Modern Analytical Cheymistry - Chapter 8 pot

When the signal is the mass of a precipitate, we call the method precipitation gravimetry.. sus-8A.3 Conservation of Mass An accurate gravimetric analysis requires that the mass of analy

or a change in mass When you step on a scale after exercising you are making, in a sense, a gravimetric determination of your mass.

Measuring mass is the most fundamental of all analytical measurements, and gravimetry is unquestionably the oldest analytical technique.

8A Overview of Gravimetry

Before we look more closely at specific gravimetric methods and their applications,

let’s take a moment to develop a broad survey of gravimetry Later, as you read

through the sections of this chapter discussing different gravimetric methods, this

survey will help you focus on their similarities It is usually easier to understand a

new method of analysis when you can see its relationship to other similar methods

8A.1 Using Mass as a Signal

At the beginning of this chapter we indicated that in gravimetry we measure mass

or a change in mass This suggests that there are at least two ways to use mass as an

analytical signal We can, of course, measure an analyte’s mass directly by placing it

on a balance and recording its mass For example, suppose you are to determine the

total suspended solids in water released from a sewage-treatment facility

Sus-pended solids are just that; solid matter that has yet to settle out of its solution

ma-trix The analysis is easy You collect a sample and pass it through a preweighed

fil-ter that retains the suspended solids Affil-ter drying to remove any residual moisture,

you weigh the filter The difference between the filter’s original mass and final mass

gives the mass of suspended solids We call this a direct analysis because the analyte

itself is the object being weighed

What if the analyte is an aqueous ion, such as Pb2+? In this case we cannot

iso-late the analyte by filtration because the Pb2+is dissolved in the solution’s matrix

We can still measure the analyte’s mass, however, by chemically converting it to a

solid form If we suspend a pair of Pt electrodes in our solution and apply a

suffi-ciently positive potential between them for a long enough time, we can force the

reaction

Pb2+(aq) + 4H2O(l) tPbO2(s) + H2(g) + 2H3O+(aq)

to go to completion The Pb2+ion in solution oxidizes to PbO2and deposits on the

Pt electrode serving as the anode If we weigh the Pt anode before and after applying

the potential, the difference in the two measurements gives the mass of PbO2and,

from the reaction’s stoichiometry, the mass of Pb2+ This also is a direct analysis

be-cause the material being weighed contains the analyte

Sometimes it is easier to remove the analyte and use a change in mass as the

analytical signal Imagine how you would determine a food’s moisture content by

a direct analysis One possibility is to heat a sample of the food to a temperature

at which the water in the sample vaporizes If we capture the vapor in a

preweighed absorbent trap, then the change in the absorbent’s mass provides a

di-rect determination of the amount of water in the sample An easier approach,

however, is to weigh the sample of food before and after heating, using the change

in its mass as an indication of the amount of water originally present We call this

an indirect analysis since we determine the analyte by a signal representing its

disappearance

The indirect determination of moisture content in foods is done by difference

The sample’s initial mass includes the water, whereas the final mass is measured

after removing the water We can also determine an analyte indirectly without its

ever being weighed Again, as with the determination of Pb2+as PbO2(s), we take

advantage of the analyte’s chemistry For example, phosphite, PO33–, reduces Hg2+

to Hg22+ In the presence of Cl–a solid precipitate of Hg2Cl2forms

2HgCl2(aq) + PO33–(aq) + 3H2O(l) tHg2Cl2(s) + 2H3O+(aq) + 2Cl–(aq)+ PO43–(aq)

gravimetry

Any method in which the signal is a mass

or change in mass.

If HgCl2is added in excess, each mole of PO33–produces one mole of Hg2Cl2 Theprecipitate’s mass, therefore, provides an indirect measurement of the mass of

PO33–present in the original sample

Summarizing, we can determine an analyte gravimetrically by directly mining its mass, or the mass of a compound containing the analyte Alternatively,

deter-we can determine an analyte indirectly by measuring a change in mass due to itsloss, or the mass of a compound formed as the result of a reaction involving the analyte

8A.2 Types of Gravimetric Methods

In the previous section we used four examples to illustrate the different ways thatmass can serve as an analytical signal These examples also illustrate the four gravi-metric methods of analysis When the signal is the mass of a precipitate, we call the

method precipitation gravimetry The indirect determination of PO33–by tating Hg2Cl2 is a representative example, as is the direct determination of Cl–byprecipitating AgCl

precipi-In electrogravimetry the analyte is deposited as a solid film on one electrode in

an electrochemical cell The oxidation of Pb2+, and its deposition as PbO2on a Ptanode is one example of electrogravimetry Reduction also may be used in electro-gravimetry The electrodeposition of Cu on a Pt cathode, for example, provides adirect analysis for Cu2+

When thermal or chemical energy is used to remove a volatile species, we call

the method volatilization gravimetry In determining the moisture content of

food, thermal energy vaporizes the H2O The amount of carbon in an organic pound may be determined by using the chemical energy of combustion to convert

com-C to com-CO2

Finally, in particulate gravimetry the analyte is determined following its

re-moval from the sample matrix by filtration or extraction The determination of pended solids is one example of particulate gravimetry

sus-8A.3 Conservation of Mass

An accurate gravimetric analysis requires that the mass of analyte present in a ple be proportional to the mass or change in mass serving as the analytical signal.For all gravimetric methods this proportionality involves a conservation of mass.For gravimetric methods involving a chemical reaction, the analyte should partici-pate in only one set of reactions, the stoichiometry of which indicates how the pre-cipitate’s mass relates to the analyte’s mass Thus, for the analysis of Pb2+and PO33–described earlier, we can write the following conservation equations

sam-Moles Pb2+= moles PbO2Moles PO33–= moles Hg2Cl2Removing the analyte from its matrix by filtration or extraction must be complete.When true, the analyte’s mass can always be found from the analytical signal; thus,for the determination of suspended solids we know that

Filter’s final mass – filter’s initial mass = g suspended solidwhereas for the determination of the moisture content we have

Sample’s initial mass – sample’s final mass = g HO

precipitation gravimetry

A gravimetric method in which the

signal is the mass of a precipitate.

electrogravimetry

A gravimetric method in which the

signal is the mass of an electrodeposit on

the cathode or anode in an

electrochemical cell.

volatilization gravimetry

A gravimetric method in which the loss

of a volatile species gives rise to the

signal.

particulate gravimetry

A gravimetric method in which the mass

of a particulate analyte is determined

following its separation from its matrix.

