GRAVIMETRIC ANALYSISQuantitative method based upon measuring the massof a pure compound to which the analyte isof a pure compound to which the analyte ischemically related.

Two types of gravimetric methods:

1. Precipitation method- based on determination of an analyte which is precipitated y p pby a reagent.-the precipitate is a slightly soluble substance with a known composition or

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composition or- it can be converted to one of known composition.

Eg.1: determination of Cl− by the addition of Ag+ to form AgCl;Ag+ + Cl− AgCl (S)

Eg.2: precipitation of Fe (III) as Fe (III) hydroxide. The hydroxide will be converted to weighable oxide.

The precipitated form and weighed forms are not be the same compound.

Fe3+ + 3NH3 + (x+3)H2O Fe(OH)3.xH2O(s) + 3NH4+

∆,1000oC2Fe(OH)3.xH2O(s) → Fe2O3(s) + (2x + 3)H2O

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2 Volatilization methods2. Volatilization methods- for determination of volatile components (H2O,

CO2, etc.) of a sample.CO2, etc.) of a sample.- The sample is warmed or ignited and the amount of

the component is found from the loss in mass of pthe sample. Or, volatile components can be absorbed by a suitable absorbent. The amount of the components are found from the increase in the weight of the absorbent.

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STEPS IN GRAVIMETRIC ANALYSIS

Preparation of the

SampledissolvedcomponentsSample

Precipitationsample

components

Calc lationDigestion precipitating agent

Calculation

Filtration and W hi WeighingWashing

4Drying or Ignition

Preparation of the solution

Solid sample must be dissolved in a suitable solvent.Some form of preliminary separation may be necessary toeliminate interfering materials.g

Requirements for Precipitatelow solubilitylow solubility.the precipitate in a form suitable for filtration.Proper adjustment of the solution condition may also mask

i l i fpotential interferences.

Factors that Must be Considered when Preparing the SolutionV l f th l ti d i i it tiVolume of the solution during precipitation.Concentration range of the test substance.The presence and concentrations of other constituents.

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TemperaturepH

Precipitation Process

• Ideally, we’d like a precipitate that forms quickly. This implies:Large, pure crystalsLow solubility Analyte + Precipitating Agent →Low solubilityEasily filteredEasily washed

y p g gSupersaturation

→

● How does precipitation occur?1. As Ksp is exceeded, solution becomes supersaturated.

Precipitation

AgCl(s) Ag+ + Cl- Ksp = [Ag+] [Cl-] / [AgCl]

2 At some point nucleation begins. (Nucleation - a process p g ( pinvolving formation of very small aggregates of a precipitating solid.

3. At the same time, crystal growth begins. (Crystal growth –formation of larger particles of a precipitating solid by deposition of i / l l h f f i i l i

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ions/molecules on the surface of existing nuclei.

The Solubility of Precipitates

The solubility of a sparingly soluble electrolyte is characterized by its Solubility Product (Ksp). y y ( sp)

(Ksp – the equilibrium constant for a reaction in which a solid dissociates into its ions)

Example:Example:AgCl, is a slightly soluble salt. The solubility equilibrium can be represented

as;

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

(aq) Ksp = [Ag+][Cl-]

The solubility product (Ksp) of a compound is used as a measure of they p ( sp) psolubility of sparingly soluble salt.

For compounds which have similar formula, the smaller Ksp, the less soluble

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pthe compound.

Example:A Cl K 1 6 10 10AgCl Ksp = 1.6 x10-10

AgBr Ksp = 7.7 x10-13

AgI Ksp = 8.3 x 10-17sp

The solubility product constant of AgI is smaller than AgBr and AgCl. Itmeans that the solubility of the AgI is smaller than the others.

But when comparing salts of different valence type, the order may bedifferent.

Solubility product expressions for more complex formulas are given as

MgF2(s) Mg2+(aq) + 2F-(aq) Ksp = [Mg2+][F-]2

y p p p gbelow;

Ca3(PO4)2(s) 3Ca2+(aq) + 2PO43-(aq) K = [Ca2+]3[PO4

3-]2

Ag2C03(s) 2Ag+(aq) + CO32-(aq) Ksp = [Ag+]2[CO3

2-]

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Ca3(PO4)2(s) 3Ca (aq) + 2PO4 (aq) Ksp [Ca ] [PO4 ]

Example:

Calculate the solubility of AgCl in a saturated solution of AgCl ifCalculate the solubility of AgCl in a saturated solution of AgCl ifKsp of AgCl at 25 oC is 1.0 x 10-10.

