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Chemistry Notes
Unit – II
Instrumental Methods of Analysis
Introduction: REPRESENTATION OF A DANIEL CELL:- i) Zn(s) | ZnSO 4 (sol) || CuSO 4 (sol) |Cu (s) ii) Zn-| ZnSO 4 (aq.) || CuSO 4 (aq.) | Cu+ iii) Zn|Zn2+||Cu2+|Cu iv) Zn|Zn2+^ (1M) ||Cu2+|Cu (1M)
The Nernst Equation:- E cell = E^0 cell - (RT/nF) lnK
Ecell = cell potential under nonstandard conditions (V) E^0 cell = cell potential under standard conditions R = gas constant, which is 8.31 (volt-coulomb)/(mol-K) T = temperature (K) n = number of moles of electrons exchanged in the electrochemical reaction (mol) F = Faraday's constant, 96500 coulombs/mol K = reaction quotient / constant Sometimes it is helpful to express the Nernst equation differently: Ecell = E^0 cell - (2.303 x RT/nF) log K at 298K, Ecell = E^0 cell - (0.0591 /n) log K Application of Nernst equation in electrochemistry:-
- The potential of an electrode and EMF of a cell can be calculated at any temperature and concentration.
- If potential of an electrode is known, the concentration of the reactant can be calculated.
- The concentration of a solution in the galvanic cell can be determined.
- The pH of a solution can be calculated by measuring the EMF. Types of Electrodes:- Reference electrodes & Indicator electrode Reference electrodes are those electrodes whose standard potential is known or strength is known. Ideal reference electrodes are:
- It is reversible and obeys the Nernst’s equation.
- Which exhibits a potential that is constant with time.
- Which returns to its original potential after being subjected to small currents.
In general, there are two types of reference electrodes, A) Primary reference electrode - e.g., Hydrogen electrode B) Secondary reference electrode – e.g. Calomel electrode , silver-silver chloride electrode. Difficulties in using standard hydrogen electrode -
- Platinum used in it, gets poisoned by absorption of impurities from the solution and the gas, hence it does not behave reversible for longer period.
- Because of presence of oxidizing agent, unsaturated organic compounds, alkaloids etc. in solution alters the potential.
- It cannot be used in presence of Cu, Ag, Au etc.
- It is difficult to prepare and maintain pure H 2 gas at 1atm pressure and concentration of electrolyte 1 M. So, we use secondary reference electrode which is convenient to handle. Secondary Reference Electrode: Calomel Electrode
Construction:
Calomel electrode consists of a narrow glass tube at the bottom of which is liquid mercury, above it is a paste of Hg – Hg 2 Cl 2 (calomel) and remaining portion of glass tube is filled with 1 N or 0.1 N or saturated solution of KCl. The potential of the calomel electrode depends upon the concentration of KCl solution. Pt. wire dipping into the mercury layer, is used to make electrical contact. Calomel electrode is represented as: Hg | Hg 2 Cl 2 (s) | KCl sat.
Reaction:-
Hg 2 Cl 2 + 2e-^2 Hg + 2 Cl- Advantages of Calomel electrode:-
- The construction is very simple.
- The cell potential is reproducible
- Potential does not vary for much long time. Disadvantages of Calomel electrode:- 1) It should not be used above 800 C, because calomel disproportionate at very high temp. into mercury and mercuric chloride. 2) It involves handling of poisonous Hg and Hg 2 Cl 2 Indicator Electrode (Glass Electrode):-
- Principle When two solutions of different [H+] are separated by a thin glass membrane, a potential difference is developed at the two surfaces of membrane. The potential difference developed is proportional to the difference in [H+] of the two solutions. The glass membrane acts as ion exchanger i.e. exchange of Na+^ of glass with H+^ of solution.
- Construction: A glass electrode is made of a long glass tube with a thin walled glass membrane bulb at the bottom. A Pt wire or AgCl coated Ag wire is dipped in the 0.1 M HCl solution in the bulb.
ii)Heterogeneous electrode consisting solid crystalline material incorporated with polymer like PVC or silicon. c) Gas-sensing electrodes : These electrodes are useful to analyze gases such as NH 3 , NO 2 , SO 2 , CO 2 and H 2 S. A nitrate ion responsive electrode is for NO 2 while a sulphide-ion selective electrode is for H 2 S. The microporous membrane is hydrophobic, made from polypropylene or any other fluorocarbon which allows only dissolved gases to pass through. The electrodes Ag/AgCl and glass pH-electrode are dipped in the inner solution. Area of membrane being small and volume of liquid being less, it reaches equilibrium with the test solution rapidly. d) Enzyme based membrane: These electrodes are used to convert substances in solution into ionic products which are measured using ion selective electrode. The enzyme is immobilized at the surface of electrode. Eg. Enzyme base membrane for determination of Urea.