A reagent that causes the precipitation of

a soluble species.

Specific details, including worked examples, are found in the sections of this chapter

covering individual gravimetric methods

8A.4 Why Gravimetry Is Important

Except for particulate gravimetry, which is the most trivial form of gravimetry, it is

entirely possible that you will never use gravimetry after you are finished with this

course Why, then, is familiarity with gravimetry still important? The answer is that

gravimetry is one of only a small number of techniques whose measurements

re-quire only base SI units, such as mass and moles, and defined constants, such as

Avogadro’s number and the mass of 12C.* The result of an analysis must ultimately

be traceable to methods, such as gravimetry, that can be related to fundamental

physical properties.1Most analysts never use gravimetry to validate their methods

Verifying a method by analyzing a standard reference material, however, is

com-mon Estimating the composition of these materials often involves a gravimetric

analysis.2

Precipitation gravimetry is based on the formation of an insoluble compound

fol-lowing the addition of a precipitating reagent, or precipitant, to a solution of the

analyte In most methods the precipitate is the product of a simple metathesis

reac-tion between the analyte and precipitant; however, any reacreac-tion generating a

pre-cipitate can potentially serve as a gravimetric method Most precipitation

gravimet-ric methods were developed in the nineteenth century as a means for analyzing ores

Many of these methods continue to serve as standard methods of analysis

8B.1 Theory and Practice

A precipitation gravimetric analysis must have several important attributes First,

the precipitate must be of low solubility, high purity, and of known composition if

its mass is to accurately reflect the analyte’s mass Second, the precipitate must be in

a form that is easy to separate from the reaction mixture The theoretical and

exper-imental details of precipitation gravimetry are reviewed in this section

Solubility Considerations An accurate precipitation gravimetric method requires

that the precipitate’s solubility be minimal Many total analysis techniques can

rou-tinely be performed with an accuracy of better than ±0.1% To obtain this level of

accuracy, the isolated precipitate must account for at least 99.9% of the analyte By

extending this requirement to 99.99% we ensure that accuracy is not limited by the

precipitate’s solubility

Solubility losses are minimized by carefully controlling the composition of the

solution in which the precipitate forms This, in turn, requires an understanding of

the relevant equilibrium reactions affecting the precipitate’s solubility For example,

Ag+ can be determined gravimetrically by adding Cl–as a precipitant, forming a

precipitate of AgCl

Ag+(aq) + Cl–(aq)tAgCl(s) 8.1

*Two other techniques that depend only on base SI units are coulometry and isotope-dilution mass spectrometry.

Coulometry is discussed in Chapter 11 Isotope-dilution mass spectroscopy is beyond the scope of an introductory text,

Figure 8.1

Solubility of AgCl as a function of pCl The

dashed line shows the predicted SAgCl,

assuming that only reaction 8.1 and

equation 8.2 affect the solubility of AgCl.

The solid line is calculated using equation

8.7, and includes the effect of reactions

8.3–8.5 A ladder diagram for the AgCl

complexation equilibria is superimposed on

the pCl axis.

If this is the only reaction considered, we would falsely conclude that the

precipi-tate’s solubility, SAgCl, is given by

8.2

and that solubility losses may be minimized by adding a large excess of Cl– In fact,

as shown in Figure 8.1, adding a large excess of Cl–eventually increases the tate’s solubility

precipi-To understand why AgCl shows a more complex solubility relationship thanthat suggested by equation 8.2, we must recognize that Ag+ also forms a series ofsoluble chloro-complexes

SAgCl= [Ag+] + [AgCl(aq)] + [AgCl2 ] + [AgCl32–] 8.6

Substituting the equilibrium constant expressions for reactions 8.3–8.5 into equation8.6 defines the solubility of AgCl in terms of the equilibrium concentration of Cl–

8.7

Equation 8.7 explains the solubility curve for AgCl shown in Figure 8.1 As Cl–

is added to a solution of Ag+, the solubility of AgCl initially decreases because of action 8.1 Note that under these conditions, the final three terms in equation 8.7are small, and that equation 8.1 is sufficient to describe the solubility of AgCl In-creasing the concentration of chloride, however, leads to an increase in the solubil-ity of AgCl due to the soluble chloro-complexes formed in reactions 8.3–8.5.*

Clearly the equilibrium concentration of chloride is an important parameter if the

concentration of silver is to be determined gravimetrically by precipitating AgCl In

particular, a large excess of chloride must be avoided

Another important parameter that may affect a precipitate’s solubility is the pH

of the solution in which the precipitate forms For example, hydroxide precipitates,

such as Fe(OH)3, are more soluble at lower pH levels at which the concentration of

OH–is small The effect of pH on solubility is not limited to hydroxide precipitates,

but also affects precipitates containing basic or acidic ions The solubility of

Ca3(PO4)2is pH-dependent because phosphate is a weak base The following four

reactions, therefore, govern the solubility of Ca3(PO4)2

Depending on the solution’s pH, the predominate phosphate species is either PO43–,

