chemical dominoes
Section 4
The Metal Activity Series: What
Can Destroy a Metal?
Section Overview
In this section, students build a circuit to light
the red LED using a magnesium strip as wire.
They then devise a way to turn off the circuit by
destroying the magnesium, using three mystery
chemicals. Students will also explore the reactions
of various metal cation solutions, as well as
acid, with different metals. They interpret their
observations based on the activity series of metals,
and use their observations and the activity series
to figure out the identities of the three mystery
chemicals. Students also invoke the activity series
to explain why some common metals are more
suitable for particular purposes, while others
are not. They practice writing and balancing
oxidation-reduction reactions.
Background Information
Metals
682
Metals have several properties that characterize
them as a class of elements. The definition of
metal is a functional one, and has grown over
time as humans have come to understand more
chemistry. Since it has been known for thousands
of years that metals are shiny solids that (with
the exception of mercury) can be beaten into
shape and that generally melt at very high
temperatures, they have a long history of use
in structures, protection, decoration, and for
making cookware. More recently, as metals
began to be used for making hardware and
electrical circuits (within the last few hundred
years), good thermal and electrical conductivity
have been added to the list of characteristics
that define them. Most recently (within the last
century), since we have known about electrons,
the definition of metal has also come to include
the characteristic of being composed of atoms
that give up electrons when they react.
Although hydrogen
Activity Series of Metals
shares this last
Name
Symbol
characteristic with
lithium
Li
metals, it doesn’t share
the other properties,
potassium
K
so hydrogen is not
calcium
Ca
a metal. However,
sodium
Na
because hydrogen
magnesium
Mg
atoms each give up
aluminum
Al
an electron when they
zinc
Zn
react, and because
hydrogen is the simplest
iron
Fe
of all the elements
tin
Sn
(meaning its behavior
lead
Pb
is easier to explain),
hydrogen*
H
metals are often
copper
Cu
compared to hydrogen.
mercury
Hg
A key question that
chemists often ask is,
silver
Ag
“does metal X give up
platinum
Pt
electrons more or less
gold
Au
easily than hydrogen?”
*Hydrogen is not a metal but it
This can easily be tested is useful for testing the activity
of metals
by placing the metal
in an acidic solution.
Because hydrogen plays such a vital role in the
testing of metals, it is often placed on the list of the
Metal Activity Series.
The Metal Activity Series
The information contained in the Metal Activity
Series originated with alchemists, who were
medieval chemists trying to create gold out of
other substances. The Metal Activity Series is
a list of metals arranged in the order of how
reactive they are. For example, lithium metal
reacts more readily with a solution of limewater
(containing calcium ions dissolved in water) than
potassium metal does, both creating calcium
metal. Because their activity is lower than
SEction 4 The METAL ACTIVITY SERIES: What can destroy a metal?
Any metal that is above another metal on the list
will react with the lower metal’s ions in solution.
For example, pure zinc metal will react with
copper ions, but pure copper metal will not react
with zinc ions. In symbolic notation, this can be
expressed as:
Zn(s) + Cu2+(aq)
→
Cu(s) + Zn2+(aq)
Two Half-Reactions Make a
Whole Redox Reaction
Redox (or reduction-oxidation) reactions are the
symbolic representation of metal activity at the
particle level. The main idea is that one kind of
metal atom (atoms of neutral metal A) gives up
electrons to form positively charged ions (An+),
while a second metal ion (Bm+) accepts those
electrons to form a pure metal (atoms of neutral
metal B). The giving up of electrons is called
oxidation, while the accepting of electrons is
called reduction.
But
Cu(s) + Zn2+(aq)
For example, the activity series tells us that zinc
metal can react with copper ions:
(WILL NOT REACT TO FORM)
Zn(s) + Cu2+(aq)
Hydrogen’s reactivity falls between that of
lead and copper. That is, pure lead metal
will react with a solution containing H+ ions
(slowly), and lead ions will form while pure
hydrogen gas is produced:
While we can pretend this happens in two
steps, in reality the two processes occur
simultaneously. In the first step, a zinc metal
atom gives up two electrons (this is considered
the oxidation half of the complete reaction,
or the oxidation half-reaction):
Zn(s) + Cu2+(aq)
Pb(s) + 2H+(aq)
→
H2(g) + Pb2+(aq)
But the reverse reaction will not occur:
H2(g) + Pb2+(aq)
(WILL NOT REACT TO FORM)
Pb(s) + 2H+(aq)
(Note: In reality, hydrogen ions do not exist in
aqueous solutions. Instead, hydronium ions,
H3O+, are the form of the acid present. H+ is
just a shorthand notation for H3O+.)
The investigation in this section provides
students with the Metal Activity Series and
challenges students to figure out how the series
is related to common facts known about the
interactions of metals with other chemicals.
Students also perform experiments and use the
Metal Activity Series to interpret their findings.
Thus, they develop both practical (common
experience) and laboratory experience in
interpreting results and learn the importance of
consistency in interpretations. Finally, they are
introduced to the particle level explanation and
the symbolic representation of metal activity.
chapter 4
calcium, none of the other pure metals on the list
react with limewater solution to produce calcium
metal. (Of course, since lithium and potassium are
so reactive, it is difficult to obtain pure lithium
metal and pure potassium metal for such a test.)
Zn(s)
→
→
Cu(s) + Zn2+(aq)
Zn2+(aq) + 2e–
In the second step, a copper ion accepts those two
electrons and becomes an atom of pure copper
metal (this is considered the reduction half of the
complete reaction, or the reduction half-reaction):
Cu2+(aq) + 2e–
→
Cu(s)
Of course, there are no naked electrons that
“float” around in the solution, and electrons are
transferred during interactions between zinc atoms
and copper ions.
