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Gravimetric Methods of Analysis

Obj. 3.0, 3.1 & 3.4 Keane Campbell

MSc; BSc; ASc

January 7, 2014

Gravimetric Analysis Gravimetric analysis describes a set of methods in analytical

chemistry for the quantitative determination of an analyte

based on the mass of a solid.

Gravimetric analysis is a technique through which the amount

of an analyte (the ion being analyzed) can be determined

through the measurement of mass via two types;

1. precipitation and

2. volatilization.

Gravimetric analyses depend on comparing the masses of two

compounds containing the analyte.

The principle behind gravimetric analysis is that the mass of an

ion in a pure compound can be determined and then used to

find the mass percent of the same ion in a known quantity of an

impure compound.

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Precipitation Gravimetry

In precipitation gravimetry, the analyte is separated from a

solution of the sample as a precipitate and is converted to a

compound of known composition that can be weighed.

The analyte is converted to a sparingly soluble precipitate and

quantitatively determined by implementing the following the me

method of gravimetric analysis that involves

◦ isolation of an ion in solution by a precipitation reaction,

◦ filtering,

◦ washing the precipitate free of contaminants,

◦ conversion of the precipitate to a product of known composition,

◦ and finally weighing the precipitate and determining its mass by difference.

From the mass and known composition of the precipitate, the

amount of the original ion can be determined.

Conditions for Gravimetric Analysis In order for the analysis to be accurate, certain conditions must be met:

1) The ion being analyzed must be completely precipitated.

2) The precipitate must be a pure, easily filtered and wash free of contaminants.

3) Unreactive with constituents of the atmosphere.

4) Must be of low solubility that no significant loss of the analyte occurs during filtration and washing.

5) The weighed form of the product should be of known chemical composition, after it is dried or, if necessary, ignited.

6) By decreasing the temperature of the solution, by using an ice bath, will also decrease the solubility.

Another factor is the "common ion" effect; this further reduces the solubility of the precipitate.

When Cl- is precipitated out by addition of Ag+

Ag+ (aq) + Cl- (aq) AgCl(s)

The low solubility of AgCl is reduced still further by the excess of Ag+ which is added, pushing the equilibrium to the right.

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In most determinations the precipitate is of

such low solubility that losses from dissolution

are negligible.

It is usually difficult to obtain a product which

is "pure", i.e. one which is free from impurities

but careful precipitation and sufficient washing

helps reduce the level of impurity.

Apparatus used in Gravimetric Analysis

Desiccator

Improperly sealed desiccator.

Note the cloudiness along the

rim.

Properly sealed desiccator.

Note that the rim is clear.

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Desiccator Purpose: The desiccator is used to store dried samples in a dry atmosphere. It should not

be used to dry an object, but to maintain an already dried object indefinitely in a dry

condition.

Usage: To open - slide lid horizontally across the top to one side until it comes off. Use

one hand to hold the bottom of the desiccator while using the other hand to grasp the

knob.

To close - place lid partly on the top and slide across until desiccator is completely

closed and then rotate lid gently in both directions.

Do not attempt to lift lid off vertically.

Make sure the lid has enough grease around the ground glass rim - if necessary, spread

Vaseline uniformly on the rim. When the lid is properly seated, the greased rim will

appear.

If the desiccant appears wet or clumpy, it probably needs to be replaced with new

dessicant. It is helpful to have a small amount of indicating desiccant present. When the

color changes to pink, the desiccant should be replaced.

Desiccants should be handled in the hood and added carefully, wearing goggles and lab

coat. Desiccant should not coat the sides or plate of the desiccator.

Balances for Gravimetric

A top-loading balance

An analytical balance

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Weighing Techniques Top-loader

Direct Weighing

◦ With nothing on the pan, set to zero by pressing the "on" button.

◦ Place weighing bottle, beaker, or vial on balance and set to zero again.

◦ Use a clean scapula to transfer sample into container slowly, until you reach the

desired mass.

Indirect weighing (Weighing by difference)

◦ Place enough of the sample in a weighing bottle, put the lid on, and place on the scale.

Record the mass.

◦ Take some out and place it in a different container (whatever you will be using for the

experiment). Record the new mass. The difference in mass is the mass of the sample

transferred.

