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Electrolytes And Nonelectrolytes Case Solution & Answer

Electrolytes And Nonelectrolytes Case Study Solution

When electricity is conducted through a solid it is generally due to the passage of electrons.  When it is conducted through a liquid, however, it is generally due to the passage of positive and negative ions.  Such ions are present when some substances are in the liquid state, and these substances are properly called electrolytes.  Other substances provide these ions when mixed with highly polar materials such as water and these are also properly classified as electrolytes.  All other substances that do not provide these ionic species in solution are classified as non-electrolytes.

IONIC COMPOUNDS

Ionic compounds are composed of positive and negative ions attracting one another and are organized into extremely stable repeating patterns called lattices.  In the solid state these ions have very little mobility and the compound does not electricity.  However, when such compounds are simply melted the attractions “holding” the ions together are weakened to such an extent that the mobility of the ions is greatly increased.  Therefore, all ionic compounds will conduct electricity in the pure liquid state and are electrolytes.  When an ionic compound dissolves, the ions are separated (dissociated) by the polar solvent and again their mobility is greatly increased.  Therefore, all aqueous solutions of ionic compounds will conduct electricity.  It can be stated that all ionic compounds (soluble or insoluble in aqueous media) are strong electrolytes.

COVALENT COMPOUNDS

In covalent compounds atoms are bonded together through the sharing of electrons.  There are no fully formed ions present.  Therefore, no pure covalent compound can conduct electricity.  However, when some covalent compounds are mixed with water, the mixture conducts electricity.  From our definition of electrolyte, it can be seen that some covalent compounds are indeed electrolytes while others are not.

In some covalent compounds ions are partially (although not certainly fully) formed.  These are called polar covalent bonds.

When covalent compounds which are highly polar are mixed with water they are ionized.  In this process, the positive part of the water molecule “pulls” at the negative end of the polar molecule and the negative part of the water molecule “pulls” at the positive end of the polar molecule.  These partially formed ions are pulled apart as fully formed ions, and the resulting mixture will conduct electricity.  It is due to the process of ionization that an aqueous solution of HCl conducts electricity while pure liquid HCl does not.

When some covalent compounds are mixed with water, a very large percentage of the molecules are ionized.  These are considered to be strong electrolytes.  Other covalent compounds (those which are non-polar or slightly polar) do not ionized to any measurable extent when mixed with water, and are therefore considered to be non-electrolytes.  There are many compounds which fall between these two extremes both in polarity and in the extent to which ionization takes place when they are mixed with water.  Ammonia (NH3) is an example of this.  When mixed with water, most ammonia molecules remain unchanged and only a small percentage are ionized.  Compounds of this type are considered to be weak or moderate electrolytes.

In summary:

  1. All ionic compounds are strong electrolytes.
  2. Covalent compounds may be strong, weak or non-electrolytes, depending upon the extent of ionization when mixed with water.

IONIC EQUATIONS

As previously explained, many compounds when mixed with water split apart into ions.  Many reactions carried out in water are, therefore, reactions involving ions as well as molecules.  There are three common methods used in writing equations for such reactions:

  1. The standard or molecular equation
  2. The ionic equation
  3. The net ionic

The following example shows each of these forms for the reaction between nitric acid and potassium hydroxide:

  1. HNO3(aq) + KOH(aq) –> KNO3(aq) + H2O(l)
  2. H+1(aq) + NO3-1(aq)  +  K+1(aq) +  OH-1(aq)  –>  K+1(aq)  +  NO3-1(aq)  +  H2O(l)
  3. H+1(aq) + OH-1(aq)  –>  H2O(l)

In order to write the ionic equation, it is necessary to know which of the substances involved are primarily in the form of separated ions when they are mixed with water.  Such substances are written as separated ions (HNO3, KOH and KNO3 in the above example) while all others (H2O from above) are written as neutral and unseparated.

The following are the solubility rules for common ionic solids, use these in making the decision. If two rules appear to contradict each other, the preceding rule takes precedence.

  1. Salts containing Group I elements (Li+, Na+, K+, Cs+, Rb+) are soluble. There are few exceptions to this rule. Salts containing the ammonium ion (NH4+) are also soluble.
  2. Salts containing nitrate ion (NO3) are generally soluble.
  3. Salts containing Cl, Br , or I are generally soluble. Important exceptions to this rule are halide salts of Ag+, Pb2+, and (Hg2)2+. Thus, AgCl, PbBr2, and Hg2Cl2 are insoluble.
  4. Most silver salts are insoluble. AgNO3 and Ag(C2H3O2) are common soluble salts of silver; virtually all others are insoluble.
  5. Most sulfate salts are soluble. Important exceptions to this rule include CaSO4, BaSO4, PbSO4, Ag2SO4 and SrSO4.
  6. Most hydroxide salts are only slightly soluble. Hydroxide salts of Group I elements are soluble. Hydroxide salts of Group II elements (Ca, Sr, and Ba) are slightly soluble. Hydroxide salts of transition metals and Al3+ are insoluble. Thus, Fe(OH)3, Al(OH)3, Co(OH)2 are not soluble.
  7. Most sulfides of transition metals are highly insoluble, including CdS, FeS, ZnS, and Ag2 Arsenic, antimony, bismuth, and lead sulfides are also insoluble.
  8. Carbonates are frequently insoluble. Group II carbonates (CaCO3, SrCO3, and BaCO3) are insoluble, as are FeCO3 and PbCO3.
  9. Chromates are frequently insoluble. Examples include PbCrO4 and BaCrO4.
  10. Phosphates such as Ca3(PO4)2 and Ag3PO4 are frequently insoluble.
  11. Fluorides such as BaF2, MgF2, and PbF2 are frequently insoluble.
  12. Strong acids (HCl, HBr, HI, HClO4, HClO3, HNO3, H2SO4) should be written as separated ions. All other acids and nitrogen bases are deemed as weak and should be written as molecules.

The following are some examples of writing ionic and net ionic equations using the above rules:

  1. (NH4)2SO4(aq) + 2KOH(aq) –> 2NH4OH(aq) + K2SO4(aq)

Ionic form:  2NH4+(aq) + SO42-(aq) + 2K+(aq) + 2OH(aq) –> 2NH4OH (aq) + 2K+(aq) + SO42-(aq)

Net ionic equation: NH4+(aq) + OH(aq) –> NH4OH(aq)

  1. 2HCl(aq) + BaCO3(s) –> BaCl2(aq) + H2CO3(aq)

Ionic form:  2H+(aq) + Cl(aq) + BaCO3(s) –> Ba2+(aq) + 2Cl(aq) + H2CO3(aq)

Net ionic equation:  2H+(aq) + BaCO3(s) –> Ba2+(aq) + H2CO3(aq)…………….

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