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![ The E.m.f of a metal immersed in a solution of its own
ions may be obtained from the Nerst equation as
follows; E = Eo + 0.0592/n * log C
Where Eo is the standard potential of the metal, n is
valency of the ions and c is the ionic concentration
e.g Calculate the emf of the galvanic cell Fe|
Fe2+
(0.200M)‖Ag+
(0.100M)|Ag
e.g The emf of the galvanic cell
Cu|Cu2+
(1.8x10-3
M)‖Ag+
(xM)|Ag is + 0.362. What is
the [Ag+
] in the silver half-cell?](https://image.slidesharecdn.com/electrochemicaltechniquesintroduction-250426192615-ddcf2265/85/Electrochemical-techniques_-Introduction-pptx-23-320.jpg)
![THEORY OF E M F & PH DETERMINATIONS
pH = -log [H+
]= log 1/[H+
]
Similarly, pOH = -log [-
OH] = log 1/ [-
OH]
Purified water ionizes in solution as follows;
2H2O H3O+
+ -
OH
The above equation may be simplified as;
H2O H+
+ -
OH
Applying the law of Mass Action;
[H+
] [-
OH] / [H2O] = Constant
Where [H+],[-OH] and [H2O] represent their concentrations
respectively](https://image.slidesharecdn.com/electrochemicaltechniquesintroduction-250426192615-ddcf2265/85/Electrochemical-techniques_-Introduction-pptx-24-320.jpg)
![PH OF A SOLUTION
Thus in pure water or neutral solution where [H+
] = [-
OH] =√Kw
= 10-7
gram ions per litre (25o
)
Also [H+
] [-
OH] = Kw = 10-14
Hence, pH + pOH = pKw = 14
One of the electrodes is referred to as the Indicator Electrode,
and the other as Reference Electrode.
The former must respond to pH changes whilst the other must
give a constant potential. Each forms a half cell.
Hydrogen electrode where n = 1,
E = Eo + 0.0592 * log [H+
]
Alternatively, E = EHo - 0.0592 * pH
But EHo is normally taken as zero, thus the above equation
becomes; E = - 0.0592 * pH](https://image.slidesharecdn.com/electrochemicaltechniquesintroduction-250426192615-ddcf2265/85/Electrochemical-techniques_-Introduction-pptx-25-320.jpg)
![Electrochemical
Techniques_Intro
5. Dehydration of glass membrane: If the
membrane has dried out, recondition it in water
for several hours before use.
6. Temperature: A pH meter should be calibrated
at the temperature at which pH measurements
will be made.
7. Na or alkaline error: when [H+
] is very low &
[Na+
] is high, the pH electrode responds to Na+
as if it were H+. So the apparent [H+
] is higher,
or apparent pH is lower.
8. Acid error: In strong acid, perhaps the glass
surface is saturated with H+
, so the apparent [H+
]
is lower & the apparent pH is higher,](https://image.slidesharecdn.com/electrochemicaltechniquesintroduction-250426192615-ddcf2265/85/Electrochemical-techniques_-Introduction-pptx-33-320.jpg)
![Electrochemical
Techniques_Intro
For Redox Reaction;
The indicator electrode appropriate is the
Platinum wire or foil.
The potential of the electrode is a function of
the ratio [Ox]/ [Red] forms of the ion. For
the general reaction.
The potential is given as;
E = Eo
+ 0.0592/n log. [Ox]/ [Red] at 25o
Where, Eo
is the standard oxidation
potential of the system](https://image.slidesharecdn.com/electrochemicaltechniquesintroduction-250426192615-ddcf2265/85/Electrochemical-techniques_-Introduction-pptx-40-320.jpg)
![For Precipitation Reactions;
The indicator electrode must readily come into equilibrium
with one of the ions in solution.
Thus Silver electrode is used for titration involving Halides
with Silver Nitrate.
