Gravimetric analysis

Common Ion Effect and Solubility Product

Common Ion Effect

• Decreases solubility of solids by adding common ions, shifting equilibrium.
• Regulates buffers by shifting the pH of the solution upon addition of conjugate ions.
• Suppresses the degree of dissociation of weak electrolytes when another electrolyte containing a common ion is added.
• Offers a valuable method to control the concentration of the desired ion given by the weak electrolyte.
• Must be considered when determining solution equilibrium upon addition of ions already present.

Solubility

• Refers to the amount of material that can be dissolved in a particular solvent.
• A solution reaches its solubility limit when no more material can be dissolved, becoming saturated.

Chemical Equilibrium

• A state where there are no net physical or chemical changes between reactants and products.
• The rates of the forward and reverse reactions are equal.

Solubility Equilibrium

• A state of chemical equilibrium between a solid compound and its dissolved form in a solution.
• Established when the rates of migration between the solid and aqueous phases are equal.

Common Ion Effect on Solubility

• Adding a common ion decreases solubility, shifting the reaction toward the reactants.
• Causes the equilibrium to shift left, toward the reactants, resulting in precipitation.

Solubility Product

• In a saturated solution of a sparingly soluble electrolyte (salt), the product of ionic concentrations, raised to proper powers, is constant at a given temperature, denoted by $K_{sp}$ or S.
• Compared to the ionic product, conclusions are as follows:
  1. Ionic product = $K_{sp}$: solution is saturated, no precipitation.
  2. Ionic product < $K_{sp}$: solution is unsaturated, no precipitation.
  3. Ionic product > $K_{sp}$: solution is supersaturated, precipitation.
• $K_{sp}$'s are equilibrium constants in heterogeneous equilibria (i.e., between two different phases).
• If several salts are present, they all ionize in the solution.
• If salts contain a common cation or anion, they contribute to the concentration of the common ion.
• Contributions from all salts must be included in the calculation of the concentration of the common ion.
• Consideration of charge balance or mass balance both leads to the same conclusion.

Common Ions Example

• When NaCl and KCl are dissolved in the same solution, the $Cl^-$ ions are common to both salts.
• Examples:
  • $NaCl \rightleftharpoons Na^+ + Cl^-$
  • $KCl \rightleftharpoons K^+ + Cl^-$
  • $CaCl_2 \rightleftharpoons Ca^{2+} + 2Cl^-$
  • $AlCl_3 \rightleftharpoons Al^{3+} + 3Cl^-$
  • $AgCl \rightleftharpoons Ag^+ + Cl^-$

Gravimetric Methods of Analysis

General Definition

• Quantitative methods based on determining the mass of a pure compound to which the analyte is chemically related.

Types of Gravimetric Methods

• Precipitation gravimetry: The analyte is separated as a precipitate and converted to a compound of known composition for weighing.
• Volatilization gravimetry: The analyte is separated by converting it to a gas of known chemical composition that can be weighed.
• Electrogravimetry: The analyte is separated by deposition on an electrode via electrical current.

Precipitation Gravimetry

• Analyte converted to a sparingly soluble precipitate.
• The precipitate is filtered, washed, and heated to form a product of known composition.
• Example: Determining $[Ca^{2+}]$ in water:
  • $2NH<em>3(aq) + H</em>2C<em>2O</em>4(aq) \rightarrow 2NH<em>4^+(aq) + C</em>2O_4^{2-}(aq)$
  • $Ca^{2+}(aq) + C<em>2O</em>4^{2-}(aq) \rightarrow CaC<em>2O</em>4(s)$
  • $CaC<em>2O</em>4(s) \xrightarrow{\Delta} CaO(s) + CO(g) + CO_2(g)$
• The precipitate is cooled and weighed to determine the mass.

Requirements for Gravimetric Analysis

• Complete Precipitation
  • The precipitate should form quantitatively, leaving negligible analyte in solution.
  • Achieved by choosing the right reagent and reaction conditions (pH, temperature, interfering substances).
• Low Solubility (Low $K_{sp}$)
  • The precipitate should have very low solubility in water to ensure maximum precipitation.
  • Example: $BaSO<em>4$ has an extremely low $K</em>{sp}$, making it suitable for sulfate determination.
• Purity of the Precipitate
  • The precipitate should be free from impurities.
  • Washing with deionized water removes adsorbed foreign ions.
• Easily Filterable Precipitate
  • The precipitate should form large, well-defined crystals (rather than fine colloidal particles) to facilitate filtration.
  • Adding an electrolyte or digesting helps in forming large crystals.
• Stable Composition
  • The precipitate must not decompose upon heating or drying.
  • If needed, it should be converted into a stable oxide (e.g., heating calcium oxalate to form calcium oxide).

