Experiment 08

Andrew T. HedrickChem 32 Section S

11/1/02

Experiment 8 – Solution Conductivity: Covalent & Ionic Solutes

Purpose: To relate bonding structure of substances to their behavior in solution, investigate electrical conductance of various aqueous solutions, and follow the conductance of a solution during the course of a chemical reaction.

Experimental Plan: The instructions for this plan were based on The Thompson General Chemistry 2 Lab Manual, pages 91100. Changes included using 5mL of deionized water as the solvent in Part IA.

Data and Observations:

Part IA:drops KI 0 1 2 3 4 5 6conductivity (in mA)

.0721 1.31 2.40 3.55 4.65 5.86 6.93

Part IB:Solution Conductivit

yelectrolyte? Solution Conductivit

yelectrolyte?

HCl 30.5mA strong Urea 70.6μA noSucrose 2.54mA weak NaOOCCH

3

6.38mA weak

NaOH 16.5mA strong C6H8O7 7.6mA weakCuSO4 10.4mA strong CH3OH 90.0μA noNH3(aq) 1.73mA weak H3PO4 16.2mA strong

Part IC:Solute ConductivityGlacial acetic acid 1.0μA.01M acetic acid .99mA.01M hydrochloric acid 31.5mA


11/1/02Part II: Please see attached sheet.

Calculations: No calculations were performed in this lab.

Results: In Part IA, conductivity increased as KI was added. In Part IB, most of the substances conducted a fair amount of electricity. Only two conducted so little that they were considered nonelectrolytes. The remaining eight compounds were evenly split between being strong and weak electrolytes. In Part IC, the less concentrated NH3 turned out to conduct more. In Part II the conductivity fell as additional H2SO4 was added, but eventually it bottomed out and began to rise again.

Conclusion: Across all of these experiments, the determining factor to increased conductivity was the amount of free ions in solution. This is shown directly in Part IA, as additional KI was added the amount of ions in solution grew. This, in turn, made the conductivity rise. In Part IB, the strong electrolytes were the solutes that completely ionized in solution (such as strong acids like HCl and strong bases like NaOH). Weak electrolytes were solutes that only partially ionized in solution (such sucrose and weak bases like NH3). Nonelectrolytes did not donate any free ions to solution and thus made extremely weak conductance occur between our electrodes. In Part IC, the concentrated NH3 had nearly no conductance. This is because there was no water present and no ions had formed to carry electrical current. Aqueous (.01M) NH3 on the other hand conducted electricity much better due to the ions being formed in the reversible reaction with water. HCl was a very good conductor due to it completely ionizing in water. In Part II, H2SO4’s conductivity dropped as it reacted with Ba(OH)2. This is because the chemical reaction was forming water and BaSO4. Once all of the BaSO4 had reacted free ions became present and conductivity rose. Generally, ionic bonds, the strong acids and bases all make very good electrolytes. Most molecular compounds make poor electrolytes and some are nonelectrolytes. As learned in Part II, you can use conductance to determine when a reaction has completed. Possible sources of error for this experiment include the multimeter, the conductance breakout box, and residual solution present on the electrodes and glasswear being used.


11/1/02

3. Suppose that Part II of the experiment were carried out with aqueous NaOH (sodium hydroxide) instead of Ba(OH)2 to form soluble Na2SO4 along with H2O as products.

(a) Write the overall balanced equation

2NaOH(aq) + H2SO4(aq) � Na2SO4(s) + 2H2O(l)

(b) Write the net ionic equation

2Na+(aq)+ 2OH(aq) + 2H+(aq) + SO42(aq) � Na2SO4(s) + 2H2O(l)

(c) How would the graph of conductivity versus volume of aqueous NaOH added differ from the graph you constructed using Ba(OH)2?

They would be similar in that conductivity would go down until the NaOH was used up to make products. The conductivity would rise again after ions are formed. The only difference is that NaOH is that the graph would not be as steep as NaOH is not as strong of an electrolyte as Ba(OH)2.

4. The electrical conductivity of deionized water is commonly used as an index of its purity. Explain.

“Pure” water should not conduct any electricity, as it would have no free ions to move charges around. However, since water dissolves an extremely large number of substances, it almost always conducts electricity. Thus, the more “pure” the water, the lower its electrical conductivity. In this way, conductivity can be a gauge to test water’s purity.


11/1/02

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Experiment 08