The potential of the electrode is given as;
E = Eo + 0.0592/n log. [Mn+] at 25°C
Where, Mn+ is the ionic conc. present during titration and in
equilibrium with the slightly soluble precipitate](https://image.slidesharecdn.com/electrochemicaltechniquesintroduction-250426192615-ddcf2265/85/Electrochemical-techniques_-Introduction-pptx-41-320.jpg)
![Electrochemistry CE [Repaired]11111.pptx](https://cdn.slidesharecdn.com/ss_thumbnails/electrochemistrycerepaired-250930194342-440b1228-thumbnail.jpg?width=640&height=640&fit=bounds)
Electrochemical techniques_ Introduction.pptx
- 1. Electrochemical Techniques_Intro ELECTROCHEMICAL TECHNIQUES Kingsley A. Nti Dept.Industrial and Health Sciences Faculty of Applied Science Takoradi Technical University
- 2. Electrochemical Techniques_Intro OVERVIEW OF ELECTROCHEMICAL TECHNIQUES Thesimplest division of electrochemical techniques is between bulk techniques, in which a property of the solution in the electrochemical cell is measured, and interfacial techniques, in which the potential, charge, or current depends on the species present at the interface between an electrode and the solution in which it sits.
- 3. Electrochemical Techniques_Intro OVERVIEW CONT’D •The measurementof a solution’s conductivity, which is proportional to the total concentration of dissolved ions, is one example of a bulk electrochemical technique. •A determination of pH using a pH electrode is an example of an interfacial electrochemical technique.
- 4. Electrochemical Techniques_Intro ELECTROCHEMICAL TECHNIQUES •Analytical techniquesthat are based on the measurement of potential, charge, or current to determine an analyte’s concentration or to characterize an analyte’s chemical reactivity. •It involves the study of the movement of electrons in an oxidation–reduction reaction.
- 5. Electrochemical Techniques_Intro ELECTROCHEMISTRY IS THESTUDY OF PHENOMENA AT ELECTRODE-SOLUTION INTERFACES
- 6. Electrochemical Techniques_Intro Electrolysis / Power consumption Electrochemicalbattery / Power generation Chemical Reactions Electric Power
- 7. Electrochemical Techniques_Intro AN INTRODUCTION TOREDOX EQUILIBRIAAND ELECTRODE POTENTIALS The more negative the value, the stronger reducing agent the metal is. The more positive the value, the stronger oxidising agent the metal ion is.
- 8.
- 9. Electrochemical Techniques_Intro REDOX POTENTIALS FORNON- METAL AND OTHER SYSTEMS Chlorine gas is the strongest oxidising agent (E° = +1.36 V). A solution containing dichromate(VI) ions in acid is almost as strong an oxidising agent (E° = +1.33 V). Iron(III) ions are the weakest of the three new ones (E° = +0.77 V). None of these three are as strong an oxidising agent as Au3+ ions (E° = +1.50 V).
- 10. Electrochemical Techniques_Intro LOOKING AT THISFROM AN EQUILIBRIUM POINT OF VIEW Suppose you have a piece of magnesium in a beaker of water. There will be some tendency for the magnesium atoms to shed electrons and go into solution as magnesium ions. The electrons will be left behind on the magnesium.
- 11. Electrochemical Techniques_Intro A dynamic equilibriumwill be established when the rate at which ions are leaving the surface is exactly equal to the rate at which they are joining it again.
- 12. Electrochemical Techniques_Intro At that pointthere will be a constant negative charge on the magnesium, and a constant number of magnesium ions present in the solution around it.
- 13. Electrochemical Techniques_Intro Copper is lessreactive and so forms its ions less readily. Any ions which do break away are more likely to reclaim their electrons and stick back on to the metal again. An equilibrium position will still be reached, however there will be less charge on the metal, and fewer ions in solution.
- 14.
- 15. Electrochemical Techniques_Intro As the hydrogengas flows over the porous platinum, an equilibrium is set up between hydrogen molecules and hydrogen ions in solution. The reaction is catalysed by the platinum. standard hydrogen electrode
- 16. Electrochemical Techniques_Intro The standard hydrogenelectrode is attached to the electrode system under investigation - for example, a piece of magnesium in a solution containing magnesium ions.
- 17. Electrochemical Techniques_Intro Magnesium has amuch greater tendency to form its ions than hydrogen does. The position of the magnesium equilibrium will be well to the left of that of the hydrogen equilibrium. That means that there will be a much greater build-up of electrons on the piece of magnesium than on the platinum.
- 18. Electrochemical Techniques_Intro Explain the effectof replacing the magnesium half cell by a copper one
- 19. Electrochemical Techniques_Intro STANDARD ELECTRODE POTENTIALS The standard electrode potential of a metal / metal ion combination is the electro-motive force (emf) measured when that metal / metal ion electrode is coupled to a hydrogen electrode under standard conditions.