The example of zinc metal reacting with copper
ions involves the transfer of two electrons from
a zinc metal atom to a copper ion. Sometimes,
however, a metal atom will give up a different
number of electrons than the other metal
ion is able to receive. When this happens,
stoichiometric coefficients must be adjusted to
reflect the conservation of electrons. For example,
if aluminum metal atoms each give up three
electrons, and copper ions can only accept two
electrons, then it takes two aluminum metal
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chemical dominoes
atoms to satisfy three copper ions. Another way
of looking at this is in terms of the least common
multiple of electrons, which is 2 × 3 = 6.
(Al
→
Al3+ + 3e-) × 2 = 2Al
(Cu2+ + 2e
→
→
2Al3+ + 6e-
Cu) × 3 = 3Cu2+ + 6e2Al + 3Cu2+
→
→
3Cu
2Al3+ + 3Cu
The overall reaction, that is the sum of the two
halves of the reaction, is shown in the box. The
objective in arriving at the overall reaction is that
the oxidation electrons and the reduction electrons
cancel out.
Where the Activity Series Breaks
Down in Predictive Power
The Activity Series of Metals is an oversimplification, and does not describe the behavior
of real metals. There are three reasons for this.
First, most metals are not available in pure form.
Instead, since most metals are somewhat reactive,
atoms on their surfaces react with oxygen in the
air over time, and thin films of metal oxide build
up on the surfaces of metals. The second reason
the Activity Series is an over-simplification is that
metals often react with hydrogen ions available in
the water solvent. Third, some metals have more
than one possible oxidation state, so depending on
the reduction potential of the agent oxidizing such
a metal, different oxidation states are possible.
The first complication explains what can be
perceived as counter-examples to the Metal
Activity Series. For example, the Activity Series
would indicate that aluminum is fairly reactive,
yet we make cooking utensils and cookware
out of aluminum, and we use aluminum foil
to wrap and protect food (which is usually
neutral to slightly acidic). In reality, aluminum
oxide films are very unreactive, and therefore
protect the aluminum metal beneath them. In
fact, concentrated nitric acid is often stored in
aluminum tanks, even though the Metal Activity
Series would predict a strong reaction between
aluminum metal and acidic hydrogen ions (recall
684
that hydrogen’s activity rests between that of lead
and copper). It is also interesting to note that, by
contrast, aluminum oxide is readily dissolved in
basic solutions. Since soaps and many cleaning
agents are strongly basic, aluminum is never used
inside dishwashers and washing machines. The
oxide layer buildup on metals is the reason why
students are advised to use sandpaper to “clean”
the surface of a metal before allowing it to react.
The second complication to the predictive power
of the Activity Series rests in the fact that water
(the solvent for aqueous solutions) contains some
hydronium ions (H+ or actually H3O+). Therefore,
when considering any reaction between two
metals in the Activity Series, what happens in
reality is that any metals above where hydrogen
would appear in the Activity Series actually react
significantly with H+. The real processes that occur
are more complicated than what is presented in
the Chem Talk section of the Student Edition.
However, the basic premise is the same, so a
simplified version is presented in order to teach
students the concept. You may want to consider
sharing the more detailed version with students
who can handle more definitive explanations.
The third complication explains why some metal
ions (such as Au+) are able to oxidize metals with
multiple oxidation states (such as Fe, which can be
oxidized to either Fe2+ or Fe3+) to different states
than other metals, while other metal ions (such
as Cu2+) do not have the power to do so. The
reduction potentials for the examples given are:
Standard Reduction Potentials
in Aqueous (Acidic) Solution at 25ºC
Reduction Half-Reaction
→ Au
Fe3+ + 1e- → Fe2+
Cu2+ + 2e- → Cu
Fe2+ + 2e- → Fe
Au+ + 1e-
Eo (volts)
1.68
0.771
0.337
-0.44
SEction 4 The METAL ACTIVITY SERIES: What can destroy a metal?
reduction:
Au+ + e- → Au
E = +1.68 V
oxidation step 1:
Fe2+ → Fe3+ + 1e-
E = –0.771 V
oxidation step 2:
Fe → Fe2+ + 2e-
E = +0.44 V
total oxidation:
Fe → Fe3+ + 3e-
o
o
reduction:
Cu2+ + 2e-
o
least common multiple of electrons, n = 3 (see
Section 7, Background Information for further
explanation on using reduction potentials).
Cu
o
E cell = +0.337 V
oxidation:
Fe → Fe2+ + 2e-
E cell = +0.44 V
net redox reaction:
Cu2+ + Fe → Fe2+ + Cu
E cell = +0.777 V
o
o
Some Basic Circuits
o
E = –0.331 V
→
chapter 4
To explain this, one must use the reduction
potentials of the half-reactions. In the first
example, Au+ is able to oxidize Fe to Fe3+ because
the reduction potential of Au+ is very large.
Simplest Circuit
It is presumed that students have already had
exposure to some basic 9th grade level physical
science, including circuits. However, if students have
not had this exposure, it may be necessary to provide
them with some background information about a
few simple circuits they might consider using for the
Chapter Challenge.
net redox reaction:
3Au+ + Fe
→
Fe3+ + 3Au o
E cell = +1.35 V
Meanwhile, copper ions, Cu2+, are not able
to oxidize Fe to Fe3+ because the reduction
potential of Cu2+, at 0.337 V, is not large enough.
o
It would lead to a value of E cell = 0.006 V,
which while positive, is very nearly zero, and
not practicable. Therefore, Cu2+ can oxidize
Fe only to Fe2+, as follows:
direction of
electricity flow
resistance
of some
kind
power
source
Figure 1. Simplest circuit
Notes
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chemical dominoes
In order to build simple circuits, the most
important idea students need to understand is
that to make electricity flow, there must be a
complete circuit that starts with a power
source and ends back at that power source.
