The following study guides are linked to Anne McKenna's website (http://employees.csbsju.edu/amckenna/CH123), with permission by the author.

Reaction Types

Most of the reactions you will see in lab (in fact, most chemical reactions) occur in solution.  For typical reactions, this means dissolving the reactants in water, mixing the two solutions and observing whether or not a reaction occurs.  In this study guide we will first learn about solutions, then we will discuss types of reactions.  Finally, using what we know about solutions and reaction types, we will predict products of reactions.

Solutions

Imagine that you are in a lab and you are asked to place a solid substance into water.  The solid substance, (the substance being dissolved) is called the solute.  Water, the dissolving medium, is called the solvent.  When the solute is placed in the solvent, it may either be soluble or insoluble.  A solute which is soluble in the solvent dissolves, forming a solution.  An insoluble solute does not dissolve in the solvent and remains a solid.  An example of a soluble solute in water would be sodium chloride.  When solid sodium chloride is placed in water it dissolves, forming a sodium chloride solution.  An example of an insoluble solute would be chalk dust (calcium carbonate), which does not form a solution with water.

How can we predict which substances will be soluble in water?  Chemists have developed a set of solubility rules which allow us to predict which substances are water-soluble.  These are given below:

Solubility Rules:

Soluble

1.  All compounds of alkali metals (Li⁺,Na⁺,K⁺) are soluble.

2.  All compounds containing the ammonium ion, NH₄⁺, are soluble.

3.  All compounds containing the following anions are soluble:

NO₃^-

ClO₃^-

ClO₄^-

CH₃COO^-

4.  Most compounds containing Cl^-, Br^-, I^- are soluble

            Exceptions

            AgCl, PbCl₂ and Hg₂Cl₂ are insoluble.

5.  Most compounds containing SO₄^2- are soluble

            Exceptions

            CaSO₄, SrSO₄, BaSO₄  and PbSO₄ are insoluble

Insoluble

1.  Metal oxides are generally insoluble

            Exceptions

            Alkali metal oxides (Li₂O, Na₂O, K₂O) are soluble

            CaO, SrO, BaO are soluble

2.  Compounds containing hydroxide ion (OH^-) are insoluble

            Exceptions

            Alkali metal hydroxides (LiOH, NaOH, KOH) are soluble

            Ca(OH)₂, Sr(OH)₂, Ba(OH)₂ are soluble

3.  Compounds containing CO₃^2-, PO₄^3-, S^2-, SO₃^2- are insoluble

            Exceptions

            Alkali metals compounds containing these anions are soluble

            NH₄⁺ compounds containing these anions are soluble

So, using the solubility rules above, we would predict that NaOH would be soluble (Rule 1), while Zn(OH)₂ would be insoluble (Rule 2).

1.  Using the solubility rules above, determine if the following substances are soluble or insoluble in water.

            a.  KNO₃         ______________

            b.  Na₂CO₃      ______________

            c.  Ni(OH)₂      ______________

            d.  CuO            ______________

Strong, Weak and Non-Electrolytes

Solubility rules only give us part of the picture of a solution.  We know that sodium chloride is soluble in water, but what does a water solution of sodium chloride “look like” at the molecular level?  To describe solutions, we first need to decide what happens to a solute as it dissolves. 

We can categorize solutions by how well they conduct an electric current.  In order to conduct a current, ions must be present in the solution.  Solutes described as strong electrolytes conduct an electric current well, those described as weak electrolytes exhibit some conductivity and solutes which are nonelectrolytes do not conduct an electric current.  Therefore, strong electrolytes dissolve to form a high concentration of ions, weak electrolytes dissolve to form a low concentration of ions, and nonelectrolytes dissolve forming no ions.

We can use the solubility rules, along with some other information, to classify solutes as strong electrolytes, weak electrolytes and nonelectrolytes.