Construction: Enzyme urease is incorporated into polyacrylamide gel which allow to set on bulb of glass electrode and is field in its position by nylon gauze. Working: When electrode is immersed in solution containing urea, NH 4 +^ ions are produced which diffused through gel. CO(NH 2 ) 2 + H 2 O+ 2H+^ 2NH 4 +^ + CO 2 Boundary potential developed due to difference in concentration of NH 4 +^ on either side of membrane. This potential is measured using glass electrode as a reference electrode. A) CONDUCTOMETRY: INTRODUCTION Important laws, definitions and Relations used in conductometry Ohm’s Law:- “The current flowing through a given solution is directly proportional to the voltage (potential difference) between the two ends of the conductor through which the current is flowing”. Mathematically it is written as, I α V or I α E E=RI (Where I = current strength, E = V = potential difference) ,R = Proportionality constant called resistance i.e. R = E / I OR V / I Unit of resistance (R) = Volts/ampere, therefore its unit is Ω (ohm).
Resistance (R): The tendency of a material to stop the flow of current is known as resistance. It is
measured in ohms (Ω).
According to Ohm's law, the resistance offered by a substance is directly proportional to
its length ( l ), but inversely proportional to its cross sectional area (A). cm^2
ρ is the proportionality constant and is known as specific resistance or resistivity of solution of length 1 cm and area of cross section 1 cm^2 , then R = ρ. In other words, ρ is the resistance of 1 c.c. of solution.
Its unit is ohm cm 𝜌 = 𝑅
𝐴
cm 2
= ohm. cm
𝑙 cm
Specific conductance / Conductivity (k): Reciprocal of specific resistance ρ is called as specific
conductance and it is the conductance of one cm^3 of solution, denoted by k (kappa).
It is measured in ohm-^1 .cm-^1 = mho. cm-^1 in (C.G.S system); or Siemens m-^1 (S.I system)
Equivalent conductivity: Conductance of solution by solution containing one gm-equivalent of an electrolyte, dissolved in volume V ml solution is known as Equivalent conductance and it is denoted by 𝖠. As one c.c. or ml of solution has conductance equal to (k) specific conductance, therefore equivalent conductance and specific conductance are related as, 𝖠 = kV (where V is ml of solution containing one gm-equivalent of electrolyte) If C is the concentration of solution as gm-equivalent per litre (normality), then volume V of the solution in ml containing 1 gm-equivalent will be 1000/C. Unit of 𝖠 is ohm–^1 cm^2 per equivalent. Molar conductivity: Conductance of solution by solution containing one gm-mole of electrolyte is known as molar conductance and it is denoted by μ. where M is concentration of solution in moles/litre. Unit of μ is ohm–^1 cm^2 per mole. Cell constant: - The ratio of the distance between the electrodes ( l) to the cross sectional area (A) of the electrodes is known as cell constant.
b) Weak Acid versus Strong Base Consider weak acid like acetic acid titration against a strong base like NaOH. In the beginning conductance of acetic acid is low and it further decreases due to depression in its dissociation by the common ion formed during early stage of neutralization. CH 3 COOH + Na+^ + OH CH 3 COO–^ + Na+^ + H 2 O After that the conductance increases, slowly due to increasing amount of completely dissociating, salt sodium acetate formed progressively upto equivalence point. Conductance at equivalence point is completely due to sodium acetate. After that, conductance increases faster due to excess of Na+^ and OH ions added (Refer the plot) from burette. 1 ml 1 N NaOH ≡ 60 mg acetic acid From the end point or equivalence point volume, normality of NaOH, we can calculate amount of acetic acid in solution. c) Weak Base Against Strong Acid Consider strong acid (HCl) from burette, against weak base (NH 4 OH) in flask. Reaction during titration is, NH 4 OH + H+^ + Cl–^ NH 4 +^ + Cl–^ + H 2 O Initially there is low conductance by the weak electrolyte but during titration, there is formation of strong electrolyte NH 4 Cl, therefore, conductance goes on increasing, upto equivalence point. After equivalence point, the conductance increases very rapidly because of fast conducting H+^ Cl– added remains unreacted in the titration mixture. Calculation: Plot a graph of conductance Vs ml of acid added from burette. From the equivalence point noted from graph, normality of HCl, we can calculate amount of base titrated. d) Weak Acid with a Weak Base: Consider weak acid (CH 3 COOH) in flask is titrated, against weak base (NH 4 OH) from burette. Reaction during titration is, CH 3 COOH + NH 4 OH CH 3 COO–^ NH 4 +^ + H 2 O The nature of curve before the equivalence point is similar to the curve obtained by titrating weak acid against strong base. After the equivalence point, conductance virtually remains same as the weak base which is being added is feebly ionized and, therefore, is not much conducting. Precipitation Titration : Precipitation titrations can be carried out conveniently by conductivity measurements, e.g. KCl versus AgNO 3 is added from burette and conductance of KCl solution observed at various occasions.