Gravimetric Determination of Copper as CuSCN

• Gravimetric analyses are highly precise due to the accuracy of contemporary analytical balances.
• High purity of the analyzed element compound should be obtained.
• The reaction must be stoichiometric.
• The weighed compound should be non-hygroscopic and stable in air.
• It is better if it has a relatively high molecular mass for more precise weighing.

Examples of Gravimetric Analyses

• Analysis of iron precipitated as $Fe(OH)<em>3$ and heated to oxide $Fe</em>2O_3$ at ca. 800°C.
• Analysis of barium precipitated as $BaSO_4$ and heated to ca. 500°C.
• Analysis of nickel precipitated as a complex with dimethylglyoxime, dried at 110°C.

Copper Determination as CuSCN

• Copper is precipitated as $CuSCN$ (solubility product $K_s = 10^{-12.7}$
• $Cu^{2+}$ ions are reduced to $Cu^+$ before precipitation using $SCN^-$.
• $2Cu^{2+} + HSO<em>3^- + H</em>2O \rightarrow 2Cu^+ + HSO_4^- + 2H^+$
• $Cu^+ + SCN^- \rightarrow CuSCN$
• Most thiocyanates of other metals are soluble, including those of Bi, Cd, As, Sb, Sn, Fe, Ni, Co, Mn, and Zn.
• Tartaric acid is added to prevent hydrolysis of Bi, Sb, or Sn salts.

Conditions for the Experiment

• Slight acidity with respect to $HCl$ or $H<em>2SO</em>4$
• The presence of a reducing agent, such as $H<em>2SO</em>3$ or $NH<em>4HSO</em>3$, to reduce $Cu(II)$ to $Cu(I)$
• A slight excess of $NH_4SCN$. A large excess increases the solubility of the copper thiocyanate due to complex formation.
• The absence of oxidizing agents.
• The precipitate is curdy and readily coagulates by boiling.
• Wash with dilute ammonium thiocyanate with an addition of $H<em>2SO</em>3$ or $NH<em>4HSO</em>3$ to avoid oxidation of $Cu(I)$.

Chemicals

• 5-6% aqueous solution of $NH<em>4HSO</em>3$
• Freshly prepared 10% aqueous solution of $NH_4SCN$

Procedure

1. Dissolve the sample in water, add a few drops of $2M HCl$ and $25 mL$ of $NH<em>4HSO</em>3$ solution.
2. Dilute to 150-200 mL, heat nearly to boiling, and add $NH_4SCN$ solution slowly, stirring constantly.
3. Filter through a glass crucible.
4. Wash the precipitate at least 10 times with a cold solution of $NH<em>4SCN$ and $NH</em>4HSO_3$ in water.
5. Dry the crucibles at 110°C for 90 min or more, cool in a desiccator, and weigh.

Calculations and Report

• Report masses of empty and full filters, calculated masses of the precipitated $CuSCN$, and the final result.
• The result is the mass of elemental copper in each sample (in grams) and the averaged weight percent of Cu in the alloy or mineral under examination.
• The ratio of molecular masses of Cu to $CuSCN$ is 0.5225.

