Experiment 7

Conductivity and Net Ionic Equations






To use electrical conductivity as a way of determining the number of free ions present in a substance and to use this information to draw conclusions regarding the type of bonding present in the substance.


Background information


Types of bonding


There are three basic types of bonds.  They are ionic, covalent, and polar covalent which is a hybrid of an ionic and a covalent bond.  Ionic bonds are formed between elements with very different electronegativities.  Generally ionic compounds form between metals and non-metals and are identified by the transfer of an electron from the metal to the non-metal to form charged ions which are held together by electrostatic interactions known as ionic bonds.  Covalent bonds form between elements with similar electronegativities.  Generally covalent bonds that form between two non-metals are characterized by the sharing of electron pairs between the atoms.  Polar bonds are formed between two elements with different electronegativities, but which still share the electrons albeit unevenly.


Electrolytes vs. non-electrolytes


Substances may be classed by their electrical conductivity.  Substances which do not conduct electricity are called non-electrolytes.  Examples of nonelectrolytes are covalent molecules such as sucrose or table sugar (C12H22O11) and acetone (CH3COCH3).  These substances are non-electrolytes because they are not composed of ions and cannot conduct an electric current.  Ionic compounds in their crystalline form are also considered to be non-electrolytes because even thought they are composed of ions, the ions are not able to move freely through the crystal and therefore are not able to conduct electricity.  Substances which can conduct electricity are called electrolytes.  When ionic compounds are melted or dissolved in water forming mobile ions they are able to conduct electricity.  Polar covalent compounds such as acids and bases will sometimes disassociate or break apart in aqueous solution to form ions as well.  Examples of these are shown below:


Text Box: Notice the difference in conductivity between strong and weak acids and bases.  What does this tell you about the relative degree of dissociation?  What are the principal species present in the solution?  How do we write these substances in ionic equations?HCl  + H2O  ˆ  H3O+  +  Cl-

HNO3  +  H2O  ˆ  H3O+  +  NO3-

H2SO4  +  H2O  ˆ  2 H3O+  +  SO4-2


CH3CO2H  +  H2O  ˆ  H3O+  +  CH3CO2-

NH3  +  H2O  ˆ  NH4+  +  OH-



Weak electrolytes are substances that are able to conduct electricity, but conduct poorly.  Examples of weak electrolytes are molecular substances that dissociate to a small extent such as weak acids and bases, and ionic compounds that have limited water solubility.


Writing ionic equations


When you write ionic equations, you need to show the principal species present in the solutions.  By testing the conductivity of a variety of solutions you can determine whether the principal species are ions or undissociated or undissolved particles.  For each of the substances tested for electrical conductivity, determine the principal species present in solution.


Experimental Procedure


This experiment will be done as an instructor demonstration.  A variety of substances will be tested for electrical conductivity.  Fill in the attached data sheet with your observations and analysis of the demonstration.



Prelab Questions


1.     Name the three principal bond types.


2.     What is the essential characteristic of a solution that is

a.     A nonconductor


b.     A good conductor


c.     A poor conductor



3.     In general, what kind of bonding is expected in solutes that are

a.     Nonelectrolytes


b.     Strong electrolytes



4.     When strong acids, strong bases, and soluble salts are dissolved in water they become _______electrolytes because they are completely _________ in aqueous solutions.


5.     What is the meaning of the term ÒhydrationÓ when used to describe what happens to an electrolyte, which is dissolved in water?



6.     What can be inferred about the degree of ionization or dissociation of a substance that is a weak electrolyte?

7.     Listed below are several substances and their conductivities when dissolved in water. Based on this information, write the formulas for all of the individual species present.  Star(*) those present in low concentration and omit water from your list







HF(aq) H+1*, F-1*





Non conductive






Name _______________________  Section_________


Experiment 7 – Data sheet


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Part 1 – Record the conductivity of each substance and solution tested.  Then, after noting the range of conductivities measured, classify each as having essentially no ions, a few ions, or many ions.  For each substance record also the principal species present in the sample.





