Chapter Four: Types of Chemical Reactions and Solution Stoichiometry
I.
Water the Common Solvent (Like Dissolves Like Rule)
a. Water able to dissolve many different substances.
b. H2O molecules with “bent” shape (105 degree angle) created by H – O – H.
c. O – H bonds in water molecule = covalent bonds formed by electron sharing between the oxygen
and hydrogen atoms.
d. Water is a polar molecule
i. Electrons are not shared equally because oxygen atom has a greater attraction for
electrons.
ii. Electrons thus are usually located closer to the oxygen atom versus the hydrogen atoms.
iii. This polarity gives water its great ability to dissolve compounds (ionic and nonionic).
e. Hydration – Addition of water to a molecule.
f. Ions
i. When ionic substances (salts) dissolve in water, they break up into individual cations and
anions. [NH4NO3 (s)  NH4+(aq) + NO3-(aq)]
g. Solubility varies for different substances depending on the relative attractions of the ions for each
other and the attractions of the ions for water molecules.
II. Nature of Aqueous Solutions
a. Solute = substance. Solvent = liquid.
b. Electrical conductivity = ability to conduct an electric current.
i. Solutions containing strong electrolytes conduct electricity efficiently.
ii. Solutions containing weak electrolytes conduct electricity poorly.
iii. Solutions that permit no electric current to flow contain nonelectrolytes.
iv. The extent to which a solution can conduct an electric current depends directly on the
number of ions present.
c. Strong electrolytes = substance completely ionized when dissolved in water.
i. Acids, when dissolved in water, behave as strong electrolytes.
1. Acid is a substance that produces protons when dissolved in water.
2. Strong Acids = completely dissociates into ions.
ii. Strong bases also completely dissociate.
d. Weak electrolytes = exhibit a small degree of ionization in water.
i. Weak acids and weak bases
1. Only dissociates to a slight extent.
e. Nonelectrolytes = substances that dissolve in water but do not produce any ions.
III.
Composition of Solutions
a. To perform stoichiometric calculations we need to know (1) the nature of the reaction and (2) the
amounts of chemicals present in the solutions.
b. Concentrations of Solution
i. Molarity (M) = moles of solute per volume of solution in liters.
c. Standard solution = a solution whose concentration is accurately known.
d. Dilution – Adding water do achieve the molarity desired for a particular solution.
i. Moles of solute after dilution = moles of solute before dilution.
ii. Only volume changers.
iii. Common instruments used in dilution include a piper and volumetric flask.
iv. Equation: M1V1 = M2V2
IV.
Types of Chemical Reactions
a. Types of Solution Reactions
i. Precipitation reactions
ii. Acid-Base reactions
iii. Oxidation-reduction reactions
V.
Precipitation Reactions
a. A solid forms from the reaction. Solid is called a precipitate.
b. To predict products, think of various possibilities and the likelihood of each possibility.
i. When Ions forma solid compound, the compound must have a zero net charge.
VI.
VII.
VIII.
ii. Most ionic materials contain only two types of ions: one type of cation and one type of
anion.
iii. When ionic substances dissolve in water, the ions separate.
iv. Ions interchange in a reaction.
Describing Reactions in Solution
a. Molecular Equation = shows reactants and products online.
i. Gives the overall reaction stoichiometry but not necessarily the actual forms of the
reactants and products in solution.
b. Complete ionic equation = shows dissociated reactants and products.
i. Represents as ions all reactants and products that are strong electrolytes.
c. Net ionic equation = only those solution components directly involved in the reaction are shown
i. Includes only those solution components undergoing a change. Spectators ions not
included.
d. Spectator ions = ions that do not participate directly in the reaction.
Stoichiometry of Precipitation Reactions
a. Always write down the species that are actually present in the solution.
b. Obtain the moles of reactants we must use the volume of the solution and its molarity.
c. Stoichiometry for reactions in solution
i. Identify the species present in the combined solution, and determine what reaction occurs.
ii. Write the balanced net ionic equation for the reaction.
iii. Calculate the moles of reactants.
iv. Determine which reactant is limiting.
v. Calculate the moles of product or products, as required.
vi. Convert to grams or other units, as required.
Acid-Base Reactions (Bronsted Lowry)
a. Acid = proton donor.
b. Base = proton acceptor.
c. When strong acids and strong bases are mixed, the net ionic equation is as follows:
H+ (aq) + OH- (aq)  H2O (l)
d. The hydroxide ion is such a strong base that it can be assumed to react completely with any weak
acid encountered.
e. Performing calculations for Acid-Base reactions
i. List the species present in the combined solution before any reaction occurs, and decide
what reaction will occur.
ii. Write the balanced net ionic equation for this reaction.
iii. Calculate the moles of reactants. For reactions in solution, use the volumes of the
original solutions and their molarities.
iv. Determine the limiting reactant where appropriate.
v. Calculate the moles of the required reactant or product.
vi. Convert to grams or volume as required.
f. Acid-base reaction is often called a neutralization reaction.
g. Acid-Base titrations
i. Volumetric Analysis – A technique for determining the amount of a certain substance by
doing titration.
ii. Titration – Delivery (from a buret) of a measured volume of a solution of known
concentration (the titrant) into a solution containing the substance being analyzed (the
analyte).
iii. Equivalence point – The point in the titration where enough titrant has been added to
react exactly with the analyte.
iv. Indicator – a substance added at the beginning of the titration that changes color at the
equivalence point.
v. Endpoint – The point where the indicator actually changes color.
vi. For a titration to be successful…
1. The exact reaction between titrant and analyte must be known
2. The stoichiometric point must be marked accurately.
3.