The electricity has to flow through materials
that are able to conduct electricity.
The simplest circuit students might build
would involve an operating circuit that is
disconnected (electricity stops flowing)
when part of the electrical circuit is destroyed
by the application of a chemical to one of
the conducting metals (see Figure 1). If
electricity is flowing through a complete
circuit, then one way to “turn off” the flow
is to interrupt the circuit by creating a break
in an electricity-conducting material that is
part of the circuit. Since electricity cannot flow
through the air that fills the break, the circuit
is no longer complete.
Circuit Within a circuit
A more useful circuit students might build as part of the
Chapter Challenge is one that contains a circuit within a
circuit. The inside circuit acts as an on-off switch. The inside
circuit stops electricity from flowing in the outside circuit,
because the inside circuit has a much lower resistance.
Breaking the inside circuit (by destroying some of the
electricity-conducting metal along the inside circuit) then
allows electricity to flow in the outside circuit.
direction of
electricity flow
power
source
inside
circuit
large
resistance
outside
circuit
small
resistance
LED
Figure 2. Circuit within a circuit
Learning outcomes
learning outcomes
686
Location in Section
Evidence of Understanding
Use proper materials to
light an LED and explain the
procedure.
Investigate
Part A, Steps 1-2
Chem Talk,
Reflecting on the Section
and the Challenge
Students are able to create an electrical circuit that lights
an LED and provide answers that match those in this
Teacher’s Edition.
Use the Metal Activity Series
to determine which metal of a
given pair is most reactive.
Investigate
Part B, 1-8, Part C, 1-5
Chem to Go
Questions 2-4, 6, 8-11
Students successfully complete the Investigate steps and
provide answers that match those in this Teacher’s Edition.
SEction 4 The METAL ACTIVITY SERIES: What can destroy a metal?
chapter 4
Notes
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Section 4
Materials, Chemicals, Preparation, and Safety
(“per Group” quantity is based on group size of 4 students)
Materials and Equipment
Materials
(and Equipment)
Quantity
per Group
(4 students)
Chemicals
(Note: the solutions below should
be stored in small dropping
bottles for student use)
Quantity
per Class
(24 students)
LED, Red, with clear lens
1
Resistor, 10 ohm
1
Aluminum chloride, AlCl3, 0.1 M
100 mL
Resistor, 100 ohm
1
Copper (II) chloride, CuCl2, 0.1 M *
100 mL
Battery, 9 V
1
Iron (III) chloride, FeCl3, 0.1 M
100 mL
Wires with alligator clips
4
Magnesium chloride, MgCl2, 0.1 M
100 mL
Zinc chloride, ZnCl2, 0.1 M
100 mL
2
Hydrochloric acid, HCl, 0.1 M *
100 mL
1
Silver nitrate, AgNO3, 0.1 M*
(store in brown dropping bottles)
100 mL
Aluminum, strip,
5 mm x 5 cm x 1mm
Copper, strip,
5 mm x 5 cm x 1mm
Iron, strip,
5 mm x 5 cm x 1mm
1
Magnesium ribbon
1
Silver (may come in a small roll or
a sheet and can then be cut
to size.)
1
Zinc, strip,
5 mm x 5 cm x 1mm
2
Steel wool (used to polish
metal strips)
1
Large microwell plate, 24 well
2
Scissors (used to cut metal
strips into small pieces for
the well plate)
* mystery solutions
Teacher Preparation
1 pair
Tweezers
1
Watch glass, 5 cm or larger
1
Materials
(and Equipment)
688
Chemicals
Quantity
per Class
Dropping bottles, small
15
Dropping bottles, small, brown
(storage of silver nitrate solution)
2
Part A – Setting up each group’s materials ahead
of time in boxes that students can carry to their
workstations will save considerable time. Test all
the LEDs beforehand. Have extra LEDs on hand
in case LEDs burn out. Prepare the three mystery
solutions* ahead of time. These are 0.1 M copper
(II) chloride, 0.1 M silver nitrate, and 1.0 M HCl.
Be careful with the silver nitrate solution, as it will
stain skin and clothing. The “mystery” chemicals
should not be labeled with correct names, but just
as A, B, C. It is the job of the students to identify
the chemical in each. To make it a little trickier,
you could put each “mystery” chemical in a
brown dropping bottle.
SEction 4 The METAL ACTIVITY SERIES: What can destroy a metal?
chapter 4
Depending on whether students have already
studied physics, you may wish to show students a
circuit that will light the LED when another part
of the circuit is destroyed. To set this up, initially
create a circuit that short-circuits the LED with
low resistance. Then, cut current through the
short-circuit, and the LED will light.
0.1 M MgCl2—Dissolve 2.03 g of magnesium
chloride hexahydrate (MgCl2•6H2O) in 90 mL
of deionized water and then adjust volume to
100 mL. Store in two small dropping bottles.
0.1 M AgNO3—Dissolve 1.7 g of silver nitrate in
90 mL of deionized water and then adjust volume
to 100 mL. Store in two dark dropping bottles.
0.1 M ZnCl2—Dissolve 1.36 g of zinc chloride
in 90 mL of deionized water and then adjust
the volume to 100 mL. Store in two small
dropping bottles.
1.0 M HCl—Use 1.0 M HCl (remaining
from Section 1 and 2) OR dissolve 8.3 mL of
concentrated HCl in 90 mL of water and then
adjust final volume to 100 mL. Store in two
dropping bottles – make certain that the HCl
cannot be confused with the metal salt solutions.
