Experiment 2 – Conductivity and Net Ionic Equations

 

 

 

Purpose

 

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.

 

 

Questions to be answered in the introduction

 

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.     Define the following terms and tell what kind of bonding is expected in solutes that are

a.     Nonelectrolytes

b.     Strong electrolytes

4.     What is meant by dissociation and what is an example of a substances that dissociates?

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?

Experimental Procedure

 

Part 1 – Test and record the conductivity of each substance and solution listed below using the method demonstrated by your instructor  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.  Your data table may be similar to the one shown below

 

Substance

Conductivity

Classification

Principal species present in solution

No ions

Few ions

Many ions

Deionized water

 

 

 

 

 

Tap water

 

 

 

 

 

Methanol

CH3OH (l)

 

 

 

 

 

CH3OH (aq)

 

 

 

 

 

Glacial acetic acid CH3CO2H(l)

 

 

 

 

 

CH3CO2H(aq)

Add water slowly and record how the conductivity changes

 

 

 

 

 

Sucrose

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

 

 

 

 

 

Part 2 – Effect of solvent

Test the conductivity of the following solvents and mixtures.

 

Substance

Conductivity

Classification

Principal species present in solution

No ions

Few ions

Many ions

Xylene

C8H10

 

 

 

 

 

HCl in Xylene

 

 

 

 

 

 

Aqueous layer (after mixing with HCl in xylene)

 

 

 

 

 

 

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

 

Part 4 – Observing changes in conductivity for ionic reactions

 

Measure the conductivities of the acid and base alone, and then titrate the acid with base using a Pasteur Pipet and observe the changes in conductivity.  Add several drops of phenolphthalein to the acid before you begin the titration.  Continue adding several more drops of the base after the phenolphthalein turns pink.  Note how the conductivity changes as you add base to the acid.  Note the changes in conductivity as well as any other changes that may occur as you perform the titration.

 

  1. Perform the experiment with 0.1 M HCl with 0.1 M NaOH

 

Conductivities

0.1 M HCl                                                     

 

0.1 M NaOH                                                 

 

HCl + NaOH                                                 

 

 

Observations

 

 

 


  1. Perform the experiment with 0.1 M CH3COOH with 0.1 M NH3

 

Conductivities

0.1 M CH3COOH                                                     

 

0.1 M NH3                                                                 

 

CH3COOH + NH3                                                     

 

 

Observations

 

 

 

 

 

 

 

  1. Perform the experiment with 0.1 M H2SO4 with 0.1 M Ba(OH)2

Conductivities

0.1 M H2SO4                                                             

 

0.1 M Ba(OH)2                                                          

 

H2SO4 + Ba(OH)2                                                      

 

 

Observations

 

 

 

Analysis to be included in the results

 

Part 1

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

o   Write an equation to describe the formation of ions in aqueous acetic acid.

 

Part 3

o   Include the equations for the reactions of CaCO3 and zinc with acetic acid and hydrochloric acid.

 

Part 4

o   Write the conventional equation, total ionic equation, and net ionic equations for each of the processes observed

 

 

Questions and topics to be addressed in the discussion

 

Part 1

o   Why is tap water more conductive than distilled water?

o   Why does acetic acid form ions when it is dissolved in water but not when it is in the pure (glacial) form? 

o   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?

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

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

 

Part 2

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

 

Part 3

o   Discuss any 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

o   Explain any changes in conductivity you detected for each reaction.

Post-lab Questions (Attach this sheet to the back of your lab report)

 

  1. 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

Substance

Conductivity

Species

HF

Poor

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

K2SO4

Good

 

CH2O

Non conductive

 

HCHO2

Poor

 

 

  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.)

 

Substance

Species present

BaSO4

BaSO4(s)

CH3COOH

CH3COOH

NaCl

Na+,Cl-

Al(OH)3

 

Fe2(SO4)3

 

HBr

 

H3PO4

 

Mg(OH)2

 

NH4C2H3O2

 

C2H5OH

 

HgCl2

 

CuS

 

 

  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
    2. Ammonium hydroxide and acetic acid
    3. Nitric acid and magnesium acetate

      4.
    4. Ammonium chloride and sodium hydroxide
    5. Barium chloride and calcium nitrate
    6. Potassium hydrogen carbonate and sulfuric acid
    7. Aluminum hydroxide and nitric acid
    8. Ammonium hydroxide and sulfuric acid
    9. Magnesium nitrate and ammonium hydroxide