Gravimetric Estimation of Nickel

• Nickel is precipitated as nickel dimethyl glyoxime by adding an alcoholic solution of dimethyl glyoxime $CH<em>3C(NOH)C(NOH)CH</em>3$ and a slight excess of aqueous ammonia solution.
• $Ni^{2+} + 2CH<em>3C(NOH)C(NOH)CH</em>3 \rightarrow Ni[CH<em>3C(NOH)C(NO)CH</em>3]_2 + 2H^+$
• When the pH is buffered between 5 and 9, the red chelate forms quantitatively.
• Chelation occurs due to the donation of electron pairs on the four nitrogen atoms.
• The solution is buffered with ammonia or citrate to prevent the pH from falling below 5.
• A slight excess of the reagent has no action on the precipitate, but a large excess should be avoided because of the possible precipitation of the reagent itself.
• The precipitate is soluble in free mineral acids, so avoid adding too large an excess of the reagent.
• The complex is slightly soluble in alcoholic solutions. Adding a small amount of chelating agents will minimize errors.
• The amount of reagent added is also governed by the presence of other metals such as cobalt, which form soluble complexes with the reagent.
• Nickel dimethylglyoximate is a very bulky precipitate. Therefore, the sample weight used in the analysis must be carefully controlled.
• The compactness of the precipitate is improved by adjusting the pH to 3 or 4, followed by the addition of ammonia solution.
• A slow increase in the concentration of ammonia in the solution causes a gradual increase in the pH and results in a denser precipitate.
• Once the filtrate has been collected and dried, the nickel content of the solution is calculated stoichiometrically from the weight of the precipitate.
• Molecular weights:
  • Dimethylglyoxime: 116.12 g/mol
  • $Ni(DMG)_2$ complex: 288.9155 g/mol
  • Nickel: 58.6934 g/mol

Determination of Chloride Ion Concentration by Gravimetry

Safety

• Wear lab coats, safety glasses, and enclosed footwear.
• Silver nitrate solution causes staining of skin and fabric (chemical burns). Rinse spills with water immediately.
• Concentrated nitric acid is very corrosive; take great care.

Introduction

• This method determines the chloride ion concentration of a solution by gravimetric analysis.
• A precipitate of silver chloride is formed by adding a solution of silver nitrate to the aqueous solution of chloride ions.
• $Ag^+ (aq) + Cl^- \rightarrow AgCl(s)$
• The precipitate is collected by careful filtration and weighed.
• Heating the reaction solution before filtering causes the solid silver chloride particles to coagulate.
• The precipitation is carried out under acidic conditions to avoid possible errors due to the presence of carbonate and phosphate ions, which would also precipitate with the silver ions under basic conditions.
• This method is best used on solutions with a fairly significant concentration of chloride ions, such as seawater.

Equipment Needed

• 100 ml volumetric flask
• Pipettes
• 250 ml conical flask
• Burette and stand
• Bunsen burner, tripod, and gauze
• Measuring cylinders
• Buchner funnel, filter paper, and a sidearm filtering flask

Solutions Needed

• Silver nitrate solution: (0.1 mol/L)
  • Dry 5 g of $AgNO_3$ for 2 hours at 100°C and allow to cool.
  • Accurately weigh about 4.25 g of solid $AgNO_3$ and dissolve it in 250 mL of distilled water in a conical flask.
  • Store the solution in a brown bottle.
• Nitric acid solutions:
  • A 6 mol/L solution of nitric acid is needed.
  • Dilute nitric acid solution: Add 1 mL of 6 mol/L $HNO_3$ solution to about 500 mL of distilled water.
• Methyl orange indicator

Method

Sample Preparation

• If the seawater contains traces of solid matter, it must be filtered before use.

Titration

1. Dilute seawater by pipetting a 50 ml sample into a 100 mL volumetric flask and making it up to the mark with distilled water.
2. Pipette a 20 ml sample of diluted seawater into a conical flask and add about 15 ml distilled water and 1 drop of methyl orange indicator. Add dilute nitric acid dropwise until the indicator turns pink, then add 1 mL of 6 mol/L nitric acid.
3. Add dropwise from a burette 55 mL of 0.1 mol/L silver nitrate solution. Allow the solution to stand for a minute and then test to see if it is completely precipitated by adding one drop of 0.1 mol/L silver nitrate solution. If more of the silver chloride precipitate forms add an additional 5 mL of the silver nitrate solution and retest for complete precipitation.
4. Heat the solution to boiling. Remove from heat and let stand in the dark for at least 1 hour until the precipitate coagulates.
5. Weigh the filter paper before filtering with the Buchner funnel and flask, and then use the equipment to filter the supernatant liquid (the liquid above the precipitate).
6. Wash the precipitate in the conical flask three times with a few ml of the very dilute nitric acid solution, pouring each washing through the Buchner funnel.
7. Finally, transfer the precipitate itself to the Buchner funnel washing any loose particles from the conical flask with a little distilled water.
8. Wash the precipitate on the Buchner funnel three times with a few mL of the very dilute nitric acid solution, and then wash the precipitate three times with a few mL of distilled water.
9. Carefully place the filter paper containing the precipitate on a watch glass and dry overnight. Weigh the dried filter paper and precipitate and calculate the weight of the dried precipitate using the known weight of the filter paper.

