Chemistry Lab Report Format
Follow this suggested format for your lab report that is due 17
August.
INTRODUCTION (10 points):
Briefly describe the purpose of the experiment. Make reference
to the appropriate section of our textbook by using the following
reference formatting (Silberberg, page #).
METHODS (20 points):
In your own words, describe the methodology that was used for
the experiment. Be sure to include the proper names of the
chemicals used.
RESULTS (25 points):
List the results of the experiment. Use tables where appropriate.
Be sure to include all data as well as observations.
CONCLUSION (10 pts):
Briefly discuss how the experiment reinforced what we have
learned in lecture conc
ncerning the topic. Again, use the following
reference in your discussion (Silberberg, page #).
Las Positas College, Chemistry 1A Lab Manual Fall 2012
Page 73
Experiment 9
Oxidation - Reduction Reactions: Predictions and Equations
Prelab - Complete the Prelab on page 79 before lab.
PURPOSE
To use patterns of reactions involved in oxidation-reduction in predicting redox
behavior.
To confirm these predictions from observations of some redox reactions.
To represent these redox reactions using balanced chemical equations.
.
Introduction
"Redox" is a convenient term for oxidation-reduction reactions. These reactions often
involve a transfer of electrons from one reactant to another. They can be recognized by
noting a change in oxidation number (charge) in one or more elements as they move
from being part of a reactant to being part of a product.
Other references are extremely helpful for writing correct chemical equations which
describe the reactions that you observe. See other sources for discussion and drill on:
1) ionic equations
2) solubility
3) assignment of oxidation numbers
4) common compounds to illustrate various oxidations states of N, Mn, O, CI, S, etc,
5) methods of balancing redox equations: the half-reaction method
6) acids and bases
Method
Redox reactions cannot be predicted as easily as simple partner exchange reactions, but it
can be done.
Transfer of electrons depends on two things:
1) the ease with which the reducing agent parts with electrons, and
2) the strength with which the oxidizing agent attracts the electrons.
You will not be required to make predictions of reaction products at this time. The factors
involved in such predictions will be discussed later. However, for this lab, we will
introduce the patterns of reactions involved in oxidation-reduction. For this purpose we
can make use of additional information that can be found at the back of this lab manual in
the Reference section:
The Activity Series of Metals
Nonmetals Common Oxidation States of Various Elements
.
An example of how the second of these tables can be used follows.
Page 74
Example: Predict the products of reaction between K2Cr2O, and H2O2 in acidic
medium. (It makes a big difference in many redox reactions whether the solution is
acidic or basic.)
Answer to above example:
System:
KCr2O7 and H2O2, acidic solution
Prediction:
From information on oxidation states, we see that Cr in Cr2022- is in its highest
oxidation state (+6). Therefore, Cr2O72- cannot be a reducing agent, but just might be a
good oxidizing agent. We also see that oxygen in H2O2 is in the - 1 state, and can go
either up to zero (as in O2) or down to -2 (as in H2O). However, if there is to be a
reaction with KCr207 H2O2 must act as a reducing agent, and yield 02. The fate of
chromium in Cr20may be revealed by a change in color.
Experiment:
A dilute solution of K2Cr207, was acidified with dilute hydrochloric acid, and H2O2
was added dropwise. The characteristic orange color of the Cr2O72- ion changed to a
greenish color, and bubbles of a gas were evolved.
Because Cr3+ is greenish in color in this case, and because a gas (02) is
expected from the reaction of H2O2, we feel justified in proposing the skeleton
equation:
(Acidic)
Cr2O72 (aq) +
H2O2 (aq)
Cr3+(aq) + O2(8)
To balance this equation, we use the technique of balancing redox equations discussed
in lecture and in your textbook. In the course of this process we will arrive at the two
half reactions:
6 e* + Cr2O72-caq) + 14 H(aq) → 2 Cr3+
(aq)
+ 7 H2O(1)
and
H2O2(aq) → O2(aq) + 2 *
Continuing the process, these two half reactions, properly multiplied and added, yield
the net ionic redox reaction:
Cr2O72(aq) + 3 H2O2(aq) + 8 H → 2 Cr3+ + 3 O2(g) + 7 H2O2(aq)
(aq) + 2 e
(ag)
(aq)
Las Positas College, Chemistry 1A Lab Manual Fall 2012
Page 75
PROCEDURE
Four possible redox systems will be investigated. Work individually! Specific directions
are given.
Redox Reaction Mixing Directions
SAFETY FIRST: Wear goggles throughout the experiment. Exercise care in handling
acids and bases. Brown fumes from reaction of nitric acid are toxic. Perform such
reactions in the hood. Do not breathe or smell the fumes. To quench (stop) a reaction,
add several volumes of water.
System 1: Cu and HNO3(aq)
Perform in HOOD!
Place 2 or 3 copper shot (about the size of BB's) in a small test tube in the hood. Add
about 5 mL dilute HNO3 (6 M). Look for a reaction over a period of 5 minutes or so.
Note the color of the fumes above the test tube. Covering mouth of test tube with
parafilm, waiting, and holding tube in front of a white background may help in seeing
the color of gas produced.
System 2: Cu and HCl(aq)
Repeat the procedure of System 1, but substitute 5 mL 6 M HCl for HNO3.
System 3: NaHSO3 and KMnO4
Do three variations: [3A], [3B] & [3C]
[3A] Acidic solution
In large test tube, 0.1 M NaHSO3 (2 mL) + 3 drops of 0.5 M H2SO4; then add 10
drops of 0.1 M KMnO4, drop by drop. The sign of reaction will be the loss of the
permanganate color.
[3B] Slightly basic solution
0.1 M NaHSO3 (2 mL) + 2 drops 1 M NaOH; then add 0.1 M KMnO4 dropwise
(about 10 drops).
[3C] Strongly basic solution
0.1 M NaHSO3 (2 mL) + 1 mL 6 M NaOH; then add 0.1 M KMnO4
dropwise (about 10 drops). Note two reactions: the first, a color change, is
unique to system [3C]; the second, a precipitate, is like that of [3B].
System 4: KI and FeCl3
In a large test tube place 0.1 M KI (2 mL) + 0.1 M FeC13 (2 mL)
Note the color and form of the precipitate of 12(s) formed. This precipitate
may be very slow to form.
Page 76
DATA
Enter the sample report form below in your notebook before coming to lab. Record
your observations, as you go along, directly into tables in your lab notebook in ink,
not into the sample tables provided here. Remember that the other references on
oxidation and reduction are extremely helpful in writing the correct equations to describe
these redox reactions. Answer as many of these questions as possible before coming to
lab.
#1: Cu and HNO3(aq)
#2: Cu and HCl(aq).
SYSTEM
Prediction: Do you expect
H2(g) to be evolved? Why
or why not?
What gas(es) might be
emitted if the nitrate ion
reacts?
Observations:
Net ionic equation for reaction, if any.
System #1: Cu and HNO3(aq)
System #2: Cu and HCl(aq)
System 3: NaHSO3 and KMnO4
Prediction:
Is the HSO3-ion a possible oxidizing agent?
If 'yes', to what species could it change?
Is it a possible reducing agent?
If 'yes', to what species could it change?
Is the MnO4 ion a possible oxidizing agent?
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