Investigating Reactions in Solution – Part II: Redox reactions
Chem. 319
Ms. Hobbie
In oxidation-reduction reactions, commonly called redox reactions, electrons are transferred from one
substance to another.
When an atom gains electrons, it is said to be REDUCED and the process is referred to as reduction.
When an atom loses electrons, it is said to be OXIDIZED and the process is referred to as oxidation.
Single Replacement Reactions. Single replacement reactions occur when one reactant is in its
elemental form and, during the course of the reaction, accepts or receives electrons to become an ion.
At the same time, an ion in the reacting solution does the opposite, loses or gains electrons to become
the neutral element. These reactions typically take the form…
A + BC → AC + B
Single replacement reactions fall into the following categories.
1. Metal Replacement
 A more active metal (i.e. more easily oxidized
metal) will react with the metal ion of a less active
metal. In these reactions, the metal ion acts as the
oxidizing agent.
2. Hydrogen Replacement
a. Reactions with Water
 The group 1 and group 2 metals react with water
to form a metal hydroxide and hydrogen gas.
b. Reactions with Acids
 All the metals above hydrogen in the activity series
can react with hydrogen ion in solution. In these
reactions, H+ ion (really the H3O+ ion) acts as an
oxidizing agent and is reduced to H2 gas.
 Note that hydrogen ion cannot act as an oxidizing
agent for the metals below hydrogen in the
activity series.
 Nitric acid is fairly unique because the NO 3- ion is a
better oxidizing agent than the H+ ion, and so
metals reacting with nitric acid are being oxidized
by the nitrate ion.
3. Halogen Replacement
 The halogens react as oxidizing agents (they are reduced to halide ions). An activity series for
the halogens is the same as their order on the periodic table, with fluorine being the strongest
oxidizing agent.
F2 > Cl2 > Br2 > I2
Some other Easily Recognizable Redox Reactions
The following categories of reactions are easily recognizable. See the patterns?

Decomposition reactions occur when one compound breaks down into two or more new substances.
This usually requires an input of energy.
BC → B + C.

Synthesis (or combination) reactions occur when two substances combine to form a new substance.
A + C → AC.