When AgCl ionizes, equal amounts of Ag+ and Cl- are formed.When AgCl ionizes, equal amounts of Ag and Cl are formed.Let say, s is molar solubility of AgCl.

A Cl A + ClAgCl(s)x xAg+

(aq) + Cl-(aq)

Ksp = [Ag+][Cl-] = (x) (x)x2 = 1.0 x 10-10

x = √ 1.0 x 10-10x √ 1.0 x 10= 1.0 x 10-5 M

So the solubility of AgCl is 1 0 x 10-5 M9

So, the solubility of AgCl is 1.0 x 10-5 M

Example:Which has greater solubility in water: A IO (K 3 0 10 8) L (IO ) (K 6 5 10 12) ?AgIO3 (Ksp = 3.0 x 10-8) or La(IO3)3 (Ksp= 6.5 x 10-12) ?

AgIO3x x

Ag+ + IO3-

x xKsp = [Ag+][ IO3

-] = 3.0 x 10-8

[Ag+] = [ IO3-] = x

(x) (x) = 3.0 x 10-8(x) (x) 3.0 x 10x = √ 3.0 x 10-8 = 1.73 x 10-4 mol L-1

La(IO3)3 La3+ + 3IO3-( 3)3

x 3x

[La3+] [IO3- ]3 = 6.5 x 10-12 (x) (3x)3 = 6.5 x 10-12

27x4 = 6 5 x 10-12

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27x 6.5 x 10x = 4√6.5 x 10-12/27 = 7.0 x 10-4 mol L-1

The Ksp for AgIO3 is greater than that for La(IO3)3, but La(IO3)3 is more

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sp g 3 g ( 3)3, ( 3)3soluble.

Effect of pH on Completeness of Precipitation

pH of solution influences the degree of precipitation.

Effect of pH on Completeness of Precipitation

Example:

Mg(OH)2(s) Mg2+(aq) + 2OH-

(aq)

Adding OH- ions (increasing the pH) shifts the equilibriumAdding OH ions (increasing the pH) shifts the equilibriumfrom right to left, thereby increasing the precipitation(decreasing the solubility) of Mg(OH)2.

Adding H+ ions (decreasing the pH) shifts the equilibriumfrom left to right and the solubility of Mg(OH)2 increases.

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Example:At what pH will Al(OH)3 begin to precipitate fromp ( )3 g p p0.10 M AlCl3

The equilibrium is;

Al(OH)3 Al3+ + 3OH-

[Al3+][OH-]3 = Ksp(0.1) [OH-]3 = 2 x 10-23

[OH-] = 3√2 x 10-23/0.1[OH ] √2 x 10 /0.1= 5.848 x 10-8 M

pOH = -log [5.848 x 10-8 ] = 7.23p g [ ]pH = 14 – 7.23

= 6.77The aluminium hydroxide precipitates when the pH just

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The aluminium hydroxide precipitates when the pH just exceeds 6.77.

Effect of Temperature on Completeness of Precipitation

The solubility product depends on temperature. If the temperature alters, the solubility product of the p , y pprecipitate also changes.

For example the solubility of AgCl at 100oC is nearlyFor example, the solubility of AgCl at 100 C is nearly 25 times as high as at 10oC.

B i i lik B SO h l bili lBut some precipitate like BaSO4, the solubility onlydoubled when the temperature is raised from 10oC to100oC.

In some instances the solubility of precipitate decreaseswith rise of temperature

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with rise of temperature.

Predicting Precipitation Condition

When two solutions are mixed or when a compound is added to a solution, any one of the following condition may exist; the solution is,

· Unsaturated· Saturated

S d ( i i i b i )· Supersaturated (precipitation begins)

A slightly soluble salt, MX contains M+ and X− ions in aqueous solution.