HPO42–, H2PO4, or H3PO4 The ladder diagram for phosphate, shown in Figure

8.2a, provides a convenient way to evaluate the pH-dependent solubility of

phos-phate precipitates When the pH is greater than 12.4, the predominate phosphos-phate

species is PO43–, and the solubility of Ca3(PO4)2will be at its minimum because

only reaction 8.8 occurs to an appreciable extent (see Figure 8.2b) As the solution

becomes more acidic, the solubility of Ca3(PO4)2increases due to the contributions

of reactions 8.9–8.11

Solubility can often be decreased by using a nonaqueous solvent A

precip-itate’s solubility is generally greater in aqueous solutions because of the ability

of water molecules to stabilize ions through solvation The poorer solvating

ability of nonaqueous solvents, even those that are polar, leads to a smaller

solu-bility product For example, PbSO4has a Ksp of 1.6×10–8in H2O, whereas in a

50:50 mixture of H2O/ethanol the Ksp at 2.6×10–12is four orders of magnitude

Figure 8.3

Schematic model of AgCl showing

difference between bulk and surface atoms

of silver Silver and chloride ions are not

shown to scale.

Avoiding Impurities Precipitation gravimetry is based on a known stoichiometrybetween the analyte’s mass and the mass of a precipitate It follows, therefore, thatthe precipitate must be free from impurities Since precipitation typically occurs in

a solution rich in dissolved solids, the initial precipitate is often impure Any rities present in the precipitate’s matrix must be removed before obtaining itsweight

impu-The greatest source of impurities results from chemical and physical tions occurring at the precipitate’s surface A precipitate is generally crystalline,even if only on a microscopic scale, with a well-defined lattice structure of cationsand anions Those cations and anions at the surface of the precipitate carry, respec-tively, a positive or a negative charge as a result of their incomplete coordinationspheres In a precipitate of AgCl, for example, each Ag+ion in the bulk of the pre-cipitate is bound to six Cl–ions Silver ions at the surface, however, are bound to nomore than five Cl–ions, and carry a partial positive charge (Figure 8.3)

interac-Precipitate particles grow in size because of the electrostatic attraction betweencharged ions on the surface of the precipitate and oppositely charged ions in solu-tion Ions common to the precipitate are chemically adsorbed, extending the crystallattice Other ions may be physically adsorbed and, unless displaced, are incorpo-rated into the crystal lattice as a coprecipitated impurity Physically adsorbed ionsare less strongly attracted to the surface and can be displaced by chemically ad-sorbed ions

One common type of impurity is an inclusion Potential interfering ions whose

size and charge are similar to a lattice ion may substitute into the lattice structure bychemical adsorption, provided that the interferent precipitates with the same crystalstructure (Figure 8.4a) The probability of forming an inclusion is greatest when theinterfering ion is present at substantially higher concentrations than the dissolvedlattice ion The presence of an inclusion does not decrease the amount of analytethat precipitates, provided that the precipitant is added in sufficient excess Thus,the precipitate’s mass is always larger than expected

Inclusions are difficult to remove since the included material is chemically part

of the crystal lattice The only way to remove included material is through itation After isolating the precipitate from the supernatant solution, it is dissolved

Bulk silver ion surrounded by 6 chloride ions Surface silver ion

surrounded by 4 chloride ions

inclusion

A coprecipitated impurity in which the

interfering ion occupies a lattice site in

the precipitate.

in a small portion of a suitable solvent at an elevated temperature The solution is

then cooled to re-form the precipitate Since the concentration ratio of interferent

to analyte is lower in the new solution than in the original supernatant solution, the

mass percent of included material in the precipitate decreases This process of

re-precipitation is repeated as needed to completely remove the inclusion Potential

solubility losses of the analyte, however, cannot be ignored Thus, reprecipitation

requires a precipitate of low solubility, and a solvent for which there is a significant

difference in the precipitate’s solubility as a function of temperature

Occlusions, which are a second type of coprecipitated impurity, occur when

physically adsorbed interfering ions become trapped within the growing precipitate

Occlusions form in two ways The most common mechanism occurs when physically

adsorbed ions are surrounded by additional precipitate before they can be desorbed

or displaced (see Figure 8.4a) In this case the precipitate’s mass is always greater than

expected Occlusions also form when rapid precipitation traps a pocket of solution

within the growing precipitate (Figure 8.4b) Since the trapped solution contains

dis-solved solids, the precipitate’s mass normally increases The mass of the precipitate

may be less than expected, however, if the occluded material consists primarily of the

analyte in a lower-molecular-weight form from that of the precipitate

Occlusions are minimized by maintaining the precipitate in equilibrium with

its supernatant solution for an extended time This process is called digestion and

may be carried out at room temperature or at an elevated temperature During

di-gestion, the dynamic nature of the solubility–precipitation equilibrium, in which

the precipitate dissolves and re-forms, ensures that occluded material is eventually

exposed to the supernatant solution Since the rate of dissolution and

reprecipita-tion are slow, the chance of forming new occlusions is minimal

After precipitation is complete the surface continues to attract ions from

solu-tion (Figure 8.4c) These surface adsorbates, which may be chemically or physically

adsorbed, constitute a third type of coprecipitated impurity Surface adsorption is

minimized by decreasing the precipitate’s available surface area One benefit of

di-gestion is that it also increases the average size of precipitate particles This is not

surprising since the probability that a particle will dissolve is inversely proportional

to its size During digestion larger particles of precipitate increase in size at the

ex-pense of smaller particles One consequence of forming fewer particles of larger size

is an overall decrease in the precipitate’s surface area Surface adsorbates also may