Combustion reactions occur when an organic substance reacts with oxygen to form (most
commonly) carbon dioxide and water.
CxHyO? + O2(g) → CO2 + H2O.
Common Oxidizing and Reducing Agents.
Some of the most useful redox reactions do not fall into any specific category, but are best recognized
by the presence of a solution that contains either a good oxidizing agent, or a good reducing agent, as
defined below:
Strong Oxidizing Agents:
The Permanganate Ion: MnO4in strongly acidic solutions MnO4- (aq, purple) → Mn2+ (aq, clear)
in neutral or basic solutions MnO4- (aq, purple) → MnO2 (s, black)
The Chromate/Dichromate Ions: CrO42-, Cr2O72in acidic solution, Cr2O72- (aq, orange) → Cr3+ (aq, green)
in basic solution, CrO42- (aq, yellow) → Cr(OH)3 (s, white)
Strong Reducing Agents:
The Sulfite, Bisulfite, and Thiosulfate ions: SO32-, HSO3-, S2O32All of these partially-oxidized sulfur compounds become fully-oxidized to the aqueous sulfate ion in
most redox reactions. This is colorless in aqueous solution:
SO32- → SO42HSO3- → SO42H2SO3 → SO42S2O32- → SO42The Oxalate ion (C2O42–)
This ion is readily oxidized to CO2, which will be easily recognized by its signature “fizzing” as the
reaction progresses.
Experiment Combine the following substances. Use 1 mL of the aqueous solutions unless otherwise
specified. In the space provided, write down simple observations.
Single Replacement Reactions Just do four of the first six… Zn and Cu
1. zinc metal (strip) + copper (II) nitrate →
(perform this and the next five reactions in different wells of a “reaction plate.”)
2. lead metal (strip) + copper (II) nitrate →
3. copper metal (strip) + zinc nitrate →
4. lead metal (strip) + zinc nitrate →
5. zinc metal (strip) + lead(II) nitrate →
6. copper metal (strip) + lead(II) nitrate →
7. sodium metal + water → (do this in an evaporating dish and add a drop of phenolphthalein)
8. zinc metal + 6 M hydrochloric acid → (in a large test tube)
9. copper metal + 6 M hydrochloric acid → (in a large test tube)
10. chlorine in water + potassium bromide → (in a small test tube)
11. iodine in water + potassium bromide → (in a small test tube)
Looking at common oxidizing and reducing agents
Perform all of the remaining reactions in small test tubes, as indicated. Again, use 1 mL of the aqueous
solutions unless otherwise specified.
12. 2 drops potassium permanganate + hydrochloric acid (6 M) → (waft for product)
13. 2 drops potassium dichromate (acidified: 3 drops 6 M H2SO4) + iron (II) sulfate →
Seeing the dichromate change to green is fun, but with this it just turns pale yellow… maybe beef up the
concentration of the iron(II) sulfate? If they really follow the directions it is greener.
14. 2 drops potassium permanganate (acidified: 3 drops 6 M H2SO4) + iron(III) sulfate →
15. Iron(III) sulfate (acidified: 3 drops of 6 M H2SO4) + a pinch of solid sodium sulfite →
Yuck – SO2 gas. And this turns red… some further reaction with the iron.
Tried dichromate acidified with solid sulfite… some SO2 but fantastic color change. Do this in the HOOD
incase you get some SO2 as a byproduct.
16. 1 drop potassium permanganate + 10 drops water (acidified: 3 drops 6 M H2SO4)
+ sodium oxalate →
NOTE: No change? Place test tube into a hot water bath and observe.
Reaction Equations:
1. For each of the first 11 reactions, write a net ionic equation if a reaction was observed. You will find
the first page of this handout quite helpful with these.
 What is oxidized?
 What is reduced?
 Who is spectating?
Reaction #7 is challenging… let your observations help you. Hint: no spectators.
2. For any of the first 11 reactions that did not occur, EXPLAIN why there was no reaction.
3. For reactions #9-13, write an unbalanced “skeleton” equation if a reaction was observed. Do this by
thinking through the following questions:
 What are the products of this reaction? Keep in mind that something must be oxidized and
something else must be reduced. Allow your observations to support your pondering. You will
find p.2 of the introduction to this lab immensely helpful!
 What is the specific “species” in solution that was oxidized? … reduced? (If it was aqueous,
was it an ion or a molecule?)
 Are there any spectators here that can be left out of the skeletal equation?
4. For any of these final reactions (#12-16) that did not occur, EXPLAIN why there was no reaction.
Follow-up Questions for discussion:
5. Choose one reaction from the second set (#12-16) and determine the balanced net ionic equation
using the half-reaction method. (AOHC)
6. Use balanced half reactions to determine the net ionic equation for the following reactions:
a. Magnesium reacts with hydrochloric acid. The magnesium disappears and bubbles are
formed.
b. Aqueous solutions of sodium iodide and excess nitric acid react to form aqueous sodium
iodate (NaIO3) and bubbles of nitrogen monoxide gas. (It is helpful here to identify any
spectator ions… and leave them out of the half reactions.)
7. Looking at page 232 of your textbook answer the following questions:
Would solid chromium dissolve in a copper(II) nitrate solution?
Would solid copper dissolve in a chromium(III) nitrate solution?
Defend your answer in a complete sentence (or two) as well as a balanced equation.
8. Explain how you know that the synthesis of carbon dioxide from its elements is a redox reaction.
9. The decomposition of hydrogen peroxide, H2O2(aq) produces water and oxygen gas. It is classified
as a special case referred to as a disproportionation reaction. Look up the definition for this term and
explain why this reaction falls into this category.