Result Calculations

1. Use the mass (in grams) of silver chloride in the dried precipitate to determine the number of moles of chloride ions in your sample.
   • $Ag^+ + Cl^- \rightarrow AgCl$
2. Calculate the concentration of chloride ions in the diluted seawater.
3. Calculate the concentration of chloride ions in the original (undiluted) seawater.
4. Calculate the concentration of sodium chloride in the seawater in mol/L, g/L, and g/100 mL (%).

Experiment 16: Estimation of Copper and Zinc in a Mixture by Gravimetry

Introduction

• In the estimation of copper and zinc in a given mixture, we first estimate copper as copper(I) thiocyanate (CuCNS) and then in the filtrate, we estimate zinc as zinc ammonium phosphate ($ZnNH<em>4PO</em>4$).
• The test solution is prepared by dissolving copper sulphate and zinc chloride in dilute HCl and making up the volume to 250 cm³ in a standard volumetric flask. 50 cm³ of this solution can be taken for estimation.

Principle

• Copper is precipitated as copper(I) thiocyanate by reducing Cu(II) to Cu(I) using sulphurous acid or ammonium hydrogen sulphite, followed by the addition of ammonium thiocyanate solution.
  • $2Cu^{2+}(aq) + HSO<em>3^-(aq) + H</em>2O(l) \rightarrow 2Cu^+(aq) + HSO_4^-(aq) + 2H^+(aq)$
  • $Cu^+(aq) + SCN^-(aq) \rightarrow CuSCN(s)$
• Zinc is precipitated as zinc ammonium phosphate by adding diammonium hydrogen phosphate solution to the filtrate.
  • $ZnCl<em>2(aq) + (NH</em>4)<em>2HPO</em>4(aq) \rightarrow ZnNH<em>4PO</em>4(s) + 2HCl(aq)$
• The precipitates are filtered, washed, dried, and weighed to calculate the concentrations of copper sulphate and zinc chloride in the solution.

Requirements

• Apparatus: Beakers, Bunsen burner, Desiccator, Filtration apparatus, Flask, Glass rod, Pair of tongs, Rubber policeman, Sintered crucible, Tripod stand, Wash bottle, Watch glass, Water bath, Wire gauze, Weighing bottle, Weight box
• Chemicals: Ammonium thiocyanate, Aqueous ammonia, Diammonium hydrogen phosphate, Hydrochloric acid, Methyl red indicator, Nitric acid, Ammonium hydrogen sulphite solution, Test Solution

Procedure

(A) Estimation of copper

1. Pipette 50 cm³ of test solution into a 500 cm³ beaker. Add 5 cm³ dilute HCl and 25 cm³ ammonium hydrogen sulphite solution to ensure a smell of $SO_2$.
2. Boil gently on a water bath. Remove the flame and add slowly 50 cm³ of 10% $NH_4SCN$ solution, stirring intermittently. Allow the solution to rest for 30-60 minutes.
3. Weigh an empty and cleaned sintered glass crucible. Filter the precipitate through this crucible by draining off the supernatant liquid first and then the precipitate.
4. Wash the precipitate with 2-3% $NH_4SCN$ solution. Check the filtrate for any precipitate formed. Wash the precipitate several times with 20% ethanol until the precipitate is free from $SCN^-$ ions. Preserve the filtrate for estimation of zinc.
5. Heat the sintered glass crucible at 100-120°C in a hot air oven for at least an hour.
6. Cool the crucible in a desiccator and then weigh. Repeat the process of heating, cooling, and weighing until a constant mass of the crucible with precipitate is obtained.