be removed by washing the precipitate Potential solubility losses, however, cannot

be ignored

Inclusions, occlusions, and surface adsorbates are called coprecipitates because

they represent soluble species that are brought into solid form along with the

de-sired precipitate Another source of impurities occurs when other species in

solu-tion precipitate under the condisolu-tions of the analysis Solusolu-tion condisolu-tions necessary

to minimize the solubility of a desired precipitate may lead to the formation of an

additional precipitate that interferes in the analysis For example, the precipitation

of nickel dimethylgloxime requires a pH that is slightly basic Under these

condi-tions, however, any Fe3+that might be present precipitates as Fe(OH)3 Finally,

since most precipitants are not selective toward a single analyte, there is always a

risk that the precipitant will react, sequentially, with more than one species

The formation of these additional precipitates can usually be minimized by

carefully controlling solution conditions Interferents forming precipitates that are

less soluble than the analyte may be precipitated and removed by filtration, leaving

the analyte behind in solution Alternatively, either the analyte or the interferent

can be masked using a suitable complexing agent, preventing its precipitation

Figure 8.4

Example of coprecipitation: (a) schematic of

a chemically adsorbed inclusion or a physically adsorbed occlusion in a crystal lattice, where C and A represent the cation–anion pair comprising the analyte and the precipitant, and is the impurity; (b) schematic of an occlusion by entrapment

of supernatant solution; (c) surface adsorption of excess C.

C AC AC AC AC AC AC AC AC AC A

AC AC AC AC AC AC AC AC AC AC CACACACAMACACACACACA ACACACACACACAMACACAC

C AC AC AC AC AC AC AC AC AC A AMACACACACACACACACAC

C AC AC AC AC AC AC AC AC AC A (a)

adsorbate

A coprecipitated impurity that adsorbs

to the surface of a precipitate.

M

Both of the above-mentioned approaches are illustrated in Fresenius’s cal method for determining Ni and Co in ores containing Pb2+, Cu2+, and Fe3+aspotential interfering ions (see Figure 1.1 in Chapter 1) The ore is dissolved in a so-lution containing H2SO4, selectively precipitating Pb2+as PbSO4 After filtering, thesupernatant solution is treated with H2S Because the solution is strongly acidic,however, only CuS precipitates After removing the CuS by filtration, the solution ismade basic with ammonia until Fe(OH)3precipitates Cobalt and nickel, whichform soluble amine complexes, remain in solution.

analyti-In some situations the rate at which a precipitate forms can be used to separate

an analyte from a potential interferent For example, due to similarities in theirchemistry, a gravimetric analysis for Ca2+may be adversely affected by the presence

of Mg2+ Precipitates of Ca(OH)2, however, form more rapidly than precipitates ofMg(OH)2 If Ca(OH)2is filtered before Mg(OH)2 begins to precipitate, then aquantitative analysis for Ca2+is feasible

Finally, in some cases it is easier to isolate and weigh both the analyte and theinterferent After recording its weight, the mixed precipitate is treated to convert atleast one of the two precipitates to a new chemical form This new mixed precipitate

is also isolated and weighed For example, a mixture containing Ca2+and Mg2+can

be analyzed for both cations by first isolating a mixed precipitate of CaCO3 andMgCO3 After weighing, the mixed precipitate is heated, converting it to a mixture

of CaO and MgO Thus

Grams of mixed precipitate 1 = grams CaCO3+ grams MgCO3Grams of mixed precipitate 2 = grams CaO + grams MgOAlthough these equations contain four unknowns (grams CaCO3, grams MgCO3,grams CaO, and grams MgO), the stoichiometric relationships between CaCO3andCaO

Moles CaCO3= moles CaOand between MgCO3and MgO

Moles MgCO3= moles MgOprovide enough additional information to determine the amounts of both Ca2+and

Mg2+in the sample.*

Controlling Particle Size Following precipitation and digestion, the precipitatemust be separated from the supernatant solution and freed of any remaining impu-rities, including residual solvent These tasks are accomplished by filtering, rinsing,and drying the precipitate The size of the precipitate’s particles determines the easeand success of filtration Smaller, colloidal particles are difficult to filter becausethey may readily pass through the pores of the filtering device Large, crystallineparticles, however, are easily filtered

By carefully controlling the precipitation reaction we can significantly increase

a precipitate’s average particle size Precipitation consists of two distinct events: cleation, or the initial formation of smaller stable particles of precipitate, and thesubsequent growth of these particles Larger particles form when the rate of particlegrowth exceeds the rate of nucleation

nu-A solute’s relative supersaturation, RSS, can be expressed as

8.12

where Q is the solute’s actual concentration, S is the solute’s expected

concentra-tion at equilibrium, and Q – S is a measure of the solute’s supersaturaconcentra-tion when

precipitation begins.3 A large, positive value of RSS indicates that a solution is

highly supersaturated Such solutions are unstable and show high rates of

nucle-ation, producing a precipitate consisting of numerous small particles When

RSS is small, precipitation is more likely to occur by particle growth than by

nucleation

Examining equation 8.12 shows that we can minimize RSS by either decreasing

the solute’s concentration or increasing the precipitate’s solubility A precipitate’s

solubility usually increases at higher temperatures, and adjusting pH may affect a

precipitate’s solubility if it contains an acidic or basic anion Temperature and pH,

therefore, are useful ways to increase the value of S Conducting the precipitation in

a dilute solution of analyte, or adding the precipitant slowly and with vigorous

stir-ring are ways to decrease the value of Q.