MX M+ + X-

The ion product, Q is the product of the concentrations of the ions at anymoment in time;

Q = [M+][X-]

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The relationships between Q and Ksp are;

Q < Ksp: Unsaturated solution

Q = Ksp: Saturated solution (the system is at equilibrium)

Q > Ksp : Supersaturated solution; MX will precipitate until the

d t f th i iproduct of the ionic concentrations is equal to Ksp

Precipitation occurred when the ion product exceeds Ksp for that substance.

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Effect of Common Ion on Precipitate

The condition of precipitation can also be affected by adding more of any of the precipitating ions.

Effect of Common Ion on Precipitate

Example:If a solution is saturated with barium sulfate,If a solution is saturated with barium sulfate,

BaSO4(s) Ba2+ + SO42-, Ksp = 9.12 x10-11

If, more sulfate or Ba2+ is added to the solution, the reactionwill shift to the left ( more precipitate formed).

This effect is called the common ion effect.

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E lExample:

Calculate the molar solubility of lead iodide;(a) in water, (b) in 0.20 M NaI solution.

The K for PbI2 is 7.9 x 10-9.The Ksp for PbI2 is 7.9 x 10 .

(a) The equilibrium is;PbI Pb2+ + 2I-PbI2 Pb2+ + 2I-

s 2sfor which Ksp = [Pb2+] [ I- ]2

p(s) (2s)2 = 7.9 x 10-9

4s3 = 7.9 x 10-9

solubility = s = 1.3 x 10-3 M

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solubility s 1.3 x 10 M

(b) The same equilibrium expression holds:(b) The same equilibrium expression holds:Ksp = [Pb2+] [ I- ]2

T f i did th N I d PbI Th t fTwo sources of iodide: the NaI and PbI2. The amount of iodide coming from the PbI2 is small compared to that from the NaI. Thus,

[ I- ] + 0 2 0 2[ I- ] = x + 0.2 ≈ 0.2

Then, [Pb2+] = solubility = Ksp /[ I- ]2p

= 7.9 x 10-9/(0.20)2

= 2.0 x 10-7 M

The solubility has decreased upon addition of an excess of I-.

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Particle Size of PrecipitatespColloidal suspensions

Tiny particles, invisible to the naked eyes (10-7 – 10-4 cm diameter)No tendency to settle from the solutionBrownian motion prevents their settling out of the solution under the influence of gravityColloidal suspensions are stable because all of colloidal particles are either positively or negatively charged because anions or cations areeither positively or negatively charged because anions or cations are adsorbed on the surface of the particlesNot easily filtered

Colloidal suspensions appear clear and contain no solid but becauseColloidal suspensions appear clear and contain no solid, but becauseparticles of colloidal suspensions scatter visible light, the path of the beamcan be seen by the eye. This phenomenon is called Tyndall effect.

Crystalline suspensionsLarger particles (tenth of a mm or greater)Settle spontaneously

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p yEasily filtered

Factors affecting particle size:g pPrecipitate solubilityTemperatureReactant concentrationsRate of mixing the reactantsRelative supersaturation

Q - SRelative supersaturation (RSS) =

SVon Weimarn equation

SQ = the concentration of the soluteS = equilibrium solubility of the solute

Particle size of a precipitate varies inversely with the relative supersaturation during mixing of reactants.When (Q – S)/S is large → small precipitate size

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When (Q – S)/S is small → larger precipitate size

Growth of precipitates andGrowth of precipitates and crystals

e.g. sodium acetate crystals from a supersaturated solutionpsmall particles and difficult to filter

In gravimetric analysis we minimize the supersaturation in porder to obtain larger and easily filterable solid particles

In a supersaturated solution, nucleation goes faster than particle growth.growth.

