Download Activity Series Unit

The Activity Series
Minnesota K-12 Academic Standards that are covered in this Unit:
9.1.1.2.3 – The Nature of Science and Engineering: The practice of science
- Scientific inquiry uses multiple interrelated processes to investigate
and explain the natural world.
o Identify the critical assumptions and logic used in a line of
reasoning to judge the validity of a claim
9.2.1.2.3 – Physical Science: Matter
- Chemical reactions involve the rearrangement of atoms as chemical
bonds are broken and formed through transferring or sharing of
electrons and the absorption or release of energy.
o Describe a chemical reaction using words and symbolic
equations.
9C.2.1.3.2 – Physical Science: Matter
- Chemical reactions describe a chemical change in which one or more
reactants are transformed into one or more products.
o Use solubility and activity of ions to determine whether a
double replacement or single replacement reaction will occur.
9C.2.1.3.4 – Physical Science: Matter
- Chemical reactions describe a chemical change in which one or more
reactants are transformed into one or more products.
o Balance chemical equations by applying the laws of
conservation of mass and constant composition.
Pre-Requisites for this lesson?
Matter
The Mole
Periodic Trends
Nomenclature
Chemical Reactions and Equations
Stoichiometry
The Activity Series
Engage: How do Metals React with Other Substances?
A. You’ve learned about characteristics of metals and nonmetals, and balancing
chemical equations. How does one determine the relative reactivities of metals?
In spontaneous redox reactions, which metal is more active?
Before Studying this Unit
After Studying this Unit
Explore 1: What reaction occurs between metals and nonmetals?
1. To each of 6 test tubes containing 0.5 mL of dilute 6 M HCl, add a small
piece of the metals Ca, Cu, Fe, Mg, Sn, and Zn.
2. a. Observe the test tubes and note any changes that occur (such as the
evolution of a gas, whether it is vigorous or not, and any other changes). Record
your observations below.
Results may vary
- Ca: Metal dissolved with very vigorous evolution of H2 - a colorless
solution remained.
- Cu: No Reaction
- Mg: Metal dissolved, but occurred slower than Ca – some H2 was
formed and a colorless solution was produced
- Fe: Very slow reaction – only some H2 was formed
- Sn: Extremely slow reaction – very little H2 formed
- Zn: Metal dissolved to produce a colorless solution – some H2was
formed
b. Perform a splint test on the gas(es) formed. From the results of this
investigation what gas(es) were formed?
Splint test should give a slight pop, this means hydrogen was formed.
3. After completing each reaction and recording your observations, dispose of
the contents of your test tubes in the designated containers.
Explain 1: Why do some ions interact with other in a reaction while others do not?
4. Consider the reactions that took place between the metals and hydrochloric acid.
Write a chemical equation of each reaction that took place. (Be sure to write each species
as it exists in solution) This is called the total ionic equation.
Ca(s) + 2HCl(aq) → CaCl2(aq) + H2(g)
Ca(s) + 2H+(aq) + 2Cl-(aq) → Ca2+(aq) + 2Cl-(aq) + H2(g)
No reaction for Copper and HCl
Mg(s) + 2HCl(aq) → MgCl2(aq) + H2(g)
Mg(s) + 2H+(aq) + 2Cl-(aq) → Mg2+(aq) + 2Cl-(aq) + H2(g)
Fe(s) + 2HCl(aq) → FeCl2(aq) + H2(g)
Fe(s) + 2H+(aq) + 2Cl-(aq) → Fe2+(aq) + 2Cl-(aq) + H2(g)
Sn(s) + 2HCl(aq) → SnCl2(aq) + H2(g)
Sn(s) + 2H+(aq) + 2Cl-(aq) → Sn2+(aq) + 2Cl-(aq) + H2(g)
Zn(s) + 2HCl(aq) → ZnCl2(aq) + H2(g)
Zn(s) + 2H+(aq) + 2Cl-(aq) → Zn2+(aq) + 2Cl-(aq) + H2(g)
5. When analyzing the reactions that took place, which ion(s) did not play an
“active” role in the reaction?
Cl6. Using your answers from Problems 4 and 5, re-write this ionic equation, but
only include the species that play an active role in the change that occurs. Do this
for the reaction that occurs between magnesium and hydrochloric acid.
Ca(s) + 2H+(aq) → Ca2+(aq) + H2(g)
Mg(s) + 2H+(aq) → Mg2+(aq) + H2(g)
Fe(s) + 2H+(aq) →Fe2+(aq) + H2(g)
Sn(s) + 2H+(aq) → Sn2+(aq) + H2(g)
Zn(s) + 2H+(aq) → Zn2+(aq) + H2(g)
7.