K+^ + Cl–^ + Ag+^ NO– 3 K+^ + NO– 3 + AgCl↓ The conductance of KCl decreases slowly upto equivalence point because greater mobility Cl are replaced by lower mobility NO 3 – ions. Because conductance difference in them is not large therefore conductance decreases slowly upto equivalence point. After that conductance increases rapidly, due to addition of Ag+^ and NO– 3 ions from burette. Calculation : Plot a graph of conductance Vs ml of titrant. 1 ml 1N AgNO 3 ≡ 35.5 mg Cl–^ or 74.5 mg KCl Equivalence point of titration is known from graph. From the known normality of AgNO 3 , equivalent point volume, we can calculate amount of Cl–^ or KCl in solution. Advantage of conductometric titrations:-
- This method can be used with very dilute solutions
- This method can be used with coloured or turbid solutions in which end point cannot be visible.
- This method can be used in which there is no suitable indicator is found to work satisfactorily can be successfully titrated by this method. B) pH-Metry: INTRODUCTION The concept of pH was first introduced by Danish chemist Sorensen. pH is a measure of the hydrogen ion concentration of a solution. Solutions with a high concentration of hydrogen ions have a low pH and solutions with low concentrations of H+^ ions have a high pH. pH= - log[H+] For example is the H+^ concentration is very low, lets say about 0.0000001M, then the pH is pH= - log[.0000001] which is the same as - log[1 X 10 - 7] The term log [1 X 10 -^7 ] = - 7 ; Thus - (-7) = 7 Solutions with a pH below 7.0 are termed acidic and solutions with a pH above 7.0 are termed basic. A pH value is a number from 0 to 14 , with 7 as the middle (neutral) point. Values towards 1 indicate more acidity while values towards 14 being the most alkaline. It is a logarithmic scale in which two adjacent values increase or decrease by a factor of 10. For example, a pH of 3 is ten times more acidic than a pH of 4, and 100 times more acidic than a pH of 5. Similarly, a pH of 9 is 10 times more alkaline than a pH of 8, and 100 more alkaline than a pH of 7. pH of blood is 7.4 and pure water is 7. A pH meter is an electronic device used for measuring the pH (acidity or alkalinity) of a liquid. Buffer solution: Solution which resists change in pH even on addition of small aquantity of acid or base. Buffer solutions are of three types : a) Acidic buffer b) Basic buffer c) Neutral buffer a) Acidic buffer: Acidic buffer solution contains equimolar quantities of a weak acid and its salt with strong base. For example: acetic acid, CH 3 COOH and sodium acetate I.e. CH 3 COONa. A solution containing equimolar quantities of acetic acid and sodium acetate maintains its pH value around 4.74. b) Basic buffer: Basic buffer solution contains equimolar quantities of a weak base and its salt with strong acid. For example: ammonium hydroxide i.e. NH 4 OH and ammonium chloride I.e. NH 4 Cl. A solution containing equimolar quantities of ammonium hydroxide and ammonium chloride maintains its pH value around 9.25.
- It has many applications, i.e. it can be used for weak acid and weak bases, redox, precipitation, or complex titrations
- More accurate results are obtained because the end point is determined graphically.
Uses / Applications of pH metry:
- a swimming pool must be constantly maintained by checking its pH. If the water becomes too acidic or too basic, must be balanced
- pH also is important when it comes to the body. Whenever someone gets heartburn, stomach acid builds up and causes deep pain near the opening of the stomach. When one takes antacid tablets (a base), the build up of acid is neutralized in the stomach, causing relief to the sufferer.
- The blood is also important when dealing with pH. Usually, the pH of blood should only be slightly basic. A fluctuation in the pH of the blood may result in serious injury to vital organs in the body.
- The pH test is an important preliminary test analysis of water.
- Also use in analysis of soil. Most plants can tolerate a wide pH range in solution culture, but they cannot tolerate a wide range of acidity in the soil.
- pH measurent also use in medical field.