This is bad because you wind up with many small particles in solution (a colloid) and few large ones

M t d thi t t ti l th21

Must do something to promote particle growth

Supersaturated SolutionUnstable solution that contains a higher soluteUnstable solution that contains a higher solute concentration than a saturated solutionBy time, the excess solute will precipitates out of the

d l isupersaturated solution

Mechanism of Precipitate Formationa) Nucleation

A process in which a number of atoms, ions, or molecules join together to give a stable solid, i.e. molecules in j g gsolution randomly form small aggregates of moleculesNucleation may also occur on the surface of suspended solid contaminants

b) Particle GrowthAfter nucleation, further precipitation takes place on the existing nuclei i e particle growth

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existing nuclei, i.e. particle growth.i.e. the addition of more molecules to the aggregate to form a crystal

Precipitates form by nucleation and particle growthp y p gIf nucleation predominates, a large number of very fine particles resultsIf particle growth predominates a smaller number ofIf particle growth predominates, a smaller number of larger particles is obtained

If: Rate of Nucleation > Rate of Particle GrowthIf: Rate of Nucleation > Rate of Particle Growththen precipitate containing a large number of small particles.

If: Rate of Nucleation < Rate of Particle Growththen precipitate containing a smaller number of largerparticles is produced.p p

Precipitates with large particle are more easily handled duringfiltration and washing

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filtration and washing.

How to Control Particle Size?Elevate the temperature to increase solubility (S) of the precipitates, decreasing supersaturationPrecipitate from dilute solutions (minimize Q), => volume of solution large, analyte and reagent concentrations are kept lowconcentrations are kept lowSlow addition of the precipitating agent (to minimize Q) with good stirring to avoid localized supersaturationsupersaturationPrecipitation from acidic solution. If the precipitate depends on pH, larger particle size can be obtained by controlling the pH.After the precipitate is formed, the particle size can be improved by digestion process or by precipitation

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p oved by d ges o p ocess o by p ec p ofrom homogeneous solution.

C ll id l P i iColloidal Precipitates

A colloidal precipitates is formed when a precipitate has a lowA colloidal precipitates is formed when a precipitate has a low solubility: S <<< Q, because relative supersaturation remains high throughout the precipitation process.

Example of colloidal precipitates (because of their very low solubilities): )

• hydrous oxide of iron(III), aluminum, and chromium(III)

• Sulfides of most heavy metal ions• Sulfides of most heavy metal ions

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Why colloid suspensions are stable andWhy colloid suspensions are stable and do not coagulate spontaneously?Particles of colloid are charged: either +ve or –veDue to cations or anions bound to (adsorbed on) the surface of the colloid particlesthe colloid particlese.g. Ag+ (or Cl-) on the surface of AgCl particle attracts anions (or cations) in the solution

How do you prove that the colloid The electrical double layeron the colloid particles y p

particles are charged?Observe their migration when placed in an electrical field

on the colloid particlesprevents particles fromcolliding and adhering

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Adsorption of ions onto the colloidal surface:

The kind of ions retained on the surface of a colloidal particle ,

Adsorption of ions onto the colloidal surface:

The kind of ions retained on the surface of a colloidal particle , and their numbers depend on:

Lattice ions are more strongly held than others, e.g if AgNO3 is first added to a solution of containing Cl- the colloidal particles of theadded to a solution of containing Cl , the colloidal particles of the precipitate are negatively charged due to adsorption of some excess Cl-.

The charge becomes positive when excess of Ag+ is then added.

The surface charge is minimum when the supernatant liquid contains an excess of neither ion.

When the concentration of the common ion becomes greater, the extent of adsorption (and thus the charge) of a particle increase rapidly

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rapidly

Precipitation of Silver Chloride (AgCl)

• AgCl colloid particle in a solution contains ana solution contains an excess of AgNO3.

• Primary adsorption layer: attached directly t th lid fto the solid surface (Ag+)

• Counter-ion layer: excess of –ve ions (NO3

-) surrounding the charged particles to just balance the charge on the particlethe particle.

• Electrical double layer: primary adsorbed layer and the counter-

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ion layer

Precipitation Using an Electrolyte

– Ionic compounds are usually precipitated in the presence of an electrolyte; for AgCl

The surface of the particle will have a small positive charge due to adsorption of excess silver ions.

h h d hThe ionic atmosphere surrounds the particle and has a slight net negative charge.

For particle growth, colloidal particles must collide to coalesce. But the negative atmospheres around the particles repulse each other electrostatically. Adding electrolyte decreases the volume of th i i t h ll ithe ionic atmospheres, allowing repulsion to be overcome.

In the concentrated AgNO3: the effective charge on thethe effective charge on the particles prevents them from approaching one another more closely than 2d1 – a distanceclosely than 2d1 a distance that is too great for coagulation to occur

In the diluted AgNO3: The two particles can approachThe two particles can approach within 2d2 of one another. As the AgNO3 becomes more diluted, the distance between ,particles becomes small enough for the forces of coagulation to take effect and

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coagulated precipitate to appear.

Coagulation of ColloidsgWe can coagulate or agglomerate individual particles of most colloids to give a filterable, amorphous mass that g , psettle out of solutionCoagulation can be hastened by heating, by stirring, and by adding an electrolyte to the mediumadding an electrolyte to the medium.Heating

Decreases the number of adsorbed ions, and thus the thickness of the double layerdouble layerThe particles gain kinetic energy at higher temperature to overcome the barrier to close approach posed by the double layer.

Increase the electrolyte concentrationyThe concentration of counter-ions increases in the vicinity of each particle, The volume of the solution that contains sufficient counter-ions to balance the charge of the primary adsorption layer decreases.Adding electrolyte shrinks the counter ion layer the particles can approach

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Adding electrolyte shrinks the counter-ion layer, the particles can approach one another more closely and agglomerate.

Properties of Precipitates and Precipitating Agents

G i t i i it ti t h ldGravimetric precipitating agent shouldReact specifically or selectivelyReact with the analyte to give product that is

easil filtered and ashed free of contaminantseasily filtered and washed free of contaminantsof low solubility• no significant loss of the analyte occurs during filtration and

washingwashing• unreactive with the constituents of the atmosphere• of known chemical composition after it is dried, or ignited.

Example of selective reagent:AgNO3: precipitates from acidic solution, Cl-, Br-, I-, and SCN

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Example of specific reagent:Dimethylglyoxime: specifically precipitates Ni2+ from alkaline solution

Inorganic Precipitantsg p

Form slightly soluble salts or hydrous oxidesith the anal tewith the analyte.Most inorganic reagents are not very selective.T o common inorganic precipitating agentsTwo common inorganic precipitating agents

are,silver nitrate which is used to precipitatesilver nitrate, which is used to precipitatehalide ions such as chloride,barium chloride which is used tobarium chloride, which is used to precipitate sulfate ion.

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Common Inorganic Precipitants

Precipitant Analyte Precipitate Formed

Precipitate weighedg

NH4OH Al Al(OH)3 Al2O3

Fe Fe(OH) Fe OFe Fe(OH)3 Fe2O3

HCl Ag AgCl AgCl

AgNO3 Cl AgCl AgCl

NaSO Ba BaSO BaSONaSO4 Ba BaSO4 BaSO4

BaCl2 SO42- BaSO4 BaSO4

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Organic Precipitants

Very useful precipitating agents for metals.Advantages by using organic reagents as a precipitant ;I f h l d i h i hi hIt forms chelate compounds with cations which are veryinsoluble in water. So, that metal ions may be quantitativelyprecipitated.p pThe organic precipitant often has a large molecular weight.Thus a small amount of metal may yield a large weight ofprecipitateprecipitate.Some of the organic reagents are fairly selective, yieldingprecipitates with only a limited number of cation. By

t lli h f t H d th t ti fcontrolling such factors as pH and the concentration ofmasking reagents, the selectivity of an organic reagent can beenhanced.

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The precipitates obtained with organic reagents are oftencoarse and bulky, and hence easily handle.

Examples:

8- hydroxyquinoline (oxine)

It can precipitates many elements but can be used for groupO

It can precipitates many elements but can be used for group separation by controlling pH.

Aluminium ion can be precipitated at pH 4. A higher pH isrequired to precipitate magnesium.

N+ Mg2+

Mg

O

N N2 + H+

N.. M

ogN N

OH

H

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Dimethyglyoxime

Principally used for determination of nickel. The reactionis,

Ni2+ + 2C H N O Ni(C H N O ) (p) + 2H+Ni2+ + 2C4H8N2O2 Ni(C4H7N2O2)2 (p) + 2H+

Dimethyglyoxime

This precipitate is so bulky that only small amount ofnickel can be handled conveniently.

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