There are, however, practical limitations to minimizing RSS Precipitates that

are extremely insoluble, such as Fe(OH)3and PbS, have such small solubilities that

a large RSS cannot be avoided Such solutes inevitably form small particles In

addi-tion, conditions that yield a small RSS may lead to a relatively stable supersaturated

solution that requires a long time to fully precipitate For example, almost a month

is required to form a visible precipitate of BaSO4under conditions in which the

ini-tial RSS is 5.4

An increase in the time required to form a visible precipitate under conditions

of low RSS is a consequence of both a slow rate of nucleation and a steady decrease

in RSS as the precipitate forms One solution to the latter problem is to chemically

generate the precipitant in solution as the product of a slow chemical reaction This

maintains the RSS at an effectively constant level The precipitate initially forms

under conditions of low RSS, leading to the nucleation of a limited number of

parti-cles As additional precipitant is created, nucleation is eventually superseded by

par-ticle growth This process is called homogeneous precipitation.5

Two general methods are used for homogeneous precipitation If the

precipi-tate’s solubility is pH-dependent, then the analyte and precipitant can be mixed

under conditions in which precipitation does not occur The pH is then raised or

lowered as needed by chemically generating OH–or H3O+ For example, the

hydrol-ysis of urea can be used as a source of OH–

CO(NH2)2(aq) + H2O(l) tCO2(g) + 2NH3(aq)

NH3(aq) + H2O(l) tNH4+(aq) + OH–(aq)

The hydrolysis of urea is strongly temperature-dependent, with the rate being

negli-gible at room temperature The rate of hydrolysis, and thus the rate of precipitate

formation, can be controlled by adjusting the solution’s temperature Precipitates of

BaCrO4, for example, have been produced in this manner

In the second method of homogeneous precipitation, the precipitant itself is

generated by a chemical reaction For example, Ba2+can be homogeneously

precipi-tated as BaSO4by hydrolyzing sulphamic acid to produce SO42–

NH2SO3H(aq) + 2H2O(l) tNH4+(aq) + H3O+(aq) + SO42–(aq)

than that expected at equilibrium (RSS).

homogeneous precipitation

A precipitation in which the precipitant

is generated in situ by a chemical reaction.

Color Plate 5 shows the difference between a precipitate formed by direct precipitation and a precipitate formed

by a homogeneous precipitation.

Figure 8.5

Schematic model of the solid–solution

interface at a particle of AgCl in a solution

containing excess AgNO3.

Homogeneous precipitation affords the dual advantages of producing largeparticles of precipitate that are relatively free from impurities These advantages,however, may be offset by increasing the time needed to produce the precipitate,and a tendency for the precipitate to deposit as a thin film on the container’s walls.The latter problem is particularly severe for hydroxide precipitates generated usingurea

An additional method for increasing particle size deserves mention When aprecipitate’s particles are electrically neutral, they tend to coagulate into larger par-ticles Surface adsorption of excess lattice ions, however, provides the precipitate’sparticles with a net positive or negative surface charge Electrostatic repulsion be-tween the particles prevents them from coagulating into larger particles

Consider, for instance, the precipitation of AgCl from a solution of AgNO3,using NaCl as a precipitant Early in the precipitation, when NaCl is the limitingreagent, excess Ag+ions chemically adsorb to the AgCl particles, forming a posi-tively charged primary adsorption layer (Figure 8.5) Anions in solution, in this case

NO3 and OH–, are attracted toward the surface, forming a negatively charged ondary adsorption layer that balances the surface’s positive charge The solution

sec-outside the secondary adsorption layer remains electrically neutral Coagulation

cannot occur if the secondary adsorption layer is too thick because the individualparticles of AgCl are unable to approach one another closely enough

Coagulation can be induced in two ways: by increasing the concentration of theions responsible for the secondary adsorption layer or by heating the solution Oneway to induce coagulation is to add an inert electrolyte, which increases the concen-tration of ions in the secondary adsorption layer With more ions available, thethickness of the secondary absorption layer decreases Particles of precipitate maynow approach one another more closely, allowing the precipitate to coagulate Theamount of electrolyte needed to cause spontaneous coagulation is called the criticalcoagulation concentration

AgCl ClAgCl AgClAgCl ClAgClAgCl AgClAgClAgCl ClAgClAgClAg AgClAgClAgCl ClAgClAgCl AgClAgClAg ClAgClAgCl AgClAgCl

Secondary adsorption layer Bulk solution

coagulation

The process of smaller particles of

precipitate clumping together to form

larger particles.

Figure 8.6

Proper procedure for filtering solids using filter paper The filter paper circle in (a) is folded in half (b), and folded in half again (c) The filter paper is parted (d), and a small corner is torn off (e) The filter paper is opened up into a cone and placed in the funnel (f) Note that the torn corner is placed

to the outside.

Heating the solution and precipitate provides a second way to induce

coagula-tion As the temperature increases, the number of ions in the primary adsorption

layer decreases, lowering the precipitate’s surface charge In addition, increasing the

particle’s kinetic energy may be sufficient to overcome the electrostatic repulsion

preventing coagulation at lower temperatures

Filtering the Precipitate After precipitation and digestion are complete, the

precip-itate is separated from solution by filtration using either filter paper or a filtering

cru-cible The most common filtering medium is cellulose-based filter paper, which is

classified according to its filtering speed, its size, and its ash content on ignition

Fil-tering speed is a function of the paper’s pore size, which determines the particle sizes

retained by the filter Filter paper is rated as fast (retains particles > 20–25 µm),

medium fast (retains particles > 16 µm), medium (retains particles > 8 µm), and

slow (retains particles > 2–3 µm) The proper choice of filtering speed is important

If the filtering speed is too fast, the precipitate may pass through the filter paper

re-sulting in a loss of precipitate On the other hand, the filter paper can become

clogged when using a filter paper that is too slow

Filter paper is hygroscopic and is not easily dried to a constant weight As a

re-sult, in a quantitative procedure the filter paper must be removed before weighing

the precipitate This is accomplished by carefully igniting the filter paper Following

ignition, a residue of noncombustible inorganic ash remains that contributes a

posi-tive determinate error to the precipitate’s final mass For quantitaposi-tive analytical

pro-cedures a low-ash filter paper must be used This grade of filter paper is pretreated

by washing with a mixture of HCl and HF to remove inorganic materials Filter

paper classed as quantitative has an ash content of less than 0.010% w/w

Qualita-tive filter paper typically has a maximum ash content of 0.06% w/w

Filtering is accomplished by folding the filter paper into a cone, which is then

placed in a long-stem funnel (Figure 8.6) A seal between the filter cone and the

Figure 8.7

Proper procedure for transferring the

supernatant to the filter paper cone.

funnel is formed by dampening the paper with water and pressing the paper to thewall of the funnel When properly prepared, the stem of the funnel will fill with thesolution being filtered, increasing the rate of filtration Filtration is accomplished bythe force of gravity

The precipitate is transferred to the filter in several steps (Figure 8.7) The first

step is to decant the majority of the supernatant through the filter paper without

transferring the precipitate This is done to prevent the filter paper from becomingclogged at the beginning of the filtration process Initial rinsing of the precipitate isdone in the beaker in which the precipitation was performed These rinsings arealso decanted through the filter paper Finally, the precipitate is transferred onto thefilter paper using a stream of rinse solution Any precipitate clinging to the walls ofthe beaker is transferred using a rubber policeman (which is simply a flexible rubberspatula attached to the end of a glass stirring rod)

An alternative method for filtering the precipitate is a filtering crucible ure 8.8) The most common is a fritted glass crucible containing a porous glassdisk filter Fritted glass crucibles are classified by their porosity: coarse (retainingparticles > 40–60 µm), medium (retaining particles > 10–15 µm), and fine (re-taining particles > 4–5.5 µm) Another type of filtering crucible is the Gooch cru-cible, a porcelain crucible with a perforated bottom A glass fiber mat is placed inthe crucible to retain the precipitate, which is transferred to the crucible in thesame manner described for filter paper The supernatant is drawn through thecrucible with the assistance of suction from a vacuum aspirator or pump

(Fig-Rinsing the Precipitate Filtering removes most of the supernatant solution ual traces of the supernatant, however, must be removed to avoid a source of deter-minate error Rinsing the precipitate to remove this residual material must be donecarefully to avoid significant losses of the precipitate Of greatest concern is the po-tential for solubility losses Usually the rinsing medium is selected to ensure thatsolubility losses are negligible In many cases this simply involves the use of coldsolvents or rinse solutions containing organic solvents such as ethanol Precipitatescontaining acidic or basic ions may experience solubility losses if the rinse solution’s

Resid-pH is not appropriately adjusted When coagulation plays an important role in

de-Vent

to Vacuum

Trap Suction

flask Rubber hose

Rubber adapter Crucible

Figure 8.8

Procedure for filtering through a filtering

crucible The trap is used to prevent water

from a water aspirator from backwashing

into the suction flask.

supernatant

The solution that remains after a

precipitate forms.

termining particle size, a volatile inert electrolyte is often added to the rinse water to

prevent the precipitate from reverting into smaller particles that may not be

re-tained by the filtering device This process of reverting to smaller particles is called

peptization The volatile electrolyte is removed when drying the precipitate.

When rinsing a precipitate there is a traoff between introducing positive

de-terminate errors due to ionic impurities from the precipitating solution and

intro-ducing negative determinate errors from solubility losses In general, solubility

losses are minimized by using several small portions of the rinse solution instead of

a single large volume Testing the used rinse solution for the presence of impurities

is another way to ensure that the precipitate is not overrinsed This can be done by

testing for the presence of a targeted solution ion and rinsing until the ion is no

longer detected in a freshly collected sample of the rinse solution For example,

when Cl–is known to be a residual impurity, its presence can be tested for by

adding a small amount of AgNO3to the collected rinse solution A white precipitate

of AgCl indicates that Cl–is present and additional rinsing is necessary Additional

rinsing is not needed, however, if adding AgNO3does not produce a precipitate

Drying the Precipitate Finally, after separating the precipitate from its

super-natant solution the precipitate is dried to remove any residual traces of rinse

solu-tion and any volatile impurities The temperature and method of drying depend on

the method of filtration, and the precipitate’s desired chemical form A temperature

of 110 °C is usually sufficient when removing water and other easily volatilized

im-purities A conventional laboratory oven is sufficient for this purpose Higher

tem-peratures require the use of a muffle furnace, or a Bunsen or Meker burner, and are

necessary when the precipitate must be thermally decomposed before weighing or

when using filter paper To ensure that drying is complete the precipitate is

repeat-edly dried and weighed until a constant weight is obtained

Filter paper’s ability to absorb moisture makes its removal necessary before

weighing the precipitate This is accomplished by folding the filter paper over the

precipitate and transferring both the filter paper and the precipitate to a porcelain

or platinum crucible Gentle heating is used to first dry and then to char the filter

paper Once the paper begins to char, the temperature is slowly increased Although

the paper will often show traces of smoke, it is not allowed to catch fire as any

pre-cipitate retained by soot particles will be lost After the paper is completely charred

the temperature is slowly raised to a higher temperature At this stage any carbon

left after charring is oxidized to CO2

Fritted glass crucibles cannot withstand high temperatures and, therefore,

should only be dried in an oven at temperatures below 200 °C The glass fiber mats

used in Gooch crucibles can be heated to a maximum temperature of

approxi-mately 500 °C

Composition of Final Precipitate The quantitative application of precipitation

gravimetry, which is based on a conservation of mass, requires that the final

precipi-tate have a well-defined composition Precipiprecipi-tates containing volatile ions or

sub-stantial amounts of hydrated water are usually dried at a temperature that is

suffi-cient to completely remove the volatile species For example, one standard

gravimetric method for the determination of magnesium involves the precipitation

of MgNH4PO4⋅6H2O Unfortunately, this precipitate is difficult to dry at lower

temperatures without losing an inconsistent amount of hydrated water and

ammo-nia Instead, the precipitate is dried at temperatures above 1000 °C, where it

decom-poses to magnesium pyrophosphate, MgP O

peptization

The reverse of coagulation in which a coagulated precipitate reverts to smaller particles.

