an oxidizing agent); therefore, ashing aids, such as Na₂CO₃, are strongly encouraged. Acid digestions using nitric, perchloric, and sulfuric acids are suggested. Wet ashing uses sulfuric acid as an ashing aid. Do not use ashing temperatures above 450 °C. If sulfuric acid is added to a petroleum sample, then slowly heat on a hot plate until foaming stops and a char is produced, i.e., wet ashing is very time-consuming, but it is a common practice in the petroleum industry. For more on ashing, please see the following paper: Ashing

Silver is listed in the scope for EPA Methods 3050A and 3050B (Open Vessel Acid Digestion) and 3051 and 3052 (Microwave Assisted Acid Digestion), and these methods are suggested for environmental samples (sediments, sludges, soils, and oils).

Samples containing mid to low ppm levels of Ag can be digested with nitric/perchloric acid. Only use trace metals grade acids, due to contamination issues. For more detailed information about acid digestions of organics, please see the following article: Acid Digestions of Organic Samples