Reactions Between Ions in Aqueous Solution

Posted on February 25, 2009

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>A solution is a homogeneous mixture in which the two or more components mix freely
The solvent is taken as the component present in the largest amount
A solute is any substance dissolved in the solvent
For example, the percentage concentration is the number of grams of solute per 100 g of solution
The relative amounts of solute and solvent are often given without specifying the actual quantities
There is usually a limit to the amount of solute that can dissolve in a given amount of solvent
For example, 36.0 g NaCl is able to dissolve in 100 g of water at 20°C
A solution is said to be saturated when no more solute can be dissolved at the current temperature
The solubility of a solute is the number of grams of solute that can dissolve in 100 grams of solvent at a given temperature
Solubilities of some common substances
Solubility usually increases with temperature
Supersaturated solutions contain more solute than required for saturation at a given temperature
They can be formed, for example, by careful cooling of saturated solutions
Supersaturated solutions are unstable and often result in the formation of a precipitate
A precipitate is the solid substance that separates from solution
Precipitates can also form from reactions
Reactions that produce a precipitate are called precipitation reactions
Many ionic compounds dissolve in water
Solutes that produce ions in solution are called electrolytes because their solutions can conduct electricity
An ionic compounds dissociates as it dissolves in water
Most solutions of molecular compounds do not conduct electricity and are called nonelectrolytes
The dissociation of ionic compounds may be described with chemical equations
The hydrated ions, with the symbol (aq), have been written separately
Since physical states are often omitted, you might encounter the equation as:
Ionic compounds often react when their aqueous solutions combine
This reaction may be represented with a molecular, ionic, or net ionic equation:
Molecular:
Ionic:
Net Ionic:
The most compact notation is the net ionic equation which eliminates all the non-reacting spectator ions from the equation
Criteria for balanced ionic and net ionic equations:
Material balance – the same number of each type of atom on each side of the arrow
Electrical balance – the net electrical charge on the left side of the arrow must equal the net electrical charge on the right side of the arrow
In the reaction of Pb(NO3)2 with KI the cations and anions changed partners
This is an example of a metathesis or double replacement reaction
Solubility rules allows the prediction of when a precipitation reaction will occur
For many ionic compounds the solubility rules correctly predict whether the ionic compound is soluble or insoluble
Solubility rules for ionic compounds in water:
Soluble Compounds
Insoluble compounds
A knowledge of these rules will allow you to predict a large number of precipitation reactions
Acids and bases are another important class of compounds
Acids and bases affect the color of certain natural dye substances
They are called acid-base indicators because they indicate the presence of acids or bases with their color
The first comprehensive theory of acids, bases, and electrical conductivity appeared in 1884 in the Ph.D. thesis of Savante Arrhenius
He proposed that acids form hydrogen ions and bases released hydroxide ions in solution
The characteristic reaction between acids and bases is neutralization
HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l)
In general, the reaction of an acid and a base produces water and a salt
We can state the Arrhenius definition of acids and bases in updated form
In general, acids are molecular compounds that react with water to produce ions
This is called ionization:
It is common to encounter the hydrogen ion (H+) instead of the hydronium ion
The previous ionization is also written as
Monoprotic acids are capable of furnishing only one hydrogen ion per molecule
Acids that can furnish more than one hydrogen ion per molecule are called polyprotic acids
Some nonmetal oxides react with water to produce acids
They are called acidic anhydrides (anhydride means without water)
Soluble metal oxides are base anhydrides
Examples include:
Ammonia gas ionizes in water producing hydroxide ions
It is an example of a molecular base
Many molecules that contain nitrogen can act as a base
Binary compounds of many nonmetals and hydrogen are acidic
In water solution these are referred to as binary acids
They are named by adding the prefix hydro- and the suffix –ic to the stem of the nonmetal name, followed by the word acid
Acids that contain hydrogen, oxygen, plus another element are called oxoacids
They are named according to the number of oxygen atoms in the molecule and do not take the prefix hydro-
When there are two oxoacids, the one with the larger number of oxygens takes the suffix –ic and the one with the fewer oxygen atoms takes the suffix –ous
The halogen can occur with up to four different oxoacids
The oxoacid with the most oxygens has the prefix per- the one with the least has the prefix hypo-
Anions are produced when oxoacids are neutralized
There is a simple relationship between the name of the polyatomic ion and the parent acid
–ic acids give –ate ions
-ous acids give –ite ions
In naming polyatomic anions, the prefixes per- and hypo- carry over from the parent acid
Polyprotic acids can be neutralized
An acidic salt contains an anion that is capable of furnishing additional hydrogen ions
The number of hydrogens that can still be neutralized is also indicated
Naming bases is much less complicated
Ionic compounds containing metal ions are named like any other ionic compound
Molecular bases are specified by giving the name of the molecule
Acids and bases can be classified as strong or weak and so as strong or weak electrolytes
Strong acids are strong electrolytes
The most common strong acids are:
Strong bases are the soluble metal hydroxides
These include:
Most acids are not completely ionized in water

They are classified as weak electrolytes
Weak acids and bases are in dynamic equilibrium in solution
Consider the case of acetic acid:
Neutralization of a strong acid with strong base gives a salt and water:
This net ionic equation applies only to strong acids and bases
The neutralization of a weak acid with a strong base involves a strong and weak electrolyte
Consider the neutralization of acetic acid with NaOH:
Note that in ionic equations the formulas of weak electrolytes are written in “molecular” form
The situation is similar when a strong acid reacts with a strong base
For ammonia and HCl the net ionic equation is:
Note that water only appears as a product if the hydronium ion is used
Both strong and weak acids react with insoluble hydroxides and oxides
The driving force is the formation of water
Magnesium hydroxide has a low solubility in water, but reacts with strong acid
The net ionic equation is:
Magnesium hydroxide is written as a solid because it is insoluble
A number of metal oxides also dissolve in acids
For example, iron(III) oxide reacts with hydrochloric acid:
Some reactions with acids or bases produce a gas
The reactions are driven to completion because the gas escapes and is unavailable for back reaction
(CO2 and SO2 are produced by the decomposition of H2CO3 and H2SO3, respectfully)
Solutions are characterized by their concentration
The molar concentration or molarity (M) is defined as
The molarity of a solution gives an equivalence relation between the moles of solute and volume of solution
Solutions provide a convenient way to combine reactants in many chemical reactions
Example: How many grams of AgNO3 are needed to prepare 250 mL of 0.0125 M AgNO3 solution?
ANALYSIS: Find moles, then mass of solute.
SOLUTION:
Solutions of high concentration can be diluted to make solutions of lower concentration
Conservation of solute mass requires:
Where dil labels the diluted and concd the concentrated solution
Stoichiometry problems often require working with volumes and molarity
Example: How many mL of 0.124 M NaOH are required to react completely with 15.4 mL of 0.108 M H2SO4?
2 NaOH + H2SO4  Na2SO4 + 2H2O
ANALYSIS: Use the mole-to-mole ratio to convert.
SOLUTION:
Limiting reagent problems are also common
Example: How many moles of BaSO4 will form if 20.0 mL of 0.600 M BaCl2 is mixed with 30.0 mL of 0.500 M MgSO4?
BaCl2 + MgSO4  BaSO4 + MgCl2
ANALYSIS: This is a limiting reagent problem.
SOLUTION:
Titration is a technique used to make quantitative measurements of the amounts of solutions
The end-point is often determined visually
Paths for working stoichiometry problems may be summarized with a flowchart:

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