Part B – Prepare solutions ahead of time and
place in small bottles. Label bottles and
droppers with colored tape, effectively colorcoding solutions. This prevents students from
contaminating solutions. If solutions become
contaminated, the experiments will not work.
Retain some of each prepared solution so that you
can empty contaminated small bottles and refresh
with new solution.
Students can share but need to take care not to
cross-contaminate any of the solutions.
0.1 M AlCl3—Dissolve 2.42 g of aluminum
chloride hexahydrate (AlCl3•6H2O) in 90 mL of
deionized water and then adjust the volume to
100 mL. Store in two small dropping bottles.
0.1 M CuCl2—Dissolve 1.7 g of copper (II)
chloride dihydrate (CuCl2•2H2O) in 90 mL
of deionized water and then adjust volume to
100 mL. Store in two small dropping bottles.
The “mystery” chemicals are the CuCl2, HCl
and AgNO3 solutions listed above. Place a
small amount of these in separate dropping
bottles marked in code (such as, A, B, C).
Safety Requirements
• A
ll activity in the laboratory requires
goggles and aprons.
• M
agnesium fragments can be digested in
dilute hydrochloric acid and poured down
the drain with plenty of water.
• O
ther metal fragments can be disposed
in the garbage.
• S mall amounts of solutions can be washed
down the drain with plenty of water.
• S ilver nitrate will stain skin and clothing
and so proper techniques must be used.
• W
ash hands and arms before leaving the
laboratory area.
0.1 M FeCl3—Dissolve 1.99 g of iron (II)
chloride tetrahydrate (FeCl3•4H2O) in 90 mL
of deionized water and then adjust volume to
100 mL. Store in two small dropping bottles.
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chemical dominoes
Meeting the Needs of All Students
Differentiated Instruction
Augmentation and Accommodations
Learning issue
Making predictions
about reactivity of
metals
Reference
Investigate
Part B, 4.
Augmentation
•C
onsider having students insert the copy of the table as Part A, Step 1. instead
of Part A, Step 4. so that they may insert “more easily affected” directly on the
table at the top left and “less easily affected” at the bottom left connecting
the two comments with a double arrow.
Accommodations
•C
opying the Activity Series of Metals table into their logs may require
accommodation for some students who experience difficulty with visual motor
integration. Provide a copy of this table they can insert in their Chem logs.
•C
onsider referring students to Chem Talk, where a paragraph under
Reduction clearly explains how relative reactivity affects reactions.
Recalling previously
learned concepts
Investigate
Part B, 5.
Augmentation
• In Chapter 1, students learned that an element that gives an electron is
positively charged because the ion that remains has one more proton (+) than
electron (–). For some who think losing an electron should make the atom
negatively charged, this may be counter-intuitive. Remind students of the
concept they learned in Chapter 1, and check that they understand the logic.
•D
rawing a diagram of an atom to show this may help students with language
issues understand the concept.
Learning selfassessment
Investigate
Part B, 6.,7.
Augmentation
• S tudents’ awareness of their mastery or lack of mastery of a skill is important in
developing executive functioning. They may be unsure of their answers to
Step 6.a-6.d). Having them check their own work by comparing their answers to
the first seven lines of Step 7 will allow them to evaluate their own work and
allow the teacher to help them with any misunderstandings.
Recalling previously
learned concepts
Investigate
Part B, 8.
Augmentation
•A
few students will remember that Al gives up three electrons because it has
three valence electrons. Others will not remember the atomic number of Al,
but will remember the principle and assume that it has three electrons in its
outer shell. Still others will have forgotten their previous learning altogether.
Reminding students of the concept they learned previously, that elements
combine in certain proportions, depending on the number of electrons in their
outer shells, and checking that they understand it, will help them build on their
prior knowledge.
Copying the table
690
Augmentation and Accommodations
SEction 4 The METAL ACTIVITY SERIES: What can destroy a metal?
Testing reactions
between metals
Reference
Investigate
Part C
Writing balanced
equations that
show the change in
electrons
chapter 4
Learning issue
Augmentation and Accommodations
Augmentation
• In Step 1.c), students are required to write balanced equations as part of
testing reactions in pairs of metals. Even if most understand the principles
taught thus far, they may not be able to write their own balanced equations
using proper notation. Check for understanding to see who can write an
equation independently and put at least one in each group to ensure all groups
can complete the next part of the lesson successfully.
• In Step 1.d), students are asked to make predictions based on inferences made
from the Activity Series of Metals table and their observations. Model the
thinking process out loud for students having difficulty, giving them all the
things you think about except the answer. Help them establish the pattern of
thought that will lead to a logical answer. Guide them through the process
with questions, not directives. Let them explore their wrong answers before
correcting them. Set up others who do not need help in other groups to work
independently while you give help to those who need it.
Accommodations
•T
his section calls for careful, legible recording of data. Make one group
member with this skill responsible for recording the group’s observations.
• Consider giving him/her a blank table to complete.
Strategies for Students with Limited English Language Proficiency
Learning issue
Reference
Augmentation and Accommodations
Background
knowledge
What Do You
Think?
Students may not have prior knowledge of electrical circuitry. Allow for some
group discussion.
Vocabulary
Investigate
Check for understanding of the term, “diode.” Check understanding of usage
for the word, “efficient” as it applies to the concept being taught. Phrases such
as “schematic diagram” and “correct orientation” may require elaboration. The
word, “apparatus” may also not be in the students’ expressive vocabulary. Check
for oral production of formula.
Background
knowledge
Vocabulary
Comprehending text
Chem Talk
Provide some history on the field of “alchemy” to help provide background
knowledge. Provide insights into the various derivatives of “oxidize” that are used
in this section. The teacher may want to isolate sections of the text and either share
the reading aloud or have small groups of students read the material in sections.