(B) Estimation of zinc

1. Evaporate the filtrate from copper estimation on a water bath to reduce the volume to nearly 100 cm³. Add 20 cm³ concentrated nitric acid and 15 cm³ concentrated HCl and again heat to dryness on a water bath in a fume hood.
2. To the residue, add 100 cm³ of distilled water and dissolve the contents by shaking. Add 2-3 drops of methyl red indicator and then add 10% aqueous ammonia solution until the smell of ammonia prevails and the color of the methyl red changes to yellow.
3. To the solution, add slowly 10% diammonium hydrogen phosphate with adequate stirring. Digest the precipitate on a water bath for at least 30 minutes.
4. Filter the precipitate through a previously weighed sintered crucible. Drain out the supernatant liquid first and then the precipitate.
5. Wash the precipitate with 2-3% diammonium hydrogen phosphate (DAHP) solution and finally with 50% alcohol to remove the excess of DAHP. Check if any phosphate ions are present in the washings.
6. Heat the crucible and the precipitate at temperature range 100-120°C in an hot air oven. Cool the crucible in a desiccator and then weigh it. Repeat heating, cooling, and weighing till a constant mass of the crucible with precipitate is obtained.

Observations

• i) 1st mass of empty crucible
• ii) 2nd mass of empty crucible
• iii) 1st mass of crucible + CuSCN
• iv) 2nd mass of crucible + CuSCN
• v) 1st mass of empty crucible
• vi) 2nd mass of empty crucible
• vii) 1st mass of crucible + ZnNH4PO4
• viii) 2nd mass of crucible + ZnNH4PO4

Calculations

Calculations for copper

• Mass of CuSCN obtained = iv) - ii) = w g
• From stoichiometry:
  $CuSO<em>4 \cdot 5H</em>2O = Cu^{2+} = CuSCN$
  $249.5 g = 63.54 g = 121.62 g$
• Therefore, w g of copper(I) thiocyanate = $\frac{249.5 \times w}{121.62}$ g of copper(II) sulphate.
• Concentration of copper(II) sulphate in test solution = $\frac{249.5 \times 1000 \times mass \ of \ CuSCN}{121.62 \times volume \ of \ test \ solution}$ gdm-3

Calculations for zinc

• Mass of $ZnNH<em>4PO</em>4$ obtained = viii) - vi) = w' g
• From stoichiometry:
  $ZnCl<em>2 = Zn^{2+} = ZnNH</em>4PO_4$
  $136.37 g = 65.37 g = 178.34 g$
• Therefore, w' g of $ZnNH<em>4PO</em>4 = \frac{136.37 \times w'}{178.34}$ g of $ZnCl_2$
• Concentration of $ZnCl_2$ in test solution = $\frac{136.37 \times w' \times 1000}{178.34 \times 50}$ g dm-3
• Concentration of $ZnCl<em>2 = \frac{136.37 \times 1000 \times mass \ of \ ZnNH</em>4PO_4}{178.34 \times Volume \ of \ test \ solution}$ gdm-3

Result

• Concentration of copper(II) sulphate in the test solution = ….. g dm-3
• Concentration of zinc chloride in the test solution = ….. g dm-3

METHOD #: 375.3

TITLE: Sulfate (Gravimetric)

• Analyte: Sulfate, SO4
• Instrumentation: Drying Oven

1.0 Scope and Application

• 1.1 This method is applicable to drinking, surface and saline water, domestic and industrial wastes.
• 1.2 This method is the most accurate method for sulfate concentrations above 10 mg/L. Therefore, it should be used whenever results of the greatest accuracy are required.

2.0 Summary of Method

• 2.1 Sulfate is precipitated as barium sulfate in a hydrochloric acid medium by the addition of barium chloride. After a period of digestion, the precipitate is filtered, washed with hot water until free of chloride, ignited, and weighed as BaSO4.
• 2.2 Preserve by refrigeration at 4°C.

3.0 Interferences

• 3.1 High results may be obtained for samples that contain suspended matter, nitrate, sulfite and silica.
• 3.2 Alkali metal sulfates frequently yield low results. This is especially true of alkali hydrogen sulfates. Occlusion of alkali sulfate with barium sulfate causes the substitution of an element of lower atomic weight than barium in the precipitate. Hydrogen sulfate of alkali metal acts similarly and decomposes when heated. Heavy metals such as chromium and iron, cause low results by interfering with complete precipitation and by formation of heavy metal sulfates.

4.0 Apparatus

• 4.1 Steam bath
• 4.2 Drying oven, equipped with thermostatic control.
• 4.3 Muffle furnace with heat indicator.
• 4.4 Desiccator
• 4.5 Analytical balance, capable of weighing to 0.1 mg.
• 4.6 Filter paper, acid-washed, ashless hard-finish filter paper sufficiently retentive for fine precipitates.