Strong electrolytes:  All soluble ionic compounds (Use Solubility Rules)

All strong acids (HCl, HBr, HI, HNO₃, H₂SO₄, HClO₄)

Weak Electrolytes:    All weak acids (CH₃CO₂H, HF, H₃PO₄, etc.).

                                    All weak bases (NH₃, other nitrogen-containing bases)

Nonelectrolytes:        All substance dissolving as molecular species (pure water, alcohols, sucrose, other organic compounds)

So, a substance like NaCl (a soluble ionic compound) is a strong electrolyte in water.  This means that a NaCl solution contains water molecules, sodium (Na⁺) ions and chloride (Cl^-) ions.  There are no NaCl “molecules” in the solution.  This is the case with all strong electrolytes:  to be correct they should be written as ions in solution.

Weak electrolytes form some ions in solution, but are mostly found in their molecular forms; so, NH₃ in water consists primarily of NH₃ molecules (with a few NH₄⁺ and OH^- ions present).  Substances such as CH₃OH (methyl alcohol) which are nonelectrolytes may be soluble in water, but they dissolve as molecules.

2.  The following substances are added to water.  Describe what makes up the solution.

             a.  NaOH

            b.  HI

            c.  HF

            d.  CaCO₃

            e.  CH₃CH₂OH (ethyl alcohol)

 Aqueous Reactions

 When two aqueous solutions are mixed, does a chemical reaction occur?  If it does, what will be the product(s) of the reaction?  To answer these questions, we need to recognize the pattern of chemical reactions.

In order for a chemical reaction to occur, the solutes must recombine to form a substance which is removed from the reaction.  What does this mean?  For electrolyte solutions, this means that the ions must “change partners” to form something that is isolated from the reaction mixture.  How can something be isolated from the reaction mixture?  Typically we think of the formation of a solid (precipitation), a covalently bound substance such as water (acid-base), or a gas (gas-forming).  In all three of these cases, the product is removed from the ions in solution.

Precipitation Reactions

In a precipitation reaction, a solid product is formed from two electrolyte solutions.  Let’s work through an example, then see if we can use the pattern to predict the products of another reaction.

Example:  A NaI solution is mixed with a Pb(NO₃)₂ solution.

Step 1:  Determine makeup of reacting solutions.

            NaI is a soluble ionic compound and exists in solution as Na⁺(aq) and I^-(aq)

Pb(NO₃)₂ is a soluble ionic compound and exists in solution as Pb²⁺(aq) and NO₃^-(aq)

 Step 2:  Look at recombining ions (pairing + ions with – ions) and write correct formulas for ion combinations.

                        NaI solution                  Pb(NO₃)₂ solution

                        Na⁺                              Pb²⁺

                        I^-                                  NO₃^-

                         Possible combinations:  NaNO₃

 PbI₂                                                                            

Step 3:  Determine if either ion combination forms a precipitate by checking solubility rules.

NaNO₃ is soluble, so cannot be a precipitate

PbI₂ is insoluble, so is a precipitate

Step 4:  Write equation for reaction, giving formulas for reactants and products.

Indicate soluble reactants and products by (aq) and insoluble ones by (s).

First (unbalanced) equation:

NaI(aq) + Pb(NO₃)₂(aq)  à NaNO₃(aq) + PbI₂(s)

Balanced equation:

2NaI(aq) + Pb(NO₃)₂(aq)  à 2NaNO₃(aq) + PbI₂(s)

Step 5:  Write total ionic and net ionic equations. 

A total ionic equation writes all strong electrolytes as ions.  Remember that a precipitate is not a strong electrolyte, so it does not form ions in solution.

            2[Na⁺(aq) + I^-(aq)] + Pb²⁺(aq) + 2NO₃^-(aq) à 2[Na⁺(aq) + NO₃^-(aq)] + PbI₂(s)

A net ionic equation eliminates all of the ions not participating in the reaction (the spectator ions).  In the reaction above, Na⁺ has the same form as a reactant and a product, so it would be a spectator ion.  One other ion, the NO₃^- ion, also is a spectator.  So, the equation written without the spectator ions would be:

            2I^-(aq) + Pb²⁺(aq) à PbI₂(s)

The net ionic equation gives us the “real picture” in the reaction.  What is actually happening is the formation of solid lead (II) iodide.  The other ions in the solution do not participate in the reaction.

3.  Use the steps given above to predict the product of the following reaction and to write balanced, total ionic and net ionic equations.

            K₂CO₃ + CaCl₂  à

Step 1:  Determine makeup of reacting solutions.

Step 2:  Look at recombining ions (pairing + ions with – ions) and write correct formulas for ion combinations.

Step 3:  Determine if either ion combination forms a precipitate by checking solubility rules.

Step 4:  Write equation for reaction, giving formulas for reactants and products.

Step 5:  Write total ionic and net ionic equations.

Acid/Base Reactions

Acid/base reactions follow a similar pattern as precipitation reactions.  The major difference between the two types of reactions involves the formation of water instead of (or in addition to) a precipitate.  First, let’s re-define acids and bases.

Acids:  donate H⁺ to a base.  Most formulas of acids write the H⁺  being lost first, such as HCl and HNO₃.

            Bases:  accept H⁺ from an acid.  Many bases contain the hydroxide ion (OH‑) or produce

hydroxide ion when added to water.  Examples of bases include metal hydroxides and ammonia (NH₃) and other nitrogen-containing compounds.

Let’s look at some aqueous solutions of acids and bases and determine the makeup of the solution.

An aqueous solution of HCl is a strong electrolyte (because HCl is a strong acid).  Therefore, we would predict that an HCl solution consists of H⁺ and Cl^- ions in water.  This is a logical prediction based on the behavior of other strong electrolytes.  However, the H⁺ ion in solution, because it is so small, associates itself with one or more water molecules.  However, although it would be more correct to write H₃O⁺, for simplicity’s sake we will use H⁺.

An aqueous solution of NaOH is a strong electrolyte because NaOH is a soluble metal hydroxide.  Therefore, a water solution of NaOH contains Na⁺ ions and OH^- ions.

Aqueous solutions of weak acids and weak bases are weak electrolytes.  The primary form of the acid or base in these solutions is the molecular (undissociated) form.  So, a solution of H₃PO₄, a weak acid, contains primarily H₃PO₄ molecules.

Now, we’re ready to start acid/base reactions.  We will use the similar steps as those used for precipitation reactions.

Example:  A solution of HCl is mixed with a solution of KOH.

Step 1:  Determine makeup of reacting solutions.

            HCl is a strong acid and exists in solution as H⁺(aq) and Cl^-(aq).

KOH is a strong base (soluble metal hydroxide) and exists in solution as K⁺(aq) and OH^-(aq)

 Step 2:  Using H⁺ transfer between acid and base determine products of reaction.

             H⁺ will be transferred to OH^-, the base. 

         H⁺ + Cl^- + K⁺ + OH^-

            H⁺ transferred to OH^-

            K⁺ and C^l- will not react (check solubility rules to see if precipitate forms)

Step 3:  Write equation for reaction, giving formulas for reactants and products.

            First (unbalanced) equation

            HCl(aq) + KOH(aq) à H₂O(l) and KCl(aq)

            Note the use of (l) by water.  This means that water is a pure liquid.

            Balanced equation:

HCl(aq) + KOH(aq) à H₂O(l) and KCl(aq)

Step 4:  Write total ionic and net ionic equations.

Total Ionic: 

H⁺(aq) + Cl^-(aq) + K⁺(aq) + OH^-(aq) à H₂O(l) + K⁺(aq) + Cl^-(aq)

Note that we don’t separate water into ions.  This is because pure water is a nonelectrolyte.

Net Ionic: