Chapter 19, worksheet 1 (ws19.1)
Review of Oxidation-reduction reactions (redox reactions) and balancing redox reactions
Redox reactions involve the transfer of electrons from one atom to another. The atom that loses
electrons is said to be oxidized. The atom that accepts electrons is reduced. If an atom is
oxidized another must be reduced (the electrons must go somewhere). The substance that is
oxidized is called the reducing agent. The substance that is reduced is called the oxidizing
agent. During a redox reaction, electrons are transferred from the reducing agent (which is
oxidized) to the oxidizing agent (which is reduced).
Memory Tricks
OIL RIG – Oxidation Is Loss; Reduction Is Gain
If an atom is reduced, its charge is reduced
(It has gained electrons)
Examples of redox reactions
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Combustion and corrosion were among the first redox reactions studied. These are reactions
with oxygen which explains why the term “oxidation” was coined.
Example of combustion: CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(g)
Example of corrosion:
4 Fe(s) + 3 O2(g) → 2 Fe2O3(s) (rust!)
Since these reactions don’t involve ions it is less obvious that they are redox reactions.
Strong acids can oxidize some metals. The metal appears to dissolve and hydrogen gas is
produced.
Molecular equation: Fe(s) + 2HCl(aq) → FeCl2(aq) + H2(g)
Net ionic equation: Fe(s) + 2H+(aq) → Fe2+(aq) + H2(g)
Conversely, hydrogen gas can reduce some metal ions (often
used to produce metals from their ores).
CuO(s) + H2(g) → Cu(s) + H2O(g)
Very reactive (easily oxidized) metals such as Li, Na, and K can be oxidized by water
2 K(s) + 2 H2O(l) → 2 KOH (aq) + H2(g)
Enough heat is produced to ignite the hydrogen gas (another redox
reaction!):
2 H2(g) + O2(g) → 2 H2O (g)
Notice that all of the above reactions are drawn with a single arrow. These reactions, like
many redox reactions, go to completion (large equilibrium constants).
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The Activity Series of Metals
Some metals lose electrons more easily than others. Table 5.5 below lists a number of metals
and hydrogen gas in order of their ease of oxidation. For example, the alkali metals and alkaline
earth metals are very easily oxidized so they are at or near the top of the table.
Strongest reducing agent
Weakest reducing agent
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Metals can be oxidized by Metal ions (salts)
The table below (not in your textbook) is a version of the activity series that includes the
oxidation “half-reactions” that occur when a metal or H2 is oxidized. Oxidation of a reducing
agent produces an oxidizing agent (do you see why?). Oxidation of a strong reducing agent
produces a weak oxidizing agent. Redox reactions will occur between the stronger oxidizing
agent and the stronger reducing agent (sound familiar?). Therefore:
A METAL WILL BE OXIDIZED BY ANY ION BELOW IT IN THE TABLE
Strongest reducing
agent
Weakest oxidizing
agent
Weakest reducing
agent
Strongest oxidizing
agent
Note: Oxidation “half-reactions” never occur in isolation. They must be coupled with a
reduction “half-reaction”. In order for a substance to lose electrons, something must
accept the electrons. The number of electrons lost must equal the number of electrons
gained. (You never end up with excess free electrons.)
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1. When zinc metal is placed in a solution of CrCl3, the following redox reaction takes place:
Molecular:
Zn(s) +
CrCl3 (aq) →
ZnCl2(aq) +
Cr(s)
(balance this equation)
Ionic:
Net ionic:
Refer to the net ionic equation and describe the transfer of electrons that occurs during the
reaction.
___________ is oxidized and ___________ is reduced.
___________ is the oxidizing agent and ____________ is the reducing agent.
Write the two balanced half reactions.
Oxidation:
Reduction:
Many redox reactions (like neutralization reactions) go to completion because one direction
occurs much more readily than the other. Why doesn’t the reverse of the reaction above occur to
a measurable extent?
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2. Table 4.5 above can be used to predict whether or not a redox reaction will occur. Write the
molecular and net ionic equations for any reaction that will occur.
a. Will zinc metal react with nickle (II) chloride or will nickel react with zinc chloride?
b. Will sodium chloride react with copper metal or will sodium react with copper (II)
chloride?
c. Name a metal that will be oxidized by a strong acid.
d. List 5 metals from table 4.5 that will NOT be oxidized by strong acids. Choose one of
these metals and find an oxidizing agent in Table 4.5 that WILL oxidize it.
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Most metals are not found as free elements in nature. Instead, they are found combined with
other elements into compounds (salts, minerals, ores). Why? (The process of purifying metals
from ores is called “metallurgy”. Metallurgy exploits the principles of redox chemistry and
electrochemistry to reduce metal ions to metals. Read sections 20.1, 20.2, 20.3, and 20.7 in your
textbook. Also read the learning objectives for chapter 20. I will not lecture on metallurgy.)
3. Which metals can be found in nature as pure elements rather than combined with other
elements? Why?
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Balancing Oxidation-Reduction Reactions Using the Half-Reaction Method
Oxidation-reduction reactions that take place in aqueous solution, particularly in acidic or basic
solutions, can be very complicated and difficult to balance by inspection. The following steps
will assist you in balancing these kinds of reactions.
Half-Reaction Method
If a redox reaction takes place in acidic solution,
1.
2.
3.
4.
5.
Divide the equation into two incomplete half-reactions, one for oxidation and the
other for reduction.
Balance each half-reaction.
a.
First, balance the elements other than H and O.
b.
Next, balance the O atoms by adding H2O.
c.
Then, balance the H atoms by adding H+.
d.
Finally, balance the charges by adding e-.
Multiply each half-reaction by an integer so that the number of electrons lost in
one half-reaction equals the number gained in the other.
Add the two half-reactions and simplify where possible.
Check the equation.
If a redox reaction takes place in neutral or basic solution, there are not enough protons around
to participate in the reaction. Thus, the final equation must contain OH- instead of H+. Rather
than developing a new set of steps,
1.
2.
Balance the reaction as if it were in acid, as described above.
Neutralize the H+ by adding the requisite number of OH- ions to both sides of the
equation.
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1. Use the half-reaction method to balance this reaction that takes place in acidic solution.
a.
Cu(s) + NO3- (aq) → Cu2+ (aq) + NO2 (g)
b.
The UNBALANCED reaction that occurs in the Breathalyzer test is:
CH3CH2OH (g) + Cr2O7 2- (aq) → CH3COOH (aq) + Cr3+
(ethanol)
c.
Mn2+ (aq) + NaBiO3 (s) → Bi3+ (aq) + MnO4- (aq)
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2. Balance the following equations which take place in neutral or basic solution.
a.
CN- (aq) + MnO4- (aq) → CNO- (aq) + MnO2 (s)
b.
Cr(OH)3 (s) + ClO- (aq) → CrO42- (aq) + Cl2 (aq)
c.
NO2- (aq) + Al (s) → NH3 (aq) + Al(OH)4- (aq)
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Rules for Assigning Oxidation Numbers
1. All atoms in an element = 0
2. All monatomic ions = their charge
3. Oxygen = -2 except in peroxide ion, O22- where it is -1 (e.g. H2O2)
4. Hydrogen = +1 when bonded to non-metals (e.g. HCl, H2O);
-1 when bonded to metals (e.g. LiH)
5. Fluorine = -1
6. All other halogens = -1 except when bonded to oxygen
(e.g. ClO3-) in which case they have positive oxidation numbers.
7. All other oxidation numbers are selected so as to make the sum of
the oxidation numbers equal to the charge on the molecule or ion.
If an atom is oxidized (loses electrons) its oxidation number increases.
If an atom is reduced (gains electrons) its oxidation number decreases.
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