◦ Continue this procedure until you have as much sample as you need.

◦ It is best to transfer small amounts at a time, so you do not take more than you need.

You should not put excess sample back into the weighing bottle.

Weighing Techniques

Analytical Balance

Use the same procedure as with a top-loader, remembering these additional points:

Close all the doors before taking measurements.

Remember the number of significant figures. It is higher than on a regular top-loader.

Make sure the sample is completely cooled when weighing. If a sample is still warm, it will weigh less because of buoyancy due to upward circulation of hot air.

For example, a 50 mL beaker 3 minutes after removal from a 110 degree oven weighs 27.0271 g. At room temperature, it weighs 27.0410 g.

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Gravimetric Experiment

Put enough unknown into a weighing bottle with the lid on sideways and dry in the oven. Cool in a desiccator.

Sample in a weighing bottle.

Note the position of the lid for heating.

Indirectly weigh some mass, determined to 0.1 mg, of unknown into beaker.

Dissolve the unknown.

Add a precipitating agent to the solution

Optional - heat the solution on a hot plate to increase the particle size for

easier filtering (see Figure 3 above). This is usually referred to as digestion.

Test for complete precipitation by adding a drop of the precipitating agent

and looking for any sign of precipitate.

Filter the solution using vacuum filtration. Use a rubber policeman to make

sure all the precipitate has been transferred from the beaker to the filter. It

is important that the precipitate is quantitatively transferred to the filter. If

any remains in the beaker, the mass obtained will be inaccurate.

Dry and weigh the precipitate.

Use stoichiometry to determine the mass of the ion being analyzed.

Find percent by mass of analyte by dividing the mass of the analyte by the

mass of the unknown.

Gravimetric Experiment

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Examples of Gravimetric Analysis

An example of a gravimetric analysis is the determination of

chloride in a compound.

In order to do a gravimetric analysis, a cation must be found

that forms an insoluble compound with chloride.

This compound must also be pure and easily filtered.

The solubility rules indicate that Ag+, Pb2+, and Hg22+ form

insoluble chlorides.

Therefore silver chloride could be used to determine %Cl-,

because it is insoluble (that is, about 99.9% of the silver is

converted to AgCl) and it can be formed pure and is easily

filtered.

Another example of precipitation method for determining calcium in natural waters is recommended by the Association of Official Analytical Chemists.

Here, an excess of oxalic acid, H2C2O4, is added to an aqueous solution of the sample. Ammonia is then added, which neutralizes the acid and causes essentially all of the calcium oxalate. The reactions are

2NH3 + H2C2O4 2NH4+ + C2O4

2-(aq)

Ca2+ (aq) + C2O42- (aq) CaC2O4 (s)

The precipitate is filtered using a weighed filtering crucible, then dried and ignited.

The process converts the precipitate entirely as calcium oxide. The reaction is

CaC2O4 (s) CaO (s) + CO2 (g) + CO (g)

After cooling the crucible and the precipitate are weighed, and the mass of calcium oxide is determined by subtracting the known mass of the crucible.

Now students in order to fully appreciate the gravimetric method we will proceed unto objective 3.4 of your syllabus to calculate the calcium content of the sample – natural waters.

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The calcium in a 200 cm3 sample of a natural water was determined by precipitating

the cation as CaC2O4. The precipitate was filtered, washed, and ignited in a crucible

with an empty mass of 26.6002 g. The mass of the crucible plus CaO (56.077 g mol-1)

was 26.7134 g. Calculate the concentration of Ca (40.078 g mol-1) in water in units

of gram per 100 cm3 of the water.

Answer:

The mass of CaO is

26.7314 𝑔 − 26.600 𝑔 = 0.1132 𝑔

The number of moles Ca in the sample is equal to the number of moles CaO or

Amount of Ca

0.1132𝑔 𝐶𝑎𝑂 × 1 𝑚𝑜𝑙 𝐶𝑎𝑂

56.077𝑔×

1 𝑚𝑜𝑙 𝐶𝑎

𝑚𝑜𝑙 𝐶𝑎𝑂

= 2.0186 × 10−3𝑚𝑜𝑙 𝐶𝑎

Conc. Ca

100 𝑚𝑙 ×2.0186 × 10−3 𝑚𝑜𝑙 𝐶𝑎 × 40.078

𝑔𝐶𝑎𝑚𝑜𝑙𝐶𝑎

200 𝑚𝑙

= 0.0566 𝑔 𝑝𝑒𝑟 100 𝑚𝑙

Gravimetric Calculations The following calculations would be done for the gravimetric determination of chloride:

Mass of sample of unknown chloride after drying: 0.0984 g

Mass of AgCl precipitate: 0.2290 g

One mole of AgCl contains one mole of Cl-.