Asking literal questions will allow the teacher to assess understanding.
Supporting details
Research skills
What Do You
Think Now?
Ask students to explain their choices in answering the questions to ensure
comprehension.
Comprehension
Vocabulary
Chem Essential
Questions,
Chem to Go
Ensure that students clearly understand their task when they are asked to “explain.”
Allow students to discuss the possible answers to the questions. Have students work
in groups to achieve a consensus on the possible answers.
Following directions
Inquiring Further
Have students design experiments in small groups or with a partner. Make sure that
they know the criteria based upon the possibilities of what “design” means. Clarify
how they are to “design” a report on their experiment.
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chemical dominoes
Section 4
Teaching Suggestions
and Sample Answers
What Do You See?
The main purpose of the What
Do You See? section is to engage
the students and open up a
discussion of their perceptions.
There may be different levels of
observations. For example, some
students may say, “The cat is
running away!” Others might try
to interpret the scientific concept
that is represented by the
illustration. Both responses are
good and it is your task to elicit
responses without judgment.
Some of the items depicted in
the illustration that will have
more relevance later are a
battery, and an LED.
What Do You Think?
These are open questions that
should bring forth a variety of
responses. You are not looking
for correct answers, but for a
discussion of ideas. The primary
purpose is to elicit students’ prior
knowledge and conceptions.
Most students will have a limited
knowledge of which and how
chemicals might destroy a metal
but some might suggest an acid.
To destroy a metal wire, you
would have to use a solution
containing the cation of a metal
(or hydrogen) that is less reactive
than the metal in the wire. If
the metal composing the circuit
is more reactive than hydrogen,
then an acid (H+) could be used
as well. The less active the metal
in the circuit, the less easily it can
be destroyed. For example, gold
is very difficult to force into a
reaction (i.e., to destroy).
Students’ Prior Conceptions
1. Electrons start in the battery and return
to the battery.
This slight misconception, and many others,
accompany the topic of circuits. Since circuit theory is
not a component of the usual chemistry curriculum,
very little time should be spent on this with students.
Circuits are used in this chapter to the minimum
extent possible, to allow for the exploration of
chemical processes that generate and use electricity.
However, since this particular misconception can get
in the way of understanding electrochemical cells
later on, it may be worth a brief demonstration to
eliminate this misconception.
One way to do this calls for a hula hoop and two or
more volunteers to demonstrate. Have the students
hold the hula hoop as shown, and explain that the
hoop represents a complete circuit in which electrons
can circulate:
Ask the students to remain where they are. Then,
guide the hula hoop to turn in a circle, sliding
through the hands of the students. Ask students if
they felt the “electrons” passing through
692
What Do You Think?
a chemist’s response
their hands. This illustrates that when a circuit is
operational, electrons all along the circuit
move simultaneously.
2. Metals and metal ions are the same.
This is the misconception that Step 3 in Part B of
Investigate attempts to dispel. The iron in staples, for
example, has very different properties from the iron
in vitamin tablets. The iron in staples is in the form
of neutral metal atoms, while the iron in vitamins
is positively charged metal ions and part of an ionic
salt with a corresponding anion. This salt can dissolve
in water and be ingested by people for nutrition
while the metallic iron cannot. The misconception
arises in large part from a confusion of the terms
in common parlance, and especially in advertising
where vitamin tablets are sold as containing iron,
rather than iron cations. Students will not be able
to understand oxidation and reduction unless they
can understand this important difference. If students
have already studied the bonding differences
between various solids (covalent, ionic, metallic),
then you can draw upon this in convincing students
of the difference. If not, the examples in Step 3 of
Part B and other examples like these, should be used.
3. Electrons are present in solution when
oxidation-reduction reactions occur.
This misconception stems from breaking down
overall redox reactions into an oxidation halfreaction and a reduction half-reaction, where
electrons are explicitly shown. There are no
SEction 4 The METAL ACTIVITY SERIES: What can destroy a metal?
chapter 4
Investigate
Part A:
Destroy the Circuit
and Light the LED
1.
Students should be able to make
the LED light by connecting the
short lead on the LED to the
negative battery terminal, and
the long lead on the LED to the
positive battery terminal. No
pieces of wire are necessary if the
LED leads are long enough. If
not, then the circuit will require
a minimum of one piece of wire.
An unexpected phenomenon that
students should note is that the
LED will not light when current
flows in the opposite direction.
unattached electrons floating around in solution.
The electrical charge is carried through the solution
by the anions in solution.
4. Oxidation means oxygen must be involved.
This point of confusion is certainly most understandable,
given the common root in “oxidation” and “oxygen.”
However, while oxygen is a common oxidizing agent in
oxidation-reduction (redox) reactions, many oxidationreduction reactions involve no oxygen. Most of the
reactions seen in this section can be pointed out to
students as examples of oxidations that occur that do
not involve oxygen. The key idea that “oxidation”
means the loss of electrons by any chemical species
should be emphasized. It is likely that the term arose
from the many examples where oxygen in air caused
the oxidation of metals.
5. Reduction is a loss of electrons.
Because students are taught that
the defining feature of a redox
reaction is the transfer of electrons
from one species to another, they
frequently assume that “reduction”
means a reduction in the number of
electrons. In fact, the opposite is true
– reduction is the gain of electrons.
It is helpful to provide students with
several examples of redox equations
and help them to identify what is
reduced in the reaction.
6. Any two chemicals will undergo a chemical reaction
when combined.