5.0 Reagents

• 5.1 Methyl red indicator solution: Dissolve 100 mg methyl red sodium salt in distilled water in a 100 mL volumetric flask and dilute to the mark with distilled water.
• 5.2 Hydrochloric acid, HCl, 1+1
• 5.3 Barium chloride solution: Dissolve 100 g BaCl · 2H O in 1 liter of distilled water. Filter through a membrane filter or hard-finish filter paper. One mL of this reagent is capable of precipitating approximately 40 mg SO4.
• 5.4 Silver nitrate-nitric acid reagent. Dissolve 8.5 g AgNO, and 0.5 mL conc. HNO, in 500 mL distilled water.

6.0 Procedure

6.1 Removal of silica: If silica concentration is greater than 25 mg/L

• 6.1.1 Evaporate sample nearly to dryness in a platinum dish on a steam bath.
• 6.1.2 Add 1 mL HCl solution (5.2), tilt dish and rotate until acid contacts all of the residue.
• 6.1.3 Continue evaporation to dryness.
• 6.1.4 Complete drying in an oven at 180°C.
• 6.1.5 If organic matter present, char over a flame.
• 6.1.6 Moisten with 2 mL distilled water and 1 mL HCl solution (5.2).
• 6.1.7 Evaporate to dryness on a steam bath.
• 6.1.8 Add 2 mL HCl solution (5.2).
• 6.1.9 Take up soluble residue in hot distilled water and filter.
• 6.1.10 Wash the insoluble silica with several small portions of hot distilled water.
• 6.1.11 Combine filtrate and washings.

6.2 Precipitation of barium sulfate

• 6.2.1 If necessary, treat clarified sample to remove interfering agents.
• 6.2.2 Adjust to contain approximately 50 mg SO, ion in a 250 mL volume.
• 6.2.3 Adjust acidity with HCl solution (5.2) to pH 4.5 to 5.0, using pH meter or orange color of methyl red indicator (5.1).
• 6.2.4 Add an additional 1 to 2 mL HCl solution (5.2).
• 6.2.5 For lower concentrations of sulfate ion fix the total volume at 150 mL.
• 6.2.6 Heat to boiling and, while stirring gently, add warm BaCl, solution (5.3) slowly, until precipitation appears to be complete; then add approximately 2 mL in excess.
• 6.2.7 If the amount of precipitate is small, add a total of 5 mL BaCl, solution (5.3).
• 6.2.8 Digest the precipitate at 80 to 90°C preferably overnight but for not less than 2 hours.

6.3 Filtration and Weighing

• 6.3.1 Mix a little ashless filter paper pulp with the BaSO, and filter at room temperature.
• 6.3.2 Wash the precipitate with small portions of warm distilled water until the washings are free of chloride as indicated by testing with silver nitrate-nitric acid reagent (5.4).
• 6.3.3 Dry the filter and precipitate.
• 6.3.4 Ignite at 800°C for 1 hour. DO NOT LET THE FILTER PAPER FLAME.
• 6.3.5 Cool in a desiccator and weigh

7.0 Calculation

• mg/L SO4= $\frac{mg \ BaSO_4 \times 411.5}{mL \ sample}$

8.0 Precision and Accuracy

• 8.1 A synthetic unknown sample containing 259 mg/L sulfate, 108 mg/L Ca, 82 mg/L Mg. 3.1 mg/L K, 19.9 mg/L Na, 241 mg/L chloride, 250 g/L nitrite N. 1.1 mg/L nitrate N and 42.5 mg/L alkalinity (contributed by NaHCO), was analyzed in 32 laboratories by the gravimetric method, with a relative standard deviation of 4.7% and a relative error of 1.9%.

Indian Standard: Method for Gravimetric Determination of Sulphates (IS:2317-1975)

Foreword

• This standard prescribes the method for the gravimetric determination of sulfates as barium sulfate.

Reagents

• Phenolphthalein Indicator Solution: Dissolve 0.1 g of phenolphthalein in 60 ml of rectified spirit and dilute with water to 100 ml.
• Sodium Hydroxide Solution: 4 percent (m/v).
• Concentrated Hydrochloric Acid.
• Barium Chloride Solution: 10 percent (m/v).
• Concentrated Nitric Acid.
• Concentrated Sulphuric Acid.

