主题：ICP-AES系列学习3：Handling, Calculations, Preparation, and Storage

3. Handling, Calculations, Preparation, and Storage
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:

Never put solution transfer devices into the standard solution. This precaution avoids possible contamination from the pipette or transfer device.

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.

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 concentrations units are in wt./wt. rather than wt./volume.

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.

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.

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.

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 礸/mL (micrograms per milliliter) or ng/mL (nanograms per milliliter). For example, a 1000 礸/mL solution of Ca+2 contains 1000 micrograms of Ca+2 per each mL of solution and a 1 礸/mL solution of Ca+2 contains 1000 ng of Ca+2 per milliliter of solution. To convert between metric concentration units the following conversions apply:

Table 3.1: Mass portion of concentration unit where g = gram

Prefix  Scientific Notation  Decimal equivalents  Example Units
kilo- (k)  = 103  = 1000 g  kilogram (kg)
milli- (m)  = 10-3  = 0.001 g  milligram (mg)
micro- (?  = 10-6  = 0.000001 g  microgram (礸)
nano- (n)  = 10-9  = 0.000000001 g  nanogram (ng)
pico- (p)  = 10-12  = 0.000000000001 g  picogram (pg) 

Table 3.2: Volume portion of concentration unit where L = liter

Prefix  Scientific Notation  Decimal equivalents  Example Units
milli- (m)  = 10-3  = 0.001 L  milliliter (mL)
micro- (?  = 10-6  = 0.000001 L  microliter (礚)
nano- (n)  = 10-9  = 0.000000001 L  nanoliter (nL)
pico- (p)  = 10-12  = 0.000000000001 L  picoliter (pL) 

The difference between ppm and 礸/mL is often confused. A common mistake is to refer to the concentration units in ppm as a short cut (parts per million) when we really mean 礸/mL. One ppm is in reality equal to 1 礸/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 礸/mL or ng/mL the density of the solution must be known. The equation for conversion between wt./wt. and wt./vol. units is:

(礸/g) (density in g/mL) = 礸/mL

and/or

(ng/g) (density in g/mL) = ng/mL
Therefore, if we have a solution that is 1000 礸/mL Ca+2 and know or measure the density to be 1.033 g/mL then the ppm Ca+2 = (1000 礸/mL) / (1.033 g/mL) = 968 礸/g = 968 ppm.

When making dilutions the following equation is useful:

(mLA)(CA) = (mLB)(CB)
For example, to determine how much of a 1000 礸/mL solution of Ca+2 required to prepare 250 mL of a 0.3 礸/mL solution of Ca+2 we would use the above equation as follows:

(mLA)(1000 礸/mL) = (250 mL)(0.3 礸/mL)

(mLA) = [(250 mL)(0.3 礸/mL)]/ (1000 礸/mL)

(mLA) = 0.075 mL = 75 礚

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 greaten 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):

dS = [(100-%) + (dA)(%)] / 100
Where:

dS = density of final solution
% = The v/v % of a given aqueous acid solution
dA = density of the concentrated acid used
For example, lets 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.

DS = [(100-%) + (dA)(%)]/100 = [(100-10) + (1.42)(10)]/100 = (90 + 14.2)/100 = 1.042 g/mL