∴0.2290 𝑔 𝐴𝑔𝐶𝑙

143.323𝑔

𝑚𝑜𝑙

= 1.598 × 10−3 𝑚𝑜𝑙 𝐴𝑔𝐶𝑙

1.598 × 10−3 𝑚𝑜𝑙 𝐴𝑔𝐶𝑙

35.453𝑔

𝑚𝑜𝑙𝐶𝑙

= 0.0566 𝑔 𝐶𝑙

0.0566 𝑔 𝐶𝑙

0.0984 𝑔 𝑠𝑎𝑚𝑝𝑙𝑒× 100% = 57.57% 𝐶𝑙 in unknown chloride sample.

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Gravimetric Calculations Notice that even though the mass of sample (0.0984) only

contains three significant figures, the number is known to one

part in a thousand (0.0001/0.0984 = 1/1000).

The number 0.0984 therefore actually is "good" to four

significant figures and the answer can be expressed to four

significant figures.

If Pb2+ had been used to precipitate the chloride, the

calculation would need to be modified to account for the fact

that each mole of PbCl2 contains two moles of chloride.

Lead would not be a good precipitating reagent, however,

because PbCl2 is moderately soluble and therefore a small

amount of chloride would remain in solution, rather than in

the precipitate.

Gravimetric Calculations General calculation of the percent by mass of analyte in a sample:

Write the balanced chemical equation for the precipitation reaction.

Calculate the moles of precipitate:

𝑚𝑜𝑙𝑒𝑠 = 𝑚𝑎𝑠𝑠 ÷ 𝑚𝑜𝑙𝑒𝑐𝑢𝑙𝑎𝑟 𝑚𝑎𝑠𝑠

Calculate moles of analyte from the balanced chemical equation using

the mole ratio of analyte : precipitate

Calculate mass of analyte:

𝑚𝑎𝑠𝑠 = 𝑚𝑜𝑙𝑒𝑠 × 𝑚𝑜𝑙𝑒𝑐𝑢𝑙𝑎𝑟 𝑚𝑎𝑠𝑠

Calculate percent by mass of analyte in sample: 𝑀𝑎𝑠𝑠 𝑜𝑓 𝑎𝑛𝑎𝑙𝑦𝑡𝑒

𝑀𝑎𝑠𝑠 𝑜𝑓 𝑠𝑎𝑚𝑝𝑙𝑒 × 100

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Gravimetric Calculations

Example:

A 2.00g sample of limestone was dissolved in hydrochloric

acid and all the calcium present in the sample was converted

to Ca2+(aq).

Excess ammonium oxalate solution, (NH4)2C2O4(aq), was

added to the solution to precipitate the calcium ions as

calcium oxalate, CaC2O4(s).

The precipitate was filtered, dried and weighed to a constant

mass of 2.43g.

Determine the percentage by mass of calcium in the

limestone sample.

Gravimetric Calculations Write the balanced chemical equation for the precipitation reaction:

Ca2+(aq) + C2O4

2-(aq) CaC2O4(s)

Calculate the moles of calcium oxalate precipitated.

n(CaC2O4(s)) = mass ÷ MM

n(CaC2O4(s)) = 2.43 ÷ (40.08 + 2 x 12.01 + 4 x 16.00)

n(CaC2O4(s)) = 2.43 ÷ 128.10

n(CaC2O4(s)) = 0.019 mol

Find the moles of Ca2+(aq).