When these species are omitted, the net ionic equation is generated.
8. The ions that do not play a part in the reaction are called spectator ions.
These specie(s) can be removed from the total ionic equation to form the net ionic
equation (as you’ve done previously). If these ions do not play an active part
during the reaction, what is the purpose of them?
Spectator Ion – atom or group of atoms with a charge that is in both the
reactant and product side of the equation.
These ions are in solution to balance the charge of the solution.
Elaborate 1: How do metals react with acids?
9. Write the complete equation, the total ionic equation, the spectator ion(s), and
the net ionic equation for each of the following reactions:
a. Tin and hydrochloric acid
Sn(s) + 2HCl(aq) → SnCl2(aq) + H2(g)
Sn(s) + 2H+(aq) + 2Cl-(aq) → Sn2+(aq) + 2Cl-(aq) + H2(g)
ClSn(s) + 2H+(aq) → Sn2+(aq) + H2(g)
b. Magnesium and sulfuric acid
Mg(s) + H2SO4(aq) → MgSO4 + H2(g)
Mg(s) + 2H+(aq) + SO42-(aq) → Mg2+(aq) + SO42-(aq) + H2(g)
SO42Mg(s) + 2H+(aq) → Mg2+(aq) + H2(g)
c. Calcium and acetic (ethanoic) acid
Ca(s) + 2CH3COOH(aq) → Ca(CH3COO)2(aq) + H2(g)
Ca(s) + 2CH3COO-(aq) + 2H+(aq) → Ca2+(aq) + 2CH3COO-(aq) +
H2(g)
CH3COOCa(s) + 2H+(aq) → Ca2+(aq) + H2(g)
d. Aluminum and Sulfuric Acid
2Al(s) + 3H2SO4(aq) → Al2(SO4)3(aq) + 3H2(g)
2Al(s) + 6H+(aq) + 3SO42-(aq) → 2Al3+(aq) + 3(SO4)2-(aq) + 3H2(g)
SO422Al(s) + 6H+(aq) → 2Al3+(aq) + 3H2(g)
10. From the knowledge you have gained thus far, what are the general products
of a reaction between a metal and an acid?
Hydrogen gas & metal ions in solution (other ion is a spectator)
Explore 2: What reactions occur between metals and solutions of metal ions?
11. (Work in pairs or groups for this phase.) Add a small piece of calcium metal
to each of 7 test tubes containing, respectively, about 0.5 mL of Ca(NO ) , CuSO ,
FeSO , Fe(NO ) , Mg(NO ) , SnCl , and Zn(NO ) solutions. Note any reaction that
occurs by observing whether a color change occurs on the surface of the metal or
in the solution or whether a gas is evolved. Pay close attention to whether not the
Ca dissolves (because there may be other side reactions). Record your
observations in the calcium row of the table provided.
Ca2+
Cu2+
Fe3+
Fe2+
Mg2+
Sn4+
Zn2+
3 2
4
3 3
3 2
2
4
3 2
Metal
ions
Metal
Soln changed
from blue to
green as Ca
dissolved
Ca
N.R.
Ca dissolved
to give dirty
green soln
Ca dissolved to
give a dirty green
soln and a brn-red
precipitate
Ca dissolved –
milky white
suspension
formed
Ca dissolved; Ca dissolved
milky soln
white ppt
formed
formed
N.R.
N.R.
N.R.
N.R.
N.R.
N.R.
Slow reaction
N.R.
Mg dissolved
and dark gray
ppt formed
Mg
Dissolved
Cu
N.R.
Very slow
reaction
N.R
Soln was
darker and
black ppt
formed on Mg
Mg slowly
dissolved as
rust ppts
Mg dissolved very
slowly
N.R
Very slow
reaction
N.R.
N.R
N.R.
Zn dissolved
slowly and a
brn-red ppt
formed
Almost no reaction
N.R.
Fe
Mg
Sn
N.R. Zn turned black
Zn
Very little
reaction
Zn dissolved
slowly
12. After completing each series of reactions dispose of the contents of your test
tubes in the designated containers.
13. Write both complete and net ionic equations for the reaction that occurs
between calcium metal and tin(IV) tetrachloride.
2Ca(s) + SnCl4(aq) → 2CaCl2(aq) + Sn(s)
2Ca(s) + Sn4+(aq) → 2Ca2+(aq) + Sn(s)
14. Repeat the preceding process by adding a small piece of copper to another 0.5
mL of each of the metal-cation solutions. Record your observations on the table
provided above.
15. Do the same for iron, magnesium, tin, and zinc, and record all your
observations in the table provided previously. Remember to pay close attention to
whether or not the metal dissolves.
16. Write both complete molecular and ionic equations for 6 reactions that took
place in this activity. (For some, there is no reaction that takes place).
a. Calcium and Mg(NO )
Ca(s) + Mg(NO3)2(aq) → Ca(NO3)2(aq) + Mg(s)
Ca(s) + Mg2+(aq) → Ca2+(aq) + Mg(s)
3 2
b. Copper and SnCl
No Reaction
4
c. Magnesium and Zn(NO )
Mg(s) + Zn(NO3)2(aq) → Mg(NO3)2(aq) + Zn(s)
Mg(s) + Zn2+(aq) → Mg2+(aq) + Zn(s)
3 2
d. Zinc and Fe(NO )
3Zn(s) + 2Fe(NO3)3(aq) → 3Zn(NO3)3(aq) + 2Fe(s)
3Zn(s) + 2Fe3+(aq) → 3Zn2+(aq) + 2Fe(s)
3 3
e. Iron and SnCl
2Fe(s) + SnCl4(aq) → 2FeCl2(aq) + Sn(s)
2Fe(s) + Sn4+(aq) → 2Fe2+(aq) + Sn
4
f. Tin and CuSO
Sn(s) + CuSO4(aq) → SnSO4(aq) + Cu(s)
Sn(s) + Cu2+(aq) → Sn2+(aq) + Cu(s)
4
Explain 2: Oxidation and Reduction Reactions
Refer to these three reactions as needed in the following questions.
I. 4Fe + 3O2 → 2Fe2O3
II. 2Cu + O2 → 2CuO
III. 2Mg + O2 → 2MgO
17. What happens to the metals in each of these reactions? Do you see any
trends?
You may hear answers like the metal is binding to oxygen, the metal and oxygen
are bonding, etc. The trend is that each metal coordinates with oxygen.
18. There are certain rules that need to be followed when going through the
reaction process. (More specifically for assigning oxidation states)
a. free atom has an oxidation number of zero. It is not sharing, gaining, or
losing electrons.
b. Polyatomic elements have an oxidation number of zero for each atom.
Elements such as H2, O2, and P4 share electrons equally among all atoms in the
molecule.
c. The sum of the oxidation numbers in a neutral molecule is zero.
d. A monatomic ion has an oxidation number equal to its charge. For example,
the oxidation number of the oxygen in the oxide ion, O2-, is –2.
e. The sum of the oxidation numbers in a polyatomic ion is equal to the charge
on the ion.
f. The oxidation number for an oxygen ion is –2. The oxidation number for a
hydrogen ion is +1. (These are general rules that apply to a wide variety of
reactions. We will not consider the exceptions to this rule in this activity).
19. Using the reactions given as references at the beginning of this phase:
a. What happened to the oxidation number of iron in the first reaction?
0 → 3+
b. What happened to the oxidation number of copper in the second reaction?
0 → 1+
c. What happened to the oxidation number of magnesium in the third
reaction?
0 → 2+
20. Do you notice a trend or pattern for the changes in oxidation numbers from
the previous problem?
The pattern is that the oxidation numbers of the metals increase.
21. Let’s called this trend oxidation. Use the following equation as another
example to consider this term. What happens to iron in reaction four? Is there a
general trend that can be stated for any species that is oxidized?
IV. 2 Fe + 3 Cl2 → 2 FeCl3
Oxidation number goes from 0 to 3+ (Number is still increasing-becoming more
positive; just as before.)
22. Now, return your attention to the original three reactions given as a reference
in this phase.
a. What happened to the oxidation number of oxygen in the first reaction?
0→2
b. What happened to the oxidation number of oxygen in the second reaction?
0 → 2c. What happened to the oxidation number of oxygen in the third reaction?
0 → 2-
23. Do you notice a trend or pattern for the changes in oxidation numbers from
the previous problem?
The oxidation numbers decrease. This term is known as reduction.
24. Let’s called this trend reduction. Use the following equation as another
example to consider this term. What happens to chlorine in reaction four? Is
there a general trend that can be stated for any species that is reduced?
IV. 2Fe + 3Cl2 → 2FeCl3
Oxidation number goes from 0 to 1- (Number is still decreasing-becoming more
negative; just as before.)
25. Look at all four reactions at once. (They are reprinted below for your
convenience). Count the number of valence electrons in each of the metal atoms
and ions before and after the reactions. You will need to make a table to organize
your results.
I. 4Fe + 3O2 → 2Fe2O3
II. 2Cu + O2 → 2CuO
III. 2Mg + O2 → 2MgO
IV. 2Fe + 3Cl2 → 2FeCl3
Reaction
Metal Electrons Before
Metal Electrons After
I
4x8=32e
4(8-3)=20eII
2x11=22e-
2(11-1)=20e-
III
2x2=4e-
2x0=0e-
IV
2x8=16e-
2x6=12e-
26. Generally speaking, what happens to the metals in these reactions?
The metals are losing electrons in these reactions.
27. When looking at electrons, what can be said about the term, oxidation?
When a species is oxidized, it loses electrons.
28. Count the number of electrons in each of the oxygens (and the chloride)
before and after the reactions. You will need to make a table to organize your
results.
Reaction O or Cl Electrons Before O or Cl Electrons After
I
6x6=36e6(6+2)=48eII
2x6=12e-
2(6+2)=16e-
III
2x612e-
2(6+2)=16e-
IV
6x7=42e-
6(7+1)=48e-
29. Generally speaking, what happens to these species in these reactions?
They are gaining electrons from the metals.
30. When looking at electrons, what can be said about the term, reduction?
Reduction means the gain of electrons.
31. In these reactions what species is causing reduction? What species cause
oxidation?
The species can be called the reducing agent – reduces the other species
The species can be called the oxidizing agent – oxidizes the other species
32. For the future, we will call the species that causes reduction, the reducing
agent and the species that causes oxidation, the oxidizing agent.
33. Consider the following reactions:
a. S + O2 → SO2
b. Cl2 + I- → I2 + Clc. AgNO3 + Cu → Ag + CuNO3
d. NaOH + HCl → NaCl + H2OFe(NO3)3 + NaOH → Fe(OH)3 + NaNO3
e. WO2 + H2 → W + H2O
f. 3 ZnS + 8 HNO3 → 3 ZnSO4 + 8 NO + 4 H2O
Are all of these oxidation-reduction (redox) reactions?
No.
34. Which reactions are redox reactions? How can you tell?
Change in oxidation numbers
a. 0
0
+4 –2
1/8S8 + O2 → SO2
Yes, it’s redox.
S atoms in S8 are oxidized, O atoms in O2 are reduced, O2 is the oxidizing
agent, and S8 is the reducing agent.
b. 0 -1 0 -1
Cl2 + I- → I2 + ClYes, it's redox.
The I- are oxidized, the Cl2 are reduced, Cl2 is the oxidizing agent, and I- is
the reducing agent.
c. +1 -1
0
0
+1 -1
AgNO3 + Cu → Ag + CuNO3
Yes, it's redox.
The Cu is oxidized, the Ag in AgNO3 is reduced, the AgNO3 is the oxidizing
agent, and Cu is the reducing agent.
d. +1 -1
+1 -1 +1 -1
+3 -1
+1 -1
+3 -1
+1 -1
NaOH + HCl → NaCl + H2OFe(NO3)3 + NaOH → Fe(OH)3 + NaNO3
Not redox
e. +4-2
0
0
+1 -2
WO2 + H2 → W + H2O
Yes, it's redox.
The H is oxidzed, the W in WO2 is reduced, the WO2 is the oxidizing agent,
and the H2 is the reducing agent.
f. +2 -2
+1 -1
+2 -2
3 ZnS + 8 HNO3 → 3 ZnSO4 + 8 NO + 4 H2O
Not redox.
35. For each of these reactions (above), label the oxidation and reducing agents.
Also label the species that is being oxidized and the species that is being reduced.
You may want to make a table to organize your thoughts.
Answers may vary
36. From the previous question, is there a relationship between which species are
oxidizing agents and what is occurring to them?
Oxidizing agents are being reduced.
37. Is there a relationship between which species are reducing agents and what is
occurring to them?
Reducing agents are being oxidized.
38. From what you have discovered in this section, why must oxidation be
accompanied by reduction?
An oxidation must always be accompanied by a reduction reaction because the
species being oxidized must transfer an electron to some other species that is
reduced. The electron cannot just be given up to free space.
Elaborate 2: How can you Determine the Relative Activities of Species Studied?
39. From the information contained in the table you constructed in the explore 2
activity and the knowledge you have attained from the explain section, you can
rank these six metals according to their relative chemical reactivities. List the
metals in terms of decreasing reactivity, starting with the most reactive (1) and
ending with the least reactive (6).
1. __Ca__ 2. __Mg__
3. __Zn__
4. __Fe__
5. __Sn__
6. __Cu__
40. How can this trend be explained?
Reactivity trends in the reaction table. Ca reacted with everything and Cu did not
react with anything.
41. Given the trend that was found previously, which of these six metals should
be the most reactive toward oxygen? Why?
Calcium
After Questions 39-41 have been answered