Representative Methods

An additional problem is encountered when the isolated solid is stoichiometric For example, precipitating Mn2+as Mn(OH)2, followed by heating

non-to produce the oxide, frequently produces a solid with a snon-toichiometry of MnOx,

where x varies between 1 and 2 In this case the nonstoichiometric product results

from the formation of a mixture of several oxides that differ in the oxidation state

of manganese Other nonstoichiometric compounds form as a result of lattice fects in the crystal structure.6

de-Representative Method The best way to appreciate the importance of the cal and practical details discussed in the previous section is to carefully examine theprocedure for a typical precipitation gravimetric method Although each methodhas its own unique considerations, the determination of Mg2+in water and waste-water by precipitating MgNH4PO4⋅6H2O and isolating Mg2P2O7provides an in-structive example of a typical procedure

theoreti-Method 8.1 Determination of Mg 2+ in Water and Wastewater 7

Description of Method. Magnesium is precipitated as MgNH 4 PO 4⋅6H 2 O using (NH 4 ) 2 HPO 4 as the precipitant The precipitate’s solubility in neutral solutions (0.0065 g/100 mL in pure water at 10 °C) is relatively high, but it is much less soluble

in the presence of dilute ammonia (0.0003 g/100 mL in 0.6 M NH3) The precipitant is not very selective, so a preliminary separation of Mg 2+ from potential interferents is necessary Calcium, which is the most significant interferent, is usually removed by its prior precipitation as the oxalate The presence of excess ammonium salts from the precipitant or the addition of too much ammonia can lead to the formation of Mg(NH 4 ) 4 (PO 4 ) 2 , which is subsequently isolated as Mg(PO 3 ) 2 after drying The precipitate is isolated by filtration using a rinse solution of dilute ammonia After filtering, the precipitate is converted to Mg2P2O7 and weighed.

Procedure. Transfer a sample containing no more than 60 mg of Mg 2+ into a 600-mL beaker Add 2–3 drops of methyl red indicator, and, if necessary, adjust the volume to 150 mL Acidify the solution with 6 M HCl, and add 10 mL of 30% w/v (NH4)2HPO4 After cooling, add concentrated NH3dropwise, and while constantly stirring, until the methyl red indicator turns yellow (pH > 6.3) After stirring for

5 min, add 5 mL of concentrated NH 3 , and continue stirring for an additional 10 min Allow the resulting solution and precipitate to stand overnight Isolate the

precipitate by filtration, rinsing with 5% v/v NH3 Dissolve the precipitate in 50 mL

of 10% v/v HCl, and precipitate a second time following the same procedure After filtering, carefully remove the filter paper by charring Heat the precipitate at 500 °C until the residue is white, and then bring the precipitate to constant weight at

to filter and difficult to adequately rinse free of impurities.

—Continued

2 Why is the solution acidified with HCl before the precipitant is added?

The HCl is added to ensure that MgNH4PO4⋅6H2O does not precipitate when the precipitant is initially added Because PO43– is a weak base, the precipitate is soluble in a strongly acidic solution If the precipitant is added under neutral or

basic conditions (high RSS), the resulting precipitate will consist of smaller, less

pure particles Increasing the pH by adding base allows the precipitate of MgNH4PO4⋅6H2O to form under more favorable (low RSS) conditions.

3 Why is the acid–base indicator methyl red added to the solution?

The indicator’s color change, which occurs at a pH of approximately 6.3, indicates when sufficient NH 3 has been added to neutralize the HCl added at the beginning of the procedure The amount of NH3added is crucial to this procedure If insufficient NH3is added, the precipitate’s solubility increases, leading to a negative determinate error If too much NH3is added, the precipitate may contain traces of Mg(NH 4 ) 4 (PO 4 ) 2 , which, on ignition, forms Mg(PO3)2 This increases the mass of the ignited precipitate, giving a positive determinate error Once enough NH3has been added to neutralize the HCl, additional NH3is added to quantitatively precipitate MgNH4PO4⋅6H2O.

4 Explain why the formation of Mg(PO 3 ) 2 in place of Mg 2 P 2 O 7 increases the mass

of precipitate.

The desired final precipitate, Mg 2 P 2 O 7 , contains two moles of Mg, and the impurity, Mg(PO3)2, contains only one mole of Mg Conservation of mass, therefore, requires that two moles of Mg(PO3)2must form in place of each mole

of Mg2P2O7 One mole of Mg2P2O7weights 222.6 g Two moles of Mg(PO3)2weigh 364.5 g Any replacement of Mg 2 P 2 O 7 with Mg(PO 3 ) 2 must increase the precipitate’s mass.

5 What additional steps in the procedure, beyond those discussed in questions 2

and 3, are taken to improve the precipitate’s purity?

Two additional steps in the procedure help form a precipitate that is free of impurities: digestion and reprecipitation.

