Lab 16: Metals and Oxidation

Introduction
Have you ever wondered why some rings will turn a fin‐
ger green, while others won’t?
The ring you just bought a couple of weeks ago is already turning your
finger green. Another ring that you have had for years still looks al‐
most new. Why is this? Knowing how reactive different metals are is
extremely important. This property helps us to decide which metal to
use for a particular application. Some metals are so nonreactive that
they will not react with even the strongest acids. Many will only react
with certain acids, while others are so reactive that they will react with
water. This is why jewelry is often made out of gold and not ordinary
iron. Iron is more reactive, and will gradually rust when exposed to the
oxygen and moisture in the air. Gold, on the other hand, will not react
easily with anything.
You know that gold is less reactive than iron, but what about other
metals? How do you know which ones are more reactive? The activity
series of metals places elements in order by their reactivity. Metals
higher on the list give up their valence electrons more easily than the
metals below them, meaning that any metal on this list will displace a
metal below it in a reaction. This can easily be observed if the less re‐
active metal is in an aqueous solution during the reaction, as it will
precipitate out of the solution when replaced. A partial activity series
list is shown in Figure 2.
As a part of this lab, you will observe the reaction of an acid and zinc
metal and compare it to the reaction of the same acid with iron metal.
The reaction between the zinc and the acid can be written as follows:
2 H+( a q ) + Zn(s) → Zn2+(aq) + H2(g)
In this reaction, the zinc transfers its valence electrons to the hydrogen in an acid solution. (Remember a superscript plus
sign indicates that the compound is missing a valence electron a minus sign indicates that it has an extra valence electron.)
This causes the hydrogen in the acid solution to separate as hydrogen gas while the zinc metal forms zinc ions dissolved in
the solution. The zinc is said to be more electropositive than the hydrogen it displaces.
Figure 1: The Statue of Liberty in New York Harbor is
made of copper, and was originally the color of a
penny. The statue has gradually acquired its green
color due to the natural oxidation of copper when
exposed to air and water.

Concepts to explore:

 Observe an oxidation‐reduction reaction
 Use the properties of a reaction product to verify its identity
 Rank the reactivity of certain metals in a weak acid, and compare it to
their order in the Activity Series of Metals
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Reactions like this that involve the transfer of electrons are called oxidation‐reduction reactions, or “redox” for short. Oxi‐
dation generally describes the loss of electrons by a molecule, while reduction describes the gain of electrons. In this reac‐
tion, since zinc is losing electrons, it is being oxidized.
Zn (s) → Zn2+ (aq) + 2eHydrogen is gaining electrons in this reaction, so it is being reduced:
2 H+(aq) + 2e– → H2 (g)
This is often remembered by the phrase “LEO the Lion goes GER.” LEO stands for Losing Electrons Oxidation, and GER
stands for Gain Electrons Reduction.
In Lab 15 we called this same reaction a single replacement reaction. How can this be? It has to do with how the reactions
are sorted. Just like laundry, there are several different ways to categorize and sort reactions depending on the applica‐
tion. For example, when doing laundry, you may need to separate red clothes from white ones. Other times you separate
the light colored clothes from the dark colored clothes. Sometimes reactions are categorized as they were in Lab 15, and
sometimes they are categorized as redox and non‐redox reactions.
But how do we know just by looking at this reaction equation that it is an redox reaction and that electrons are being trans‐
ferred? The oxidation number must first be assigned to each of the atoms involved on both sides of the reaction equation.
Let’s look at a reaction between aluminum and hydrochloric acid.
6 HCl ( a q ) + 2 Al(s) → 2 AlCl 3 (aq) + 3 H2(g)
There are several rules to help you determine oxidation numbers in a reaction:
 The oxidation number of an element by itself is zero. This means that the aluminum metal (Al) on the right side of
the above equation has an oxidation number of 0. The hydrogen gas (H2) on the right side of the equation also has
an oxidation number of 0.
 When an atom exists as a simple ion in a substance, the oxidation number is the same as its charge in the com‐
pound. The chloride (Cl) ion has a negative charge, so its oxidation number in is ‐1 in each instance it appears.
 For a neutral compound, the sum of the oxidation numbers is always zero. A polyatomic ion’s oxidation number
equals the charge on the ion. Since HCl is neutral, the sum of the oxidation numbers should equal zero. We know
from above that the Cl ion has an oxidation number of ‐1; For the sum of the oxidation numbers in HCl to equal 0,
the H must have an oxidation number of +1. Similarly , for the oxidation of AlCl3 to be zero, the Al must have an
oxidation number of +3 to balance out the three Cl atoms (each with oxidation numbers of ‐1).

Lab 16: Metals and Oxidation

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Loses electrons easily
(more easily oxidized)
Do not lose electrons easily
(not easily oxidized)
Figure 2: Activity Series of Metals (Partial List)
+1 ‐1 0 +3 3(‐1) 0
Lab 16: Metals and Oxidation
Sometimes it is helpful to write the oxidation numbers over the atoms as is shown below:
6 HCl ( a q ) + 2 Al (s) → 2 AlCl 3 (aq) + 3 H2(g)
The activity series (Figure 2) of metals also indicates how easily a metal will cause a release of hydrogen in redox reactions.
Metals at the top portion of the list release hydrogen merely by being placed in cold water. As you go down the list more
harsh conditions are required to release hydrogen. Metals towards the bottom do not release hydrogen even when placed
in acids.
Sometimes a metal forms a thin oxide coating that protects it from reacting any further with its surroundings. Aluminum is
fairly reactive and does this quickly. On the other hand, when iron is exposed to oxygen it corrodes entirely into iron oxide
(rust) over time. For this reason, iron is commonly coated in zinc to prevent oxidation.
In this laboratory exercise, you will compare the reactivity of the redox reactions of zinc and iron with a weak acid solution:
saturated citric acid. Citric acid is a different type of acid than HCl, but works in a similar way.
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Pre‐lab Questions

1. What is the oxidation state for each atom in the following reaction:
4Fe + 3O2 → 2Fe2O3
a. Elemental iron (Fe)
b. Elemental oxygen (O2)
c. One iron atom in Fe2O3
d. One oxygen atom in Fe2O3
2. Which element was oxidized and which element was reduced in the above reaction equation?
a. Element oxidized
b. Element reduced
3. From the Activity Series of Metals, determine the order of reactivity of the following metals: Ni, Au, Fe, Ca, Zn,
and Al.
Most Reactive → Least Reactive
Lab 16: Metals and Oxidation
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Procedure
1. Label two test tubes Zn and Fe with the permanent marker, and place them in a test tube rack.
2. Tip the test tube labeled Zn slightly and let a galvanized nail gently slide into it with the top of the nail going
down first.
3. If necessary, lightly sand the iron nail that is not galvanized to remove any rust and wipe it clean. Gently
place it in the test tube labeled Fe the same way you did previously.
4. Add approximately 3 to 5 mL of saturated citric acid solution to each test tube.
5. Make initial observations and continue recording observations after one minute, three minutes, and five
minutes. HINT: The most notable observations are how quickly bubbling occurs and how violently the bub‐
bling of each continues.
6. After observing the reactions for five minutes, rank the two metals in order of their reactivity. Compare your
results with the actual reactivity series.
7. To clean up, separate the acid solution from the metals by pouring them into a 250 mL beaker while leaving
the metals in their test tubes. This is called decanting. Rinse the test tube containing the metals several
times with water and add the rinses to the beaker. To neutralize the acid, add small amounts of baking soda
to the acid solution in the beaker and stir. Continue this until no more gas forms. Pour the liquid down the
drain, and throw the metals in the trash.
Experiment: Metal Reactivity
In this lab you will use what you know about chemical reactions and oxidation to examine how two metals (zinc and iron)
react differently with a citric acid solution. You will then draw conclusions about the reactivity of the metals based on what
you observe.
Lab 16: Metals and Oxidation
Materials
Safety Equipment: Safety goggles, gloves

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Table 1: Observations of Reactions

 Summarize your observations and list the order of reactivity of the metals that you observed:
Data
Lab 16: Metals and Oxidation
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Post‐lab Questions
1. Based on what you observed, what is one of the products formed in Part 1? How do you know?
2. Did the order of reactivity you determined in Part 2 match the order given in the Activity Series of Metals?
Explain.
2. Do you think the following reaction would occur? Explain your answer.
FeCl 2 ( a q ) + Cu(s) → Fe(s) + CuCl 2 (aq)
4. How do you think acid rain might affect the rate of rusting of metal?

Lab 16: Metals and Oxidation

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Last Updated on April 25, 2020