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ELECTROLYTE CONDUCTANCE

ELECTROLYTE CONDUCTANCE. CMT552 ELECTROCHEMISTRY AND CORROSION SCIENCE. What is electrolyte?. Any substance that produce ions when dissolved in a solvent (usually water) is an electrolyte. It is the electrically conductive solution that must be present for corrosion to occur.

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ELECTROLYTE CONDUCTANCE

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  1. ELECTROLYTE CONDUCTANCE CMT552 ELECTROCHEMISTRY AND CORROSION SCIENCE

  2. What is electrolyte? • Any substance that produce ions when dissolved in a solvent (usually water) is an electrolyte. • It is the electrically conductive solution that must be present for corrosion to occur. Types of electrolytes • Strong electrolyte • Weak electrolyte • Non-electrolyte

  3. Strong Electrolytes • Strong electrolytes are substances that only exist as ions in solution. • They completely dissociate to their ions when dissolved in solution. • Ionic compounds are typically strong electrolytes. • Strong acids, strong bases and salts are strong electrolytes.

  4. They conduct electricity when molten or in aqueous solution. • Example: Hydrochloric acid, Sodium chloride

  5. Weak Electrolytes • A weak electrolyte only partially dissociates in solution and produces relatively few ions (exist in water as a mixture of individual ions and incontact molecules). • Polar covalent compounds are typically weak electrolytes. • Weak acids and weak bases are weak electrolytes.

  6. They conduct electricity weakly. • Example: Acetic acid, ammonia

  7. Non-electrolytes • A non-electrolyte does not dissociate at all (present entirely as intact molecules) in solution and therefore does not produce any ions. • Non-electrolytes are typically polar covalent substances that do dissolve in water as molecules instead of ions. • They do not conduct electricity at all. • Example: Sugar

  8. Acids • Are molecular compounds which ionize or turn into ions in water. • The properties of acids were due to the presence of hydrogen ions, H+. • All acids are soluble in water • Some acids are strong electrolytes and some are weak electrolytes. • No acids are non-electrolytes.

  9. Bases • Can be molecular compounds or ionic compounds. • Some bases are soluble and some are not. • The soluble bases ionize or dissociate into ions in water. • The properties of bases were due to the presence of hydroxide ions, OH-. • All of the ionic bases which are soluble are also strong electrolytes.

  10. Salts • Are ionic compounds which are not acids or bases. • In other words, the cation is not hydrogen and the anion is not hydroxide. • Some salts are soluble in water and some are not. • All of the salts which are soluble are also strong electrolytes.

  11. Electricals Terms

  12. Other Symbols and Terms

  13. Electrolytic conductance • Electrolytic conductance occurs when a voltage is applied to the electrode dipped into an electrolyte solution, ions of the electrolyte move and electric current flows through the electrolytic solution. • This power of the electrolyte to conduct electricity is known as conductance or conductivity. • Electrolytic solution also obey Ohm’s Law just like metallic conductor.

  14. Ohm’s Law: It states that the current flowing through a conductor is directly proportional to the potential difference across it: where, V = applied potential (V) I = current measured (A) R = solution resistance () between the two electrodes V=IR

  15. Solution Resistance (R) • The increase the [ions] presence in the solution, the lower the solution resistance, R, will be. • A strong electrolyte like KCl is dissolve in water, the no. of ions per unit volume increase and the solution resistance, R, is lowered, thus increasing the current measured for a particular applied potential.

  16. Thus, current can be related to the [ions] in a particular solution. • However, • the distance between the electrodes, • the surface area of the electrodes and • the identity of the ions also affect the solution resistance, R.

  17. Solution Conductance (G) • The reciprocal of solution resistance (1/R) is called Conductance, G. • Conductance is expressed as Siemens (S) or ohm-1 (-1) or mho. • Where, A = surface area of each electrode l = distance btwn electrode  = conductivity

  18. Values of conductivity,, increased with T and concentration. • The conductivity of a solution of water is highly dependent on its concentration of dissolved salts and sometimes other chemical species which tend to ionize in the solution.

  19. Electrical conductivity of water samples is used as an indicator of how salt free or impurity free the sample is; the purer the water, the lower the conductivity.

  20. Molar Conductivity () • Defined by: Example 1: Molar conductivity of 0.005 M KCl is 144 Scm2 mol-1. Calculate its electrolytic conductivity in SI units (Sm-1). *(Hint: 1m2 = 104 cm2; mol/L or mol/dm3convert to mol m-3). Units: Sm-1 mol m-3

  21. Measurement of Conductivity • The conductivity of a solution is measured in a cell called conductance cell or conductivity cell. • Since l and A are difficult to measure, the usual procedure is to treat as a cell constant, K cell

  22. Therefore, Example 2 In a certain conductivity cell, the resistance of a 0.01 M KCl solution is 150 . The known molar conductivity of the solution is 141.27 -1 cm2 mol-1. Calculate the cell constant (Kcell). *(Kcell unit is cm-1)

  23. Exercise 1 Using the same conductance cell as in example 2, a student measured the resistance of a 0.10 M NaCl solution to be 19.9 . Calculate the experimental value of the molar conductivity of this solution.

