Analytical Chemistry PDF

# Analytical Chemistry - is the science of the characterization and measurement of chemicals and involve separating, identifying and determining the relative amounts of the components in a sample of matter - Chemical analysis is divided into two types: ## Qualitative analysis - Identify the chemi...

# Analytical Chemistry - is the science of the characterization and measurement of chemicals and involve separating, identifying and determining the relative amounts of the components in a sample of matter - Chemical analysis is divided into two types: ## Qualitative analysis - Identify the chemical identity of the species in the sample. ## Quantitative analysis: - Establishes the relative amount of one or more of the species or analytes(sample) in numerical terms. ### Quantitative analysis: - **Gravimetric analysis** - The mass of the analyte(sample) is determined. - **Volumetric analysis** - The volume of a solution containing sufficient reagent to react completely with the analyte is measured. - **Instrumental analysis** - Based on the measurement of optical or electrical or thermal properties ## Gravimetric analysis - Gravi - Metric - Weighing - Measure - Most accurate analytical technique - It is an ABSOLUTE method. - Precise methods of macro quantitative analysis - Possible sources of errors can be checked ### Precipitation Gravimetry - In precipitation gravimetry an insoluble compound forms when we add a precipitating reagent, or precipitant, to a solution containing our analyte. - Any reaction generating a precipitate can potentially serve as a gravimetric method. - Most precipitation gravimetric methods were developed in the nineteenth century, or earlier, often for the analysis of ores ## Theory and Practice - All precipitation gravimetric analysis share two important attributes: - **First**, the precipitate must be of: - low solubility - high purity - known composition if its mass is to accurately reflect the analyte's mass. - **Second**, the precipitate must be easy to separate from the reaction mixture. ## Solubility Considerations: - To provide accurate results, a precipitate's solubility must be minimal. The accuracy of a total analysis technique typically is better than ±0.1%, which means that the precipitate must account for at least 99.9% of the analyte - We can minimize solubility losses by carefully controlling the conditions under which the precipitate forms. This, in turn, requires that we account for every equilibrium reaction affecting the precipitate's solubility. - For example, we can determine Ag gravimetrically, by adding NaCl as a precipitant forming a precipitate of AgCl. $Ag (aq) + Cl (aq) → AgCl(s)$ - If this is the only reaction we consider, then we predict that the precipitate's solubility, S, is given by the following equation: $SAgCl = [Ag+] = Ksp / [Cl-]$ - Solubility equation suggests that we can minimize solubility losses by adding a large excess of Cl - In fact, addition of a large excess of Cl- increases the precipitate's solubility. $Ag (aq) + Cl-(aq) = AgCl(s)$ $log K1 = 3.70$ $AgCl(aq) + Cl-(aq) → AgCl2-(aq)$. $log K2 = 1.92$ $AgCl2-(aq) + Cl-(aq) = AgCl32-(aq)$. $log K3 = 0.78$ - As we add NaCl to a solution of Ag+, the solubility of AgCl initially decreases because of reaction: $Ag+(aq) + Cl-(aq) → AgCl(s)$ - At higher concentrations of Cl-, reactions (2 & 3) increase the solubility of AgCl. Clearly the equilibrium concentration of chloride is important if we want to determine the concentration ofsilver by precipitating AgCl. In particular, we must avoid a large excess of chloride. ## Procedure for gravimetric analysis - Analyte (or) Sample ↓(selectively convert) - Precipitate form (Insoluble form) ↓ - Filtration (Separate the precipitate) ↓ - Drying ↓ - Igniting ↓ - Weighing ↓ - Calculations ## PRINCIPLE & STEPS INVOLVED INGRAVIMETRIC ANALYSIS - Preparation of the sample solution - Precipitation process - Digestion (or) Ostwald ripening - Filtration → Washing → Drying → Igniting → Weighing - Calculation ## Preparation of sample solution - **Factors to be considered during preparation of sample solution:** - Volume of solution during precipitation - pH - Temperature ### Volume of solution during precipitation - Solution condition (volume) must be adjusted to maintain low solubility of precipitate - Excess volume of solution during precipitation leads the precipitate to be coagulate. ### pH - Influence the solubility of the analytical precipitate - Possibility of interference from other substance **Example:** Calcium oxalate precipitate; soluble in acidic medium as oxalate ion combines with hydrogen ion at low pH to form soluble oxalic acid (weak acid). - But insoluble in basic medium **Example:** 8 - Hydroxy quinoline: precipitating reagent 8-Hydroxy quinoline (oxine) used to precipitate large number of elements. - In pH 4 condition; - aluminum ions can be easily precipitated. - For magnesium ion; - Concentration of anion form of oxine is too low at pH 4 to precipitate magnesium ion. - Higher pH is required to shift the ionization step and precipitate the magnesium. - If the pH is too high magnesium hydroxide will precipitate causing interference. ### temperature - Increase the temperature Increase solubility of precipitate; precipitate will dissolve or coagulate & Difficult to filter ## Gravimetric calculations: $Gf = \dfrac{formal \ weight \ of \ analyte(\frac{a}{mol})}{formal \ weight \ of \ precipitate(\frac{b}{mol})} \times \dfrac{a(mole analyte)}{b(mole ppt)}$ $Gf = \dfrac{analyte(g)}{ppt(g)}$ $analyte(g) = ppt(g) \times Gf$ %$analyte = \dfrac{analyte(g)}{sample(g)} \times 100$ %$analyte = \dfrac{ppt(g) \times Gf}{sample(g)} \times 100$ ## Determine Gf: - $Cl2 → AgCl(s)$ $Gf = \dfrac{1}{2}Cl2 / AgCl$ $Gf = Cl2/2AgCl$ $Gf = 35.5 \times 2 / 2(108 + 35.5) = 0.25$ - $AlCl3 → AgCl(s)$ $Gf = \dfrac{1}{3}AlCl3 / AgCl$ $Gf = AlCl3 / 3AgCl$ - $I → Hg5(106)2$ $Gf = 21/Hg5(106)2$ ## Problems: - Determine the gravimetric factors in term of symbols for the determination of: - (a) Aluminum as its hydroxyquinolate, Al → AI(C9H60N)3 - (b) Phosphorus as phosphomolybdic anhydride, P → P2M024078 - (c) Potassium as its chloroplatinate, K → K2PtCl6 - (d) Sulfur as barium sulfate, S→BaSO4 - (e) Nickel as nickel dimethylglyoxime, Ni → Ni(C4H7O2N2)2 ## Example: - In an organic sample (0.352g) phosphorous was dissolved and converted to Mg2P2O7 precipitate (0.223g). Calculate the percentage of %P in the original sample. $2P(Fw,31g) → Mg2P2O7(Fw, 222.6g)$ $Gf = 2 × 31/222.6 = 0.2783$ %$analyte = ppt(g) × Gf / sample(g) × 100$ %$P = (0.223 × 0.2783 / 0.352) × 100 = 17.1%$ ## Example: - When a sample of impure potassium chloride (0.4500g) was dissolved in water and treated with an excess of silver nitrate, 0.8402 g of silver chloride was precipitated. Calculate the percentage KCI in the original sample. $KCl (Fw, 74.50g) AgCl (Fw, 143.50g)$ $Gf = \dfrac{74.50}{143.50} = 0.519$ %$KCl = (mass \ of \ AgCl \times Gf/mass \ of \ KCl) \times100$ $=(0.8402 \times 0.519/0.4500) \times 100 = 96.90%$ ## Some of inorganic precipitating agents | Precipitate | Analyte and (formed precipitate, weighed) | |-------------|:---------------------------------------------------| | NH3(aq) | Be (BeO), Al (Al2O3), Cr(Cr2O3), Fe(Fe2O3) | | (NH4OH) | Sn (SnO2), Zr (ZrO2) | | H2S | Zn (ZnS → ZnO), As (AS2S3 → As2O3 or As2O5), Bi (BizS3) | | (NH4)2HPO4 | Mg (Mg2P2O7), Zn(Zn2P2O7), Cd(Cd2P2O7) | | H2SO4 | Sr, Cd, Pb, Ba (all as sulphate) | | H2PtCl6 | K (K2PtCl6) | | HCI | Ag (AgCl), Hg(Hg2Cl2) | | AgNO3 | CI(AgCl), Br(AgBr), I(AgI) | | BaCl2 | SO42- (BaSO4) | | (NH4)2S | Hg (HgS) | | ΗΝΟ3 | Sn4+ (SnO2) | | H5106 | Hg (Hg5(106)2 | | NaCl, Pb(NO3)2 | F (PbCIF) | | MgCl2, NH4Cl | PO43- (Mg2P2O7) | | Ca2+ | H2C2O4 CaCO3 or CaO | ## Volumetric analysis: - Is a general term for a method in quantitative chemical analysis in which the amount of a substance is determined by the measurement of the volume that the substance occupies. - It is commonly used to determine the unknown concentration of a known reactant. - Volumetric analysis is often referred to as titration. What is the meaning of Titration? - Titration is a common laboratory method of quantitative chemical analysis that is used to determine the unknown concentration of a known reactant. - Because volume measurements play a key role in titration, it is also known as volumetric analysis. - A reagent of a know concentration called the titrant or titrator, (a standard solution) volume is used to react with a solution of the analyte or titrand, whose concentration is not known using a calibrated burette. - | Titrant | Analyte | |-------------------------|--------------------| | (a standard solution) | (titrand) | | Known concentration solution | the unknown concentration solution | - (Burette) the titrant (a standard solution) Known concentration solution - (Conical flask) the analyte (titrand) the unknown concentration solution - Titrant solution (in a burette) is delivered slowly to the reaction flask with shaking. - Delivery of the titrant is called a titration. - The titration is complete when equivalent titrant has been added to react with all the analyte according to the balanced titration reaction equation. - This is called the equivalence point - The reaction is generally carried out in a conical flask containing the liquid or dissolved sample (the analyte). - During the titration, two important stages are known. - An equivalence point: - a point at which the added titrant is chemically equivalent to the analyte in the sample. - Or is a point where the chemical reaction comes to an end. (theoretical point). - **Endpoint:** - is the point at which the titration is complete, as determined by an indicator. - Or is the point where the colour change occurs in a system; (experimental point) - **An indicator** is often added to the reaction flask (conical flask) to signal when all the analyte has reacted. The titrant volume where the signal is generated is called the end point volume (Veq.p) - There are different methods to determine the equivalence point include: - **Color change** that can be detect by naked eye : - In some reactions; redox titrations (oxidation-reduction titrations) the solution changes color without any added indicator. $MnO4 + 5 Fe2+ + 8H+ → Mn2+ + 5 Fe3+ + 4 H2O$ (violet) (colorless) - **Specific indicators :** - starch forms blue complex with iodine (I) and thiocyanate (SCN¯). - starch forms red complex with iron (III). - **Equilibrium Indicators:** - These indicators are found in two forms of different colors and depend only on the change in the physical property, such as: pH (acid - base titration) : HIn ↔ H+ + In (color A) (color B) - potential (redox titration): Inox + ne ↔ Inred (color A) (color B) - **Measured properties :** - an instrument to detect the end point (not noticed by the naked eye). - such as: - electrical conductivity of the solution - absorbance of electromagnetic radiation. - pH meter ## Acid - Base Indicator: - Is a molecule that changes color based on pH. - Indicators are weak acids that lose a proton (causing the color change) when [OH-] reaches a certain concentration. $HIn + OH → In + H2O$ - Choose an indicator that changes color at the equivalence point - End Point occurs when the indicator changes color. - If you have chosen the wrong indicator, the end point will be different than the equivalence point. - Phenolphthalein, a Common Indicator <start_of_image> diagrams are not included in markdown output