Handling, Calculations, Preparation and Storage of Standards

Overview

This section contains basic information on the handling, preparation and storage of standards, as well as basic calculations and nomenclature.

Handling

Observing the following recommendations will save considerable time, money, and frustration:

1. Never put solution transfer devices into the standard solution. This precaution avoids possible contamination from the pipette or transfer device.
2. Always pour an aliquot from the standard solution to a suitable container for the purpose of volumetric pipette solution transfer, and do not add the aliquot removed back to the original standard solution container. This precaution is intended to avoid contamination of the stock standard solution.
3. Perform volumetric pipette solution transfer at room temperature. Aqueous standard solutions stored at "lower" temperature will have a higher density. Weight solution transfers avoid this problem, provided the density of the standard solution is known, or the concentration units are in wt./wt. rather than wt./volume.
4. Never use glass pipettes or transfer devices with standard solutions containing HF. Free HF attacks glass, but it is sometimes considered safe to use glass when the HF is listed as trace and/or as a complex. However, many fluorinated compounds will attack glass just as readily as free HF.
5. Don't trust volumetric pipette standard solution transfer. Weigh the aliquot of the standard taken. This can be easily calculated, provided the density of the standard solution is known. There are too many possible pipetting errors to risk a volumetric transfer without checking the accuracy by weighing the aliquot.
6. Uncap your stock standard solutions for the minimum time possible. This is to avoid transpiration concentration of the analytes, as well as possible environmental contamination.
7. Replace your stock standard solutions on a regular basis. Regulatory agencies recommend or require at least annual replacement. Why is this precaution taken, in view of the fact that the vast majority of inorganic standard solutions are chemically stable for years? This is due to the changing concentration of the standard through container transpiration and the possibility of an operator error through general usage (more info). A mistake may occur the first time you use the stock standard solution, or it may never occur, with the probability increasing with use and time. In addition, the transpiration concentration effect occurs whether the standard solution is opened/used or not, and increases with use and increased vapor space (transpiration rate is proportional to the ratio of the circumference of the bottle opening to vapor space).

Calculations

The concentration units for chemical standard solutions used for ICP applications are typically expressed in µg/mL (micrograms per milliliter) or ng/mL (nanograms per milliliter). For example, a 1000 µg/mL solution of Ca²⁺ contains 1000 micrograms of Ca²⁺ per each mL of solution, and a 1 µg/mL solution of Ca²⁺ contains 1000 ng of Ca²⁺ per milliliter of solution. To convert between metric concentration units, the following conversions apply:

The difference between ppm and µg/mL is often confused. A common mistake is to refer to the concentration units in ppm as a shortcut (parts per million) when we really mean µg/mL. One ppm is in reality equal to 1 µg/g. In similar fashion, ppb (parts per billion) is often equated with ng/mL. One ppb is in reality equal to 1 ng/g. To convert between ppm or ppb to µg/mL or ng/mL, the density of the solution must be known. The equation for conversion between wt./wt. and wt./vol. units is:

(µg/g) × (density in g/mL) = µg/mL

and/or

(ng/g) × (density in g/mL) = ng/mL

Therefore, if we have a solution that is 1000 µg/mL Ca²⁺ and know or measure the density to be 1.033 g/mL, then the ppm Ca²⁺ = (1000 µg/mL) / (1.033 g/mL) = 968 µg/g = 968 ppm.

When making dilutions, the following equation is useful:

(mL_A)(C_A) = (mL_B)(C_B)

For example, to determine how much of a 1000 µg/mL solution of Ca²⁺ is required to prepare 250 mL of a 0.3 µg/mL solution of Ca²⁺, we would use the above equation as follows:

(mL_A)(1000 µg/mL) = (250 mL)(0.3 µg/mL)

(mL_A) = [(250 mL)(0.3 µg/mL)] / (1000 µg/mL)

(mL_A) = 0.075 mL = 75 µL

Preparation

Weight ≠ Volume

Standard chemical solutions can be prepared to weight or volume. The elimination of glass volumetric flasks may be necessary to eliminate certain contamination issues with the use of borosilicate glass, or to avoid chemical attack of the glass. It is often assumed that 100 grams of an aqueous solution is close enough to 100 mL to not make a significant difference, since the density of water at room temperature is very close to 1.00 (0.998203 at 20.0 °C). Diluting/preparing standard solutions by weight is much easier. Still, the above assumption should not be made. The problem is that trace metals standards are most commonly prepared in water + acid mixtures, where the density of the common mineral acids is significantly greater than 1.00. For example, a 5% v/v aqueous solution of nitric acid will have a density of ~1.017 g/mL, which translates into a fixed error of ~1.7%. Higher nitric acid levels will result in larger fixed errors. This same type of problem is true for solutions of other acids, to a degree that is a function of the density and concentration of the acid in the standard solution, as described by the following equation (to be used for estimation only):

d_S = [(100 − %) + (d_A)(%)] / 100

Where:

• d_S = density of final solution
• % = the v/v % of a given aqueous acid solution
• d_A = density of the concentrated acid used

For example, let's estimate the density of a 10% v/v aqueous solution of nitric acid made using 70% concentrated nitric acid with a density of 1.42 g/mL.

d_S = [(100 − %) + (d_A)(%)] / 100 = [(100 − 10) + (1.42)(10)] / 100 = (90 + 14.2) / 100 = 1.042 g/mL

Acid Content

Another area of confusion is the expression of the acid content of the solution. We all agree that it is important to matrix match the standard and sample solutions, to avoid a fixed error in the solution uptake rate and/or nebulization efficiency, sometimes referred to as a matrix interference. If a solution is labeled as 5% HNO₃, what does this mean? If we take 5 mL of 70% concentrated nitric acid and dilute to a volume of 100 mL, then this is 5% HNO₃ (v/v), where the use of 70% concentrated acid is assumed. However, nitric acid can be purchased as 40%, 65%, 70%, and >90%. Therefore, note the concentration of the concentrated acid used if different from the "norm," as well as the method of preparation, i.e., v/v, wt/wt, wt/v, or v/wt. The wt. % concentrations of the common mineral acids, densities, and other information are shown in the following table:

Acid Content in Molarity

It is important to know what the concentration units of the concentrated acid being used mean. Taking 70% concentrated nitric acid as an example means that 100 grams of this acid contains 70 grams of HNO₃. The concentration is expressed at 70% wt./wt. or 70 wt. % HNO₃. Some analysts prefer to work in matrix acid concentration units of Molarity (moles/liter). To calculate the Molarity of 70 wt. % nitric acid, we calculate how many moles of HNO₃ are present in 1 liter of acid. Let's say that we tare a 1 liter volumetric flask and then dilute to the mark with 70.4 wt. % HNO₃. We would then measure the weight of the solution to be 1420 grams. Knowing that the solution is 70.4 wt % would then allow us to calculate the number of grams of HNO₃, which would be (0.704)(1420 g) = 999.7 grams HNO₃ per liter. Dividing the grams HNO₃ by the molecular weight of HNO₃ (63.01 g/mole) gives the moles HNO₃/L, or Molarity, which is 15.9 M. The above logic explains the following equation used for calculating the Molarity of acids where the concentration of the acid is given in wt %:

[(% × d) / MW] × 10 = Molarity

Where:

• % = wt. % of the acid
• d = density of acid (specific gravity can be used if density not available)
• MW = molecular weight of acid

Using the above equation to calculate the Molarity of the 70 wt % nitric acid, we have:

[(70.4 × 1.42) / 63.01] × 10 = 15.9 M

Dilutions of the concentrated acid to prepare specific volumes of specified Molarity can be made using the (mL_A)(C_A) = (mL_B)(C_B) equation.

Avoiding Precipitates

In the preparation of mixtures of the elements, it is good to avoid the formation of precipitates. It is common to form precipitates when concentrates of elements that are considered compatible (see part 1 of this series) are mixed. Many precipitates are not reversible (i.e., will not go into solution upon dilution). It is therefore best to add all of the acid and most of the water to the volumetric flask or standard solution container (dilutions to weight) before adding the individual element concentrate aliquots. Mixing after each aliquot addition is strongly advised. When diluting to volume, it is often found that the solution is above room temperature. Therefore, allow the solution to cool to room temperature and adjust to the mark with DI water. It is best to prepare the dilution the day before needed to allow for proper volume adjustment.

Storage

The following are some considerations you may want to make before the storage of chemical standard solutions:

1. Know the chemical stability of your standard. Chemical stability can be altered by changes in starting materials and preparation conditions. It is therefore advisable to perform stability studies on all standard solutions to avoid time-consuming and costly delays or mistakes, and to strictly adhere to preparation methodology, including order of addition for multi-component standard solutions.
2. Note the temperature during storage and attempt to maintain a storage temperature at or around 20 °C. Some standards are not stable for long periods at room temperature and require refrigeration or even freezing.
3. Perform the stability study in the container material selected for storage. It is not advisable to use volumetric flasks as storage containers, due to expense, contamination, and transpiration issues.
4. Determine if the standard is photosensitive, and/or store in the dark if there is a concern. This is an issue with some of the precious metals, and is a function of matrix. Photosensitivity will increase in the presence of higher energy light (sunlight as opposed to artificial light) and trace or minor amounts of organics, especially if there is an extractable proton alpha to an electron-withdrawing functional group, such as a carbonyl group. The presence of chloride may increase instability to photoreduction. A classic example is Ag⁺ in HCl solutions.
5. Store the standard in containers that will not contribute to contamination of the standard. LDPE is an excellent container for most inorganic standards.
6. Weigh the standard solution before storage, and then just before the next use. If there is measurable transpiration, the weight will decrease with time.