6 Why is the precipitate rinsed with a solution of 5% v/v NH 3 ?

This is done for the same reason that precipitation is carried out in an ammonical solution; using dilute ammonia minimizes solubility losses when rinsing the precipitate.

8B.2 Quantitative Applications

Although not in common use, precipitation gravimetry still provides a reliable

means for assessing the accuracy of other methods of analysis or for verifying the

composition of standard reference materials In this section we review the general

application of precipitation gravimetry to the analysis of inorganic and organic

compounds

Inorganic Analysis The most important precipitants for inorganic cations are

chromate, the halides, hydroxide, oxalate, sulfate, sulfide, and phosphate A

sum-mary of selected methods, grouped by precipitant, is shown in Table 8.1 Many

in-organic anions can be determined using the same reactions by reversing the analyte

and precipitant For example, chromate can be determined by adding BaCl2 andprecipitating BaCrO4 Methods for other selected inorganic anions are summarized

in Table 8.2 Methods for the homogeneous generation of precipitants are shown inTable 8.3

The majority of inorganic precipitants show poor selectivity Most organic cipitants, however, are selective for one or two inorganic ions Several common or-ganic precipitants are listed in Table 8.4

pre-Precipitation gravimetry continues to be listed as a standard method for theanalysis of Mg2+and SO42–in water and wastewater analysis A description of theprocedure for Mg2+was discussed earlier in Method 8.1 Sulfate is analyzed by pre-cipitating BaSO, using BaCl as the precipitant Precipitation is carried out in an

Based on Precipitation

Table 8.3 Reactions for the Homogeneous

Preparation of Selected Inorganic Precipitants

Ni 2+ dimethylgloxime Ni(C 4 H 7 O 2 N 2 ) 2 Ni(C 4 H 7 O 2 N 2 ) 2

K + sodium tetraphenylborate K[B(C 6 H 5 ) 4 ] K[B(C 6 H 5 ) 4 ]

N N

NO N

N OH

OH

O H NO

acidic solution (acidified to pH 4.5–5.0 with HCl) to prevent the possible tion of BaCO3or Ba3(PO4)2and performed near the solution’s boiling point Theprecipitate is digested at 80–90 °C for at least 2 h Ashless filter paper pulp is added

precipita-to the precipitate precipita-to aid in filtration After filtering, the precipitate is ignited precipita-to stant weight at 800 °C Alternatively, the precipitate can be filtered through a fine-porosity fritted glass crucible (without adding filter paper pulp) and dried to con-stant weight at 105 °C This procedure is subject to a variety of errors, includingocclusions of Ba(NO3)2, BaCl2, and alkali sulfates

con-Organic Analysis Several organic functional groups or heteroatoms can be mined using gravimetric precipitation methods; examples are outlined in Table 8.5.Note that the procedures for the alkoxy and alkimide functional groups are exam-ples of indirect analyses

deter-Quantitative Calculations In precipitation gravimetry the relationship betweenthe analyte and the precipitate is determined by the stoichiometry of the relevantreactions As discussed in Section 2C, gravimetric calculations can be simplified

by applying the principle of conservation of mass The following exampledemonstrates the application of this approach to the direct analysis of a singleanalyte

and Heteroatoms Based on Precipitation

Organic halides Oxidation with HNO 3 in presence of Ag + AgNO 3 AgX R-X

X = Cl, Br, I

Organic Halides Combustion in O2(with Pt catalyst) in presence of Ag + AgNO3 AgX R-X

X = Cl, Br, I

Organic sulfur Oxidation with HNO 3 in presence of Ba 2+ BaCl 2 BaSO 4

Organic sulfur Combustion in O 2 (with Pt catalyst) to produce SO 2 and SO 3 ,

which are collected in dilute H 2 O 2 BaCl 2 BaSO 4

EXAMPLE 8.1

An ore containing magnetite, Fe3O4, was analyzed by dissolving a 1.5419-g

sample in concentrated HCl, giving a mixture of Fe2+and Fe3+ After adding

HNO3to oxidize any Fe2+to Fe3+, the resulting solution was diluted with water

and the Fe3+ precipitated as Fe(OH)3by adding NH3 After filtering and

rinsing, the residue was ignited, giving 0.8525 g of pure Fe2O3 Calculate the

%w/w Fe3O4in the sample

SOLUTION

This is an example of a direct analysis since the iron in the analyte, Fe3O4, is

part of the isolated precipitate, Fe2O3 Applying a conservation of mass to Fe,

we write

3×moles Fe3O4= 2×moles Fe2O3Using formula weights, FW, to convert from moles to grams in the preceding

equation leaves us with

which can be solved for grams of Fe3O4and %w/w Fe3O4in the sample

As discussed earlier, the simultaneous analysis of samples containing two analytes

requires the isolation of two precipitates As shown in Example 8.2, conservation of

mass can be used to write separate stoichiometric equations for each precipitate

These equations can then be solved simultaneously for both analytes

EXAMPLE 8.2

A 0.6113-g sample of Dow metal, containing aluminum, magnesium, and other

metals, was dissolved and treated to prevent interferences by the other metals The

aluminum and magnesium were precipitated with 8-hydroxyquinoline After

filtering and drying, the mixture of Al(C9H6NO)3and Mg(C9H6NO)2was found

to weigh 7.8154 g The mixture of dried precipitates was then ignited, converting

the precipitate to a mixture of Al2O3and MgO The weight of this mixed solid was

found to be 1.0022 g Calculate the %w/w Al and %w/w Mg in the alloy

SOLUTION

This is an example of a direct analysis in which the two analytes are determined

without a prior separation The weight of the original precipitate and the

ignited precipitate are given by the following two equations

g Al(C9H6NO)3+ g Mg(C9H6NO)2= 7.8154