- 20. Electrochemical Techniques_Intro Ecell = Ecathode- Eanode In the copper case:
- 21. Electrochemical Techniques_Intro SCOPE OF ELECTROCHEMISTRY Investigation of chemical phenomena associated with a charge transfer reaction To assure electroneutrality, two half-reactions take place in opposite directions (oxidation/reduction) If the sum of free energy changes at both electrodes is negative electrical energy is released battery If it is positive, external electrical energy has to be supplied to oblige electrode reactions electrolysis
- 22. Electrochemical Techniques_Intro REACTIONS AND ELECTRODES Theoverall chemical reaction taking place in a cell is made up of two independent half-reactions, which describe the real chemical changes at the two electrodes. Most of the time one is interested in only one of these reactions, and the electrode at which it occurs is called the working (or indicator) electrode, coupled with an electrode that approaches an ideal nonpolarizable electrode of known potential, called the reference electrode.
- 23. The E.m.fof a metal immersed in a solution of its own ions may be obtained from the Nerst equation as follows; E = Eo + 0.0592/n * log C Where Eo is the standard potential of the metal, n is valency of the ions and c is the ionic concentration e.g Calculate the emf of the galvanic cell Fe| Fe2+ (0.200M)‖Ag+ (0.100M)|Ag e.g The emf of the galvanic cell Cu|Cu2+ (1.8x10-3 M)‖Ag+ (xM)|Ag is + 0.362. What is the [Ag+ ] in the silver half-cell?
- 24. THEORY OF EM F & PH DETERMINATIONS pH = -log [H+ ]= log 1/[H+ ] Similarly, pOH = -log [- OH] = log 1/ [- OH] Purified water ionizes in solution as follows; 2H2O H3O+ + - OH The above equation may be simplified as; H2O H+ + - OH Applying the law of Mass Action; [H+ ] [- OH] / [H2O] = Constant Where [H+],[-OH] and [H2O] represent their concentrations respectively
- 25. PH OF ASOLUTION Thus in pure water or neutral solution where [H+ ] = [- OH] =√Kw = 10-7 gram ions per litre (25o ) Also [H+ ] [- OH] = Kw = 10-14 Hence, pH + pOH = pKw = 14 One of the electrodes is referred to as the Indicator Electrode, and the other as Reference Electrode. The former must respond to pH changes whilst the other must give a constant potential. Each forms a half cell. Hydrogen electrode where n = 1, E = Eo + 0.0592 * log [H+ ] Alternatively, E = EHo - 0.0592 * pH But EHo is normally taken as zero, thus the above equation becomes; E = - 0.0592 * pH
- 26. Electrochemical Techniques_Intro REFERENCE ELECTRODE: Areference electrode is used in measuring the working electrode potential of an electrochemical cell. The reference electrode acts as a reference point for the redox couple. The internationally accepted primary reference is the standard hydrogen electrode (SHE) or normal hydrogen electrode (NHE), which is Pt/H2(a=1)/H+ (a=1,aqueous)
- 27. Electrochemical Techniques_Intro REFERENCE ELECTRODE By farthe most common reference is the saturated calomel electrode (SCE) and the Silver/Silver Chloride (Ag/AgCl) electrodes. SCE is Hg/Hg2Cl2/KCl (sat’d in water). Its potential is 0.242 V vs. NHE.
- 28. Electrochemical Techniques_Intro WORKING ELECTRODE •A fixedpotential difference is applied between the working electrode and the reference electrode. This potential drives the electrochemical reaction at the working electrode's surface. •The current produced from the electrochemical reaction at the working electrode is balanced by a current flowing in the opposite direction at the counter electrode.
- 29. Electrochemical Techniques_Intro MATERIALS OF WORKING ELECTRODE Carbon - based electrodes Metals such as platinum, gold, silver, nickel, mercury, gold-amalgam and a variety of alloys are now also commonly used as working electrode materials.
- 30. INDICATOR ELECTRODES These include; Hydrogen, Antimony, Glass and Specific Ion Electrodes. The use of each depends on factors like pH range, type of reaction involved or specific ion detection. GLASS ELECTRODE The potential of the glass electrode; E = K + 0.0592 (pH1 - pH2) at 25o C Where K is a constant, pH1 is that of the solution in the bulb and pH2 is that of the test solution. But pH1 is constant for a given glass electrode, thus E = K – 0.0592 pH2
- 31. Thin pH sensitiveglass bulb 0.1M HCl Ag/AgCl wire Thick walled glass tube ::::::::::::::::: ::::::::::::::::: :::::::::::::::::::. ::: ::: :::::::::::::::::
- 32. Electrochemical Techniques_Intro Errors in pHmeasurements 1. pH of buffer standards: only accurate to ±0.01 pH unit. 2. Junction potential: exists if µ of the analyte is different from that of the pH standards. To minimize this, use pH standards with the same µ. 3. Junction potential drift: exists when there is formation of AgCl (precipitation) or Ag (reduction) at the porous plug. To minimize this, recalibrate the electrode every 2 h. 4. Equilibration time: It takes time for an electrode to equilibrate with the analyte solution, esp. in a poorly buffered solution (pH varies greatly).
- 33. Electrochemical Techniques_Intro 5. Dehydration ofglass membrane: If the membrane has dried out, recondition it in water for several hours before use. 6. Temperature: A pH meter should be calibrated at the temperature at which pH measurements will be made. 7. Na or alkaline error: when [H+ ] is very low & [Na+ ] is high, the pH electrode responds to Na+ as if it were H+. So the apparent [H+ ] is higher, or apparent pH is lower. 8. Acid error: In strong acid, perhaps the glass surface is saturated with H+ , so the apparent [H+ ] is lower & the apparent pH is higher,
- 34. REJUVENATION OF GLASSELECTRODES Symptoms of a faulty Glass Electrode; Slow electrode response Undue sensitivity of the pH reading to physical movement of electrode. Failure of the electrode to check against a pair of buffer solutions. Inability to standardize in the range of the meter’s asymmetric potential.
- 35. REJUVENATION; A faultyglass electrode may be rejuvenated by momentarily immersing the bulb in 0.1M HCl or by cycling the bulb between acid and alkaline solution to reduce residual Sodium ion effects. Explain how residual sodium ion is accumulated When that fails, then the bulb must be immersed in 20% NH4F solution for 3 min or 10% HF for 15sec. After this it must be thoroughly rinsed in a stream of tap water then dipped momentarily in 5M HCl to remove Fluoride impurities. Finally the bulb is rinsed in purified water and the electrode stored in 0.1M HCl.
- 36. EMF VRS. PHMEASUREMENTS. At a given temperature, there is a linear relationship between pH of a solution and the Emf of a cell containing a reference and a suitable indicator electrode. Since E = K- 0.0592 pH at 25o C, Then ∆E/∆pH = -0.0592 Thus a meter calibrated in mV may be converted to pH units by dividing with 0.0592
- 37. POTENTIOMETRIC TITRATION: Visualindicators have been used to detect the end point of most titrimetric analysis but the method may be inaccurate in very dilute or coloured solutions. Under such conditions, potentiometric detection of end point yields very accurate results. The apparatus requires a potentiometer/pH meter with suitable indicator and reference electrodes, burette, a beaker and magnetic stirrer. Any reference electrode may be used provided it gives a constant potential. The indicator electrode must be appropriate for the type of titration ie, glass electrode for acid-base titration and Platinum electrode for redox titrations.
- 38. Typical Titration CurvesTypical Differential Titration Curves EMF/pH ' ' ' ' ' ' ' Volume of titrant ( ml) ' ' ' ' ' ' ' ' Equivalent point Volume of titrant ( ml) Equivalent point E V
- 39. For Neutralizationreactions; Any pH responsive indicator electrode may be used but a Glass electrode is preferable. The potential at the equivalence point is given by the equation; E = K – 0.0592 pH (25o C) Where, K is the asymmetric potential which depends on the type of electrode used.
- 40. Electrochemical Techniques_Intro For Redox Reaction; Theindicator electrode appropriate is the Platinum wire or foil. The potential of the electrode is a function of the ratio [Ox]/ [Red] forms of the ion. For the general reaction. The potential is given as; E = Eo + 0.0592/n log. [Ox]/ [Red] at 25o Where, Eo is the standard oxidation potential of the system
- 41. For Precipitation Reactions; Theindicator electrode must readily come into equilibrium with one of the ions in solution. Thus Silver electrode is used for titration involving Halides with Silver Nitrate. The potential of the electrode is given as; E = Eo + 0.0592/n log. [Mn+] at 25°C Where, Mn+ is the ionic conc. present during titration and in equilibrium with the slightly soluble precipitate
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