Principal species present in solution

Minor species present in solution

No ions

Few ions

Many ions

Distilled water







Tap water








CH3OH (l)







CH3OH (aq)







Glacial acetic acid CH3CO2H(l)















C12H22O11 (s)







C12H22O11 (aq)







NaCl (s)







NaCl (aq)







KClO3 (s)







KClO3 molten







0.1 M HgCl2







0.1 M HCl







0.1 M NaOH







0.1 M CH3CO2H







0.1M NH3







0.1 M NaCl








1.     Write the equation for the reaction that forms the few (but important) ions in pure water.







2.     Why is tap water more conductive than distilled water?







3.     Why does acetic acid form ions when it is dissolved in water but not when it is in the pure (glacial) form?  Write an equation to describe the formation of ions in aqueous acetic acid.







4.     The models of solid NaCl describe it as consisting of Na+ cations and Cl- anions.  If this is the case, why is NaCl(s) not an electrolyte?







5.     Explain the distinctly different behavior in the conductivity of NaCl(s) and NaCl(aq).







6.     Explain the different behavior in the conductivity of KClO3(s) and KClO3(l).





Part 2 – Effect of solvent





Principal species present in solution

No ions

Few ions

Many ions








HCl in Xylene







Aqueous layer (after mixing with HCl in toluene)

















Why is the conductivity of HCl different in the two solvents?  What causes this difference?








Part 3 – Correlating Chemical and Conductivity Behavior


  1. Compare the rates of reaction of CaCO3(s) with 6 M acetic acid and 6 M hydrochloric acid. 







  1. Write the equation for the reaction of CaCO3 with acetic acid.








  1. Write the equation for the reaction of CaCO3 with hydrochloric acid.





  1. Compare the rates of reaction of Zn(s) with 6 M acetic acid and 6 M hydrochloric acid. 





  1. Write the equation for the reaction of zinc with acetic acid





  1. Write the equation for the reaction of zinc with hydrochloric acid.





  1. Is there a correlation between the rates of the chemical reactions of 6M acetic acid and 6 M hydrochloric acid and the conductivities observed for acetic acid and hydrochloric acid in part 1?






Part 4 – Observing changes in conductivity for ionic reactions


  1. 0.01 M HCl with 0.01 M NaOH


0.01 M HCl                                                   


0.01 M NaOH                                                           


HCl + NaOH                                                 


Conventional equation




Total ionic equation




Net ionic equation



Explain any changes in conductivity you detected.



  1. 0.1 M CH3COOH with 0.1 M NH3


0.1 M CH3COOH                                                     


0.1 M NH3                                                                 


CH3COOH + NH3                                                     


Conventional equation




Total ionic equation




Net ionic equation




Explain any changes in conductivity you detected.




  1. 0.1 M H2SO4 with 0.1 M Ba(OH)2


0.1 M H2SO4                                                             


0.1 M Ba(OH)2                                                          


H2SO4 + Ba(OH)2                                                      


Conventional equation




Total ionic equation




Net ionic equation




Explain any changes in conductivity you detected.



Post-lab Questions


  1. Recognition of Ionic or Molecular Species Present.  For each substance, write the formula(s) of the principal species (molecules or ions) present in major amount in the aqueous solution if the substance is soluble; if it is only slightly soluble, use the molecular or ionic formula, followed by (s).  (1st three are done as examples.)



Species present


























  1. Separate solutions of 0.01 M Ba(C2H3O2)2 and 0.01 M H2SO4 are tested for electrical conductivity.  Equal volumes of these solutions are then mixed and the conductivity tested again.  Predict the result and justify your answer with an appropriate equation.














  1. The electrodes of the conductivity apparatus of this experiment are immersed in 10 mL of 0.01 M Ba(OH)2, and the conductivity is noted.  A student blows his breath through a glass tube dipping into the solution.  Can the student Òblow out the lightÓ?  Write equations for any reactions, and justify your conclusions.  (HINT:  CO2  +  H2O  ˆ  H2CO3)





















  1. Interpretation of Reactions by Ionic Type Equations.  Aqueous solutions of the following substances or their mixtures with water if they are only slightly soluble are mixed.  Write first the conventional equation, second the total ionic equation, and last the net ionic equation.  If you predict no appreciable reaction, indicate this and state why.
    1. Magnesium chloride and sodium carbonate











    1. Ammonia and acetic acid









    1. Nitric acid and magnesium acetate













    1. Ammonium chloride and sodium hydroxide













    1. Barium chloride and calcium nitrate













    1. Potassium hydrogen carbonate and sulfuric acid

    2. Aluminum hydroxide and nitric acid











    1. Ammonia and sulfuric acid











    1. Magnesium nitrate and ammonia