From the balanced chemical equation, the mole ratio of Ca2+ : CaC2O4(s) is 1 : 1

So, n(Ca2+(aq)) = n(CaC2O4(s)) = 0.019mol

Calculate the mass of calcium in grams, mass (Ca) = n x MM

mass (Ca) = 0.019 x 40.08 = 0.76g

Calculate the percentage by mass of calcium in the original sample:

%Ca = (mass Ca ÷ mass sample) x 100

%Ca = (0.76 ÷ 2.00) x 100 = 38%

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Gravimetric Calculations

(Graded Homework)

1. A 0.4960 g sample of a CaCO3 is dissolved in an acidic solution. The calcium is

precipitated as CaC2O4∙H2O and the dry precipitate is found to weigh 0.6186 g.

What is the percentage of CaO in the sample?

2. A compound of Iron (Fe) and Chlorine (Cl) is soluble in water. An excess of

Silver Nitrate (AgNO3) was added to precipitate the chloride ion as silver

chloride. If a 134.8 mg sample of the compound gave 304.8 mg of AgCl, what is

the formula of the compound?

3. When a sample of impure potassium chloride (0.4500g) was dissolved in water

and treated with an excess of silver nitrate, 0.8402 g of silver chloride was

precipitated. Calculate the percentage KCl in the original sample.

4. A certain barium halide exists as the hydrated salt BaX2.2H2O, where X is the

halogen. The barium content of the salt can be determined by gravimetric

methods. A sample of the halide (0.2650 g) was dissolved in water (200 cm3) and

excess sulfuric acid added. The mixture was then heated and held at boiling for

45 minutes. The precipitate (barium sulfate) was filtered off, washed and dried.

Mass of precipitate obtained = 0.2533 g. Determine the identity of X.

Volatilization Gravimetry In volatilization gravimetry, the analyte is separated from other constituents of a sample by conversion to a gas of known chemical composition.

The weight of this gas then serves as a measure of the analyte concentration.

The two most common gravimetric methods based on volatilization are those for determining water and carbon dioxide.

Water is quantitatively distilled from many materials by heating. In direct determination, water vapour is collected on any of the several solid desiccants, and its mass is determined from the mass gain of the desiccant.

The indirect method, in which the amount of water vapour is determined by the loss of mass of the sample during heating, is less satisfactory because it must be assumed that water is the only component volatilized.

This assumption is frequently unjustified, however because heating of many substances results in their decomposition and a consequent change in mass, irrespective of the presence of water.

Nevertheless, the indirect method has found wide use for the determination of water in items of commerce.

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Examples of Volatilization Gravimetry A semi-automated instrument for the determination of moisture in cereal

grains can be purchased.

It consists of a platform balance on which a 10-g sample is heated with an

infrared lamp. The percent moisture is read directly.

Another example of gravimetric procedure involving volatilization of CO2 is

the determination of the NaHCO3 content of antacid tablets.

Here a weighed sample of the finely ground tablets is treated with dilute H2SO4 to

convert the NaHCO3 to CO2.

NaHCO3 (aq) + H2SO4 (aq) CO2 (g) + H2O (l) + NaHSO4 (aq)

As shown below, this reaction is carried out in a flask connected to a tube

containing the absorbent on a non-fibrous silicate.

This material retains CO2 by the reaction

2NaOH (aq) + CO2 (g) Na2CO3 (aq) + H2O (l)

The absorption tube must also contain a desiccant to prevent loss of the

water produced by the reaction.

Sulphides and sulphites can also be determined by volatilization. Hydrogen

sulphide or sulphur dioxide evolved from the sample after treatment with

acid is collected in a suitable experiment.

Finally, the classical method for the determination of carbon and hydrogen in

organic compounds is a gravimetric volatilization procedure in which the

combustion products (H2O and CO2) are collected selectively on weighed

absorbents.

The increase in mass serves as an analytical parameter.

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Uses of gravimetric analysis in quality control

Measurement of the “essential” elements in plant

foods (phosphorus, for example, is converted into the

insoluble salt, magnesium ammonium phosphate).

Estimation of pollutants in the air, such as sulphur

dioxide (by conversion to insoluble barium sulphate).

Estimation of sulphur dioxide (used to prevent

microbial spoilage) in soft drinks, such as orange juice.

Estimation of chloride ions in water supplies (by

conversion to insoluble silver chloride).

Reference

Skoog, West, Holler & Crouch,

‘Fundamentals of Analytical Chemistry’ 6th Ed.

Chang