Procedure

• Take a clear solution of the material containing the sulfate. Neutralize acidic solutions with sodium hydroxide solution and alkaline solutions with concentrated hydrochloric acid to phenolphthalein indicator.
• To the neutral solution, add 4 ml of concentrated hydrochloric acid. Dilute to about 400 ml with water and heat to boiling. Remove from heat and add hot barium chloride solution in a fine stream with constant stirring. Add the reagent in excess of that required for precipitation.
• Place the beaker on a steam-bath for 4 hours and allow the precipitate to settle. Filter through a weighed sintered glass crucible (G-4), a Gooch crucible, or a filter paper (Whatman No. 42 or equivalent), by decantation retaining the precipitate in the beaker as far as possible. Wash the precipitate with hot water until the filtered washings are nearly free from chlorides and then transfer the precipitate fully onto the filter. Wash two or three times more with hot water.

Filtration has been done through a sintered glass crucible

• Dry the precipitate at 110 ± 5°C. Cool in a desiccator and weigh to constant mass. In the case of filtration through Gooch crucible, dry the precipitate and ignite it over a burner or in the muffle furnace at 600 to 700°C for half an hour. Cool in a desiccator and weigh to constant mass.

When filtration is carried out on a filter paper

• Dry the precipitate and ignite it over a burner or in the muffle furnace at 600 to 700C for half an hour. If the ignited precipitate smells of sulfides (BaS) or has a grey appearance, indicating the presence of carbon black, moisten it with one drop of concentrated sulfuric acid and evaporate to dryness on a hot plate in the crucible, fitted with its lid to avoid loss. Return to the furnace and ignite again at 800 +25°C for 15 minutes. Cool in a desiccator and weigh to constant mass Conduct a blank.

Calculation

• Sulphates (as SO4), percent by mass = $\frac{41.15 \times M<em>1}{M</em>2}$
• $M_1$ = mass in g of the precipitate, and
• $M_2$ = mass in g of the material contained in the solution taken for precipitation.

Experiment 4: Gravimetric Determination of Lead Chromate

Solutions

• Concentrated NaOH
• pH meter
• 0. 1% Nitric Acid (Conc. HNO3 = 76%, so about 1 mL conc./100 mL water)
• 0. 10 M chromium nitrate, Cr(NO3)3·9H2O (4 g/100 mL)
• 0. 12 M potassium bromate, KBrO3 (2 g/100 mL)
• Acetate buffer solution: 6 M in acetic acid, 0.6 M in sodium acetate.

Glassware

• 2 erlenmeyer flasks
• 2 porous porcelain filter crucibles
• suction adapters

Procedure

1. To eache flask add 4 ml of 1000 ppm Pb. Bring to 20 mL vol.. To the fifth flask add 20 ml of one of your soil sample digests. To one of the first four (lead standard) flasks add 4 mL of 1000 ppm Zn.
2. If necessary, neutralize 20mL solution with NaOH to pH 7 (use indicator paper) in a 100 mL beaker. Solution will be slightly cloudy.
3. Adjust the volume of sample to about 20 mL.
4. Add 10 mL Cr(NO3)3·9H2O solution and 10 mL of KBrO3. Solution will be clear blue.
5. Heat but do not boil for 30 minutes.
6. When solution is clear and yellow (= measure of the extent of chromic oxidation), add 10 mL of buffer and heat 5 more minutes.
7. Cool the mixture and filter off the lead chromate on a sintered glass or porous porcelain filter crucible.
8. Weigh your 6 filter crucibles.
9. Wash the precipitate with 2 or 3 small portions of 0.1% nitric acid.
10. Dry at 120C for 30 minutes.
11. Cool and weigh as PbCrO4 (Pb/PbCrO4 = 0.641108).

Gravimetric and Organic Analysis, Physical Constant Determination and Organic Preparation

I-Gravimetric Analysis

Estimation of Lead as Lead Chromate

Aim

• To estimate gravimetrically, the amount of lead present in the whole of the given solution of lead acetate or lead nitrate by precipitating it as lead chromate.

Principle

• Lead is precipitated as lead chromate in dilute acetic acid medium by the addition of Potassium chromate solution. The precipitated lead chromate is filtered in a sintered crucible, washed, dried at 120°C and weighed.
• $$Pb(CH3COO)2 + K2CrO4 \rightarrow PbCrO4 + 2CH3COOK