The volume of titrant required to reach the stoichiometric point must be known
accurately.
IX.
Oxidation-Reduction Reactions
a. Reactions in which one or more electrons are transferred or called redox reactions.
b. Oxidation States/Numbers = the imaginary charges the atoms would have if the shared electrons
were divided equally between identical atoms bonded to each other with the greatest attraction for
electrons.
c. Oxidation states of ions are equal to the ion charges.
Table 4.2 Rules for Assigning Oxidation States
1. The oxidation state of an atom in an element is 0.
2. The oxidation state of a monatomic ion is the same as its charge.
3. In its compounds, fluorine is always assigned an oxidation state of -1.
4. Oxygen is usually assigned an oxidation state of -2 in its covalent compounds. Exceptions to this rule include
peroxides, where each oxygen is assigned an oxidation state of -1.and OF2 in which oxygen is assigned a +2
oxidation state.
5. In its covalent compounds with nonmetals, hydrogen is assigned an oxidation state of +1.
6. The sum of the oxidation states must be zero for an electrically neutral compound. For an ion, the sum of the
oxidation states must equal the charge of the ion.
d. Non-integer oxidation states are rare but do occur because of the arbitrary way that electrons are
divided up.
e. Oxidation – increase in oxidation state (a loss of electrons).
f. Reduction – decrease in oxidation state (a gain of electrons)
g. Oxidizing agent – Electron acceptor
h. Reducing Agent – Electron donor
i. And always remember LEO THE LION SAYS GER
i. Lose electrons in oxidation and Gain electrons in reduction.
ii. When the oxidizing or reducing agent is named, the whole compound is specified, not just
the element that undergoes the change in oxidation state.
X.
Balancing Oxidation-Reduction Equations
a. It is useful to split up reactions to balance redox reactions, one involving oxidation and the other
involving reduction.
b. Steps in balancing redox equations in an acidic solution.
i. Write separate equations for the oxidation and reduction half-reactions.
ii. For each half reaction balance all the elements except hydrogen and oxygen, then balance
oxygen using H2O, then balance hydrogen using H +, then balance the charge using
electrons.
iii. If necessary, multiply one or both balanced half reactions by an integer to equalize the
number of electrons transferred in the two half=reactions.
iv. Add the half-reactions, and cancel identical species.
v. Check that the elements and charges are balanced
c. Steps in balancing redox equations in a basic solution.
i. Use the half-reaction method specified for acidic solutions to obtain the final balanced
equation as if H+ ions were present.
ii. To both sides of the equation obtained above, add a number of OH- ions that is equal to
the number of H+ ions.
iii. Form H2O on the side containing both H+ and OH- ions, and eliminate the number of
H2O molecules that appear on both sides of the equation.
iv. Check that the elements and charges are balanced.
Multiple Choice Questions
I.
II.
III.
IV.
V.
You have a solution of table salt in water. What happens to the salt concentration as the solution boils?
a. The concentration decreases
b. The concentration increases
c. The concentration stays the same
d. None of the above
The oxidation number of Xe in the XeOF4 is
a. +6
b. -6
c. +4
d. -4
The oxidation number of Ni in NiO2 is
a. +1
b. +2
c. +4
d. -4
When a strong base and strong acid react, the solution will most likely be
a. Strongly acidic
b. Basic
c. Neutral
d. Slightly acidic
What volume of .100 M HCL will be needed to react completely with 50.00 mL of 0.200 M NaOH?
a. .100 mL
b. 10.0 mL
c. 100. mL
d. none of the above
Free-Response
A chemist analyzed the effluent (the released waste material) from an industrial process known to produce the
compounds carbon tetrachloride (CCl4) and benzoic acid (HC7H5O2), a weak acid that has one acidic hydrogen atom
per molecule. A sample of this effluent weighing 0.3518g was shaken with water, and the resulting aqueous
solution required 10.59 mL of 0.1546 M NaOH for neutralization. Calculate the mass percent of HC7H5O2 in the
original sample.
I.
Write down the equation for this reaction.
II.
Calculate the moles of NaOH
III.
Calculate the moles of HC7H5O2
IV.
Calculate grams of HC7H5O2
V.
Calculate the mass percent of HC7H5O2 in the original sample