  24. Use the same value of kcell

  25. Ex • In order to determine the molar conductivity of a 0.05 M solution of AgNO3, you need to measure the solution resistance in a conductivity cell and found that R = 75.8 . Then, in the same cell, a 0.02 M KCl solution had a resistance of 157.9 . Given that the accepted molar conductivity of the KCl solution is 0.013834 -1 m2 mol-1, calculate the molar conductivity of the AgNO3 solution.

  26. Variation of Molar Conductivity with Concentration • Molar conductivity () of electrolytes increases with dilution. • The variation is different for strong and weak electrolytes.

  27. Strong electrolytes • Fully ionized in solution •  increases slowly with dilution and there is a tendency for  to approach a certain limiting value when the concentration approaches zero(i.e. When dilution is infinite).

  28. The  when the concentration approaches zero (infinite dilution) is called molar conductivity at finite dilution or limiting molar conductivity (°).  = ° when C → 0 (at infinite dilution) • For strong electrolytes molar conductivity increase slowly with dilution and can be represented by: DEBYE HUCKEL ONSAGER equation

  29.  = Molar conductivity at a given concentration ° = Molar conductivity at infinite dilution  = constant

  30. From the graph, it has been noted that the variation of molar conductivity () with concentration (C) is small so that the plot can be extrapolated to zero concentration. • The intercept is equal to (°) and the slope is -.

  31. Weak electrolytes • Not fully ionized in solution • In weak electrolyte like acetic acid they have low degree of dissociation as compared to strong electrolyte.

  32. However, the variation of molar conductivity ()with concentration (C) is very large and we can’t obtain molar conductivity at infinite dilution (°) by extrapolation of  versus C plots.

  33. Explanation for the variation of Molar Conductivity with concentration • Conductance behaviour of strong electrolyte: • No increase in the no. of the ions with the dilution ( completely ionized in the solution at all concentration). • In concentrated solution: • strong inter-ionic forces • Molar conductivity is low

  34. In dilute solution: • Inter-ionic forces low • Molar conductivity increases with dilution When concentration very low, inter-ionic interaction becomes almost negligible and molar conductance approaches the limiting value, °.

  35. Conductance behaviour of weak electrolyte: • The no. of ions produced in solution depends upon the degree of dissociation with dilution. • Higher the degree of dissociation, larger is the molar conductance. • With increase in dilution • Degree of dissociation increases as a result molar conductivity increases. • At infinite dilution, the electrolyte is completely dissociated so that the degree of dissociation become one.

  36.  = ° (at C → 0) Thus, if  = Molar conductivity at a given concentration ° = Limiting molar conductivity or molar conductivity at infinite dilution Then, degree of dissociation

  37. Ostwald Dilution Law & Dissociation Constant of Weak Electrolyte • Consider an aqueous solution of a weak binary electrolyte, AB, of concentration C mol dm-3 and degree of dissociation of . • At equilibrium:

  38. Therefore, the dissociation constant can be expressed as: Ostwald Dilution Law

  39. However, for weak electrolyte;  is very small. • Hence, (1- )  1 • Therefore, • Since H+ = C •  [H+ ] = C =

  40. KOHLRAUSCH’S LAW • At infinite dilution the ions act completely independently, and the ° obeys a rule of additivity: where AX, AY, BX and BY are strong electrolytes.

  41. ° for a weak electrolyte can be deduced from ° values obtained from strong electrolytes. • For example, consider CH3COOH denoted as HAc’: where HX, Mac and MX are strong electrolytes.

  42. Table 1: Limiting Molar Conductivity, °, of some strong electrolytes

  43. Exercise 2 Calculate ° for a weak electrolyte NH4OH from the ° values for these strong electrolytes: NH4Cl: 149.7; NaCl: 126.5 and NaOH: 248.10

  44. Kohlrausch also stated at infinite dilution when the dissociation complete, • each ion makes a definite contribution towards molar conductance of the electrolyte irrespective of the nature of the other ion with which it is associated.

  45. It means that the molar conductivity at infinite dilution for a given salt can be expressed as the sum of the individual contributions from the ions of the electrolyte.

  46. where v+ and v-: stoichiometric coefficients for the cation and anion in the electrolyte. °+ and °-: ionic conductance of individual ions (cation and anion)

  47. Example 3 For NH4OH electrolyte: v+ = 1 and v- = 1 Since 1NH4+ ion present for each OH- ion present in solution. Example 4 For K4Fe(CN)6 electrolyte: v+ = 4 and v- = 1 Since there are 4K+ ions present for each Fe(CN)4-6 ion present in solution.

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