Although common sense and experience runs to the
contrary, students may not have had the occasion in
previous science classes to consider situations in which
substances are combined but do not react. Instruction
typically focuses on situations where reactions do occur
when substances are mixed. Thus, it is logical that many
students assume that if some metals and metal ions
react, then any combination will work. To facilitate
learning that redox reactions only occur when there
is a force driving the reaction, encourage students to
connect and resolve both common and laboratory
experiences – ultimately this should lead to transfer of
knowledge beyond the classroom.
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chemical dominoes
2.
Connected as the diagram shows,
electricity will flow through
the magnesium strip. It does so
because magnesium is a metal,
and metal conducts electricity.
A “safe” surface might mean
carrying out the experiment on a
watch glass, so it is easy to clean
up. A spot plate will also serve.
teaching tip
The magnesium strip can be filed
or sanded down to save time,
because a thin piece of magnesium
will react and break more quickly.
2.a)
Students should notice that
the mystery solution (that only
you know is CuCl2) is a blue
solution, while the other two are
colorless. They will also observe
that one solution is in a brown
bottle (which you know to be
silver nitrate). Those that react
with the metal will interrupt
the circuit, lighting the LED,
but each “mystery” chemical
will react at its own speed. The
copper solution (blue) and the
silver solution (brown bottle)
will likely also produce black or
reddish-brown precipitates of
the corresponding salt.
2.b)
Theoretically, the mystery
solution that is silver nitrate
should do the best job (quickest)
of destroying the circuit.
The mystery solution that
is hydrochloric acid should
produce bubbles (of H2 gas) as
the reaction proceeds. All three
solutions should react with the
magnesium strip and “destroy”
it so that the LED lights up
teaching tip
In one 45 minute class period,
students can finish Part A and
begin Part B. Part B can be stopped
and continued the following day
at any step before Step 8.
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Part B:
What Kinds of Chemicals
Affect Metals?
1.
Students may not be able to
answer these immediately. Some
discussion of the first few will
give students an idea of the
pattern by which to reason out
the rest.
SEction 4 The METAL ACTIVITY SERIES: What can destroy a metal?
4.
chapter 4
Be sure students make the
column titled “Changes when
atoms give up electrons” wide
enough to hold a half-reaction.
5.
At the risk of overwhelming the
students with detail, you can
mention that some reversible
reactions may require energy
input (from a battery) in order to
make them happen.
6.a)
Each zinc atom gives up two
electrons.
6.b)
Each silver ion accepts only one
electron.
6.c)
No, because one electron would
be left over.
6.d)
Two silver ions would be
required to accept two electrons
from a zinc atom.
• C
opper is not very reactive,
so it holds up over the years
when used in circuits.
• Gold, silver and platinum
are very unreactive, so
jewelry made with them
endures for generations.
• Aluminum reacts fairly
easily with other chemicals
and the dull finish is the
product of a reaction with
oxygen in air, Al2O3.
• Zinc is more reactive than
copper, but cheaper in cost.
2.
The most active metals are
at the top of the list (lithium,
potassium) and the least active
metals (platinum, gold) are at
the bottom.
3.
This a crucial point and
because of the many
misconceptions that some
students may have, it should be
emphasized. Metals and salts
of metals have vastly different
physical and chemical properties.
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695
chemical dominoes
7.a)
In nature, the reaction that
actually occurs is Ag+ and Zn.
This is because zinc metal is
more reactive than silver metal
(see the activity series).
8.a)
In nature, the reaction that
actually occurs is Al and Cu2+.
This is because aluminum metal
is more reactive than copper
metal (see the activity series).
Part C:
Investigating the Activity
Series of Metals
1.a)
The 21 pairs are:
Al3+ / Cu
Cu2+ / Al
Al3+ / Fe
Fe2+ / Al
Al3+ / Mg
Mg2+ / Al
Al3+ / Ag
Ag+ / Al
Al3+ / Zn
Zn2+ / Al
*
H+ / Al
Cu2+ / Fe
Fe2+ / Cu
Cu2+ / Mg
Mg2+ / Cu
Cu2+ / Ag
Ag+ / Cu
Cu2+ / Zn
Zn2+ / Cu
*
H+ / Cu
Fe2+ / Mg
Mg2+ / Fe
Fe2+ / Ag
Ag+ / Fe
Fe2+ / Zn
Zn2+ / Fe
*
H+ / Fe
Mg2+ / Ag
Ag+ / Mg
Mg2+ / Zn
Zn2+ / Mg
*
H+ / Mg
Ag+ / Zn
Zn2+ / Ag
*
H+ / Ag
*
H+ / Zn
Or, viewed in a different way:
696
Neutral metals
Al
Cu
Fe
Mg
Ag
Zn
(H2)
Al3+
Cu2+
X
1
2
3
4
5
—
1
X
7
8
9
10
—
Metal ions in solution
Fe2+
Mg2+
2
3
7
8
X
12
12
X
13
16
14
17
—
—
Ag+
4
9
13
16
X
19
—
Zn2+
5
10
14
17
19
X
—
H+
6
11
15
18
20
21
X
SEction 4 The METAL ACTIVITY SERIES: What can destroy a metal?
1.c)
Activity Series: Possible Reactions
→ 3Cu2+ + 2Al
2Al3+ + 3Fe → 3Fe2+ + Al
2Al3+ + 3Mg → 3Mg2+ + 2Al
Al3+ + 3Ag → 3 Ag+ + Al
2Al3+ + 3Zn → 3Zn2+ + 2Al
2Al3+ + 3H2 → 6H+ + 2Al
Cu2+ + Fe → Fe2+ + Cu
Cu2+ + Mg → Mg2+ + Cu
Cu2+ + 2Ag → 2Ag+ + Cu
Cu2+ + Zn → Zn2+ + Cu
Cu2+ + H2 → 2H+ + Cu
Fe2+ + Mg → Mg2+ + Fe
Fe2+ + 2Ag → 2Ag+ + Fe
Fe2+ + Zn → Zn2+ + Fe
Fe2+ + H2 → 2H+ + Fe
Mg2+ + 2Ag → 2Ag+ + Mg
Mg2+ + Zn → Zn2+ + Mg
Mg2+ + H2 → 2H+ + Mg
2Ag+ + Zn → Zn2+ + 2Ag
2Ag+ + H2 → 2H+ + 2Ag
Zn2+ + H2 → 2H+ + Zn
2Al3+ + 3Cu
or
or
or
or
or
or
or
or
or
or
or
or
or
or
or
or
or
or
or
or
or
→ 2Al3+ + 3Cu
3Fe2+ + 2Al → 2Al3+ + 3Fe
3Mg2+ + 2Al → 2Al3+ + 3Mg
3Ag+ + Al → Al3+ + 3Ag
3Zn2+ + 2Al → 2Al3+ + 3Zn
6H+ + 2Al → 2Al3+ + 3H2
Fe2+ + Cu → Cu2+ + Fe
Mg2+ + Cu → Cu2+ + Mg
2Ag+ + Cu → Cu2+ + 2Ag
Zn2+ + Cu → Cu2+ + Zn
2H+ + Cu → Cu2+ + H2
Mg2+ + Fe → Fe2+ + Mg
2Ag+ + Fe → Fe2+ + 2Ag
Zn2+ + Fe → Fe2+ + Zn
2H+ + Fe → Fe2+ + H2
2Ag+ + Mg → Mg2+ + 2Ag
Zn2+ + Mg → Mg2+ + Zn
2H+ + Mg → Mg2+ + H2
Zn2+ + 2Ag → 2Ag+ + Zn
2H+ + 2Ag → 2Ag+ + H2
2H+ + Zn → Zn2+ + H2
3Cu2+ + 2Al
4-4a
Blackline Master
4-4b
Blackline Master
chapter 4
Students should write both
balanced equations for each pair.
The equations are:
1.d)
The equations representing
reactions that should occur
according to the activity series
are highlighted in the preceding
table. The logic behind each
prediction is based on the
premise: X is more reactive than
Y, so the reaction X metal with
+
cation Y should proceed, while
the opposite should not.
2.
Explanations for findings should
follow the same logic as outlined
in 1.d) above.
This table is provided as a
Blackline Master on the
Teacher Resources CD.
A blank data table for students
to complete this exercise is also
provided as a Blackline Master
on the Teacher Resources CD.
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697
chemical dominoes
3.a)
White paper placed underneath
the wells should work best for
visualization. The reactions
observed should occur according
to the table of Possible
Interactions provided.
3.b)
Reactions should turn out
according to the predictions.
However, some reactions may be
slow and easy to miss. If you are
able to wait one day to observe
findings, there may be more
accuracy in the results. Also,
if the spot plates are made of
transparent plastic, viewing from
underneath can be helpful.
Notes
698
You may wish to hand out a copy
of the table of predicted reactions
for students to check off which
reactions occurred. (Blackline
Master 4-4a, noted on previous
page). Or, you may choose to let
students try writing their own
balanced equations first before
handing out the table.
5.
Students should be able to
identify all three of the mystery
chemicals now. One solution
was blue – which must be Cu2+.
One solution produced a gas
and the only possibility is a
+
solution containing H (acid).
One solution was the only one
to produce brown precipitate
+
Most reactions should match
(AgCl), so it must contain Ag .
predictions. Those that are very
It was also stored in a brown
slow or are easy to miss might be bottle, which is characteristic of
overlooked. Discuss reasons why silver solutions.
observations might not match
+
predictions.
The Ag solution should
have been the best and quickest
at destroying Mg. This makes
sense because of the three
mystery chemicals, Ag is the
farthest away from Mg in the
activity series.
SEction 4 The METAL ACTIVITY SERIES: What can destroy a metal?
chapter 4
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chemical dominoes
Chem Talk
This section elaborates on the
definition of metal and the
historical development of the
activity series. Examples of
redox reactions including those
involving zinc and copper are
highlighted. It concludes with
a discussion of the unique
properties of hydrogen.
700
SEction 4 The METAL ACTIVITY SERIES: What can destroy a metal?
chapter 4
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701
chemical dominoes
Checking Up
1.
Possible answers include:
• S hiny – used for jewelry
and other ornaments
• C
onducts electricity – used
for electrical circuits
• C
onducts heat – used in
cookware
• M
elts at high temperatures
– used for structures
• M
alleable – used to make
nails, boxes, flat surfaces,
sheet metal
2.
In rocks and minerals and in
the ocean; in the form of ions,
usually oxides.
3.
The activity series grew out of
alchemy through the careful
study of the properties of metals.
4.a)
Oxidation is the loss of electrons.
4.b)
Reduction is the gain of electrons.
4.c)
Both processes have to occur
simultaneously.
5.
that composes the circuit. The
identity of the metal wire affects
which metal ions will be less
active than the circuit. Students
might note that none of the
What Do You Think Now? metal ions on the list could
destroy a circuit made of gold.
In order to destroy a metal
wire with a chemical, students
Not all metals would be
should indicate that they would
destroyed equally well. Reactivity
use a chemical containing a less
would be dictated by the activity
active metal ion than the metal
series for metals. Each metal
The activity series can predict
which direction a redox reaction
will go.
702
would have its own reactivity.
Gold, for instance, would not be
expected to react at all.
You may want to share the
answer provided in A Chemist’s
Response and discuss the
implications. You may also want
to return to the illustration to
see if students are able to
identify the major components
of their circuit.
SEction 4 The METAL ACTIVITY SERIES: What can destroy a metal?
chapter 4
Why do you believe?
Students might indicate situations
in which metal ions are desired
such as those mentioned in the
Investigate section of the activity.
Other answers are also possible,
such as explosives or fireworks.
Why should you care?
Chem Essential
Questions
MACRO — Observations that
students might list include color
change of solution, production
of gas (bubbles), and/or formation
of new solids on the surface of
the metal pieces.
NANO — Metal atoms (and
hydrogen) are more stable when
they have given up a certain
number of electrons. Some
atoms are better able to give
up electrons than others. This is
referred to as being more active.
If metal A is more active than
metal B and atoms of metal A
come in contact with ions of metal
B+, electrons will be transferred from
atom A to ion B+. This will produce
ion A+ and atom B.
SYMBOLIC — Students should
provide a correctly written set of
half-reactions that could combine
to give a redox equation.
How do you know?
This will depend on students’
results. It is unlikely that their
results will be identical to the
activity series due to the slow
speed and low spontaneity of
some of the reactions.
Students should indicate that
circuits could be made from any
metal although if it is very inactive
(like gold), it might be difficult to
destroy because it won’t react with
most chemicals. To be effective as a
circuit, the metal should be stable,
inexpensive, readily available, and
non-toxic. Aluminum and zinc
are good examples. If you plan to
destroy the circuit, the metal
should be active, inexpensive,
readily available, and non-toxic.
Again, aluminum and zinc are
good choices.
The drawing should indicate two
paths for the electricity to flow.
When the path with less resistance
is destroyed, the alternate path
will conduct electricity. Depending
on your students’ familiarity with
circuitry, they may require guidance
with designing such a circuit.
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chemical dominoes
Reflecting on the
Section and the
Challenge
Students should read this
section for a specific, direct
connection between the section
and the Chapter Challenge.
While students do not answer
any questions in this section,
it will provide them with
valuable direction in the Chapter
Challenge. You may want to
provide some class time for
students to read this paragraph
silently or aloud.
Chem to Go
1.
The metals are: Na, Ni, Co, Sn,
Cd, Sr. (The nonmetals are:
Cl, Ar, P.)
2.a)
The metals tested that are more
reactive than hydrogen are Mg,
Al, Zn and Fe.
2.b)
The metals tested that are less
reactive than hydrogen are Cu
and Ag.
3.
a) A
l
c) A
l
Also, balanced redox reactions
with oxidation to Fe3+ are
possible, since gold ions are
strong oxidizing agents.
4.
c) 3Sn2+ + 2Al
b) Fe
Students should give one of the
following reactions:
a)
3
Pb2+
+ 2Al
→
2Al3+
+ 3Pb
b) 2Au+ + Fe → Fe2+ + 2Au
or
2Au3+ + 3Fe → 3Fe2+ + 2Au
704
→
2Al3+ + 3Sn
5.
The aluminum reacted and went
into solution as the Al3+ ion. A
chemical reaction occurred. The
aluminum did not change from
solid to liquid (this could only
occur at very high temperatures).
6.
Zn is more reactive than Fe, so
Zn will react preferentially with
any intruders. Should any Fe
react, the presence of Zn will
cause the Fe to re-form:
Fe2+ + Zn
→
7.
c) the same.
Zn2+ + Fe
SEction 4 The METAL ACTIVITY SERIES: What can destroy a metal?
chapter 4
Inquiring Further
Reacting metals with bases
The reaction of Al with
hydroxides is vigorous and
presents another way to
“destroy” a circuit. The
reaction is (assuming
potassium hydroxide):
2Al(s) + 2KOH(aq) + 6H2O
2KAl(OH)4(aq) + 3H2(g)
→
Extreme caution must be
used with caustic bases such
as NaOH and KOH. Metals
other than Al also react well
with hydroxides. Before
students begin to experiment,
check to see that their
experiment plans are safe.
Student procedures will likely
be similar to that in Part B
of the Investigate. They
would substitute a solution
of a strong base for the various
metal cations.
8.
b) Zn(s)
→
Zn2+(aq) + 2e–
9.
d) loses an electron
10.
a) e lectrons from Al to Cr3+
11.
a) m
agnesium
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chemical dominoes
Section 4 – QUiz
SECTION 4 – QUIZ ANSWERS
4-4c
Blackline Master
1. If you have a balanced ionic equation,
the total charge of the products should be:
a) greater than the total charge of
the reactants
b) equal to the total charge of
the reactants
c) less than the total charge of
the reactants
❶
b) e qual to the total charge of
the reactants
❷
❸
c) calcium
b) Answers will likely focus on the
non-reactivity of copper since
magnesium is not a typical
household chemical. Electrical
circuits are made of Cu, which is
fairly unreactive. Other possibilities
include that copper cooking pots
are frequently found in the kitchen.
❹
a) Ag
2. Which metal reacts spontaneously with
a solution containing magnesium ions?
a) aluminum
b) gold
c) calcium
d) hydrogen
3. In the lab, you saw that all the acid reacted
with magnesium, while the acid did not react
with copper.
a) Why does copper behave differently
than magnesium?
b) Describe a practical everyday
household use for this information.
4. Silver loses one electron when it oxidizes.
Iron can lose three electrons when it oxidizes.
706
a) Write the oxidation half-reactions for
each metal.
b) Write the balanced chemical reaction that
occurs naturally, according to the activity
series, when one of the metals is placed in
a solution containing the positive ion of
the other metal.
a) Mg is more active than H and can
reduce it, Cu is less active than H
and cannot displace it.
→ Ag+ + 1 eFe → Fe3+ + 3eb) Fe + 3Ag+ → Fe3+ + 3Ag
SEction 4 The METAL ACTIVITY SERIES: What can destroy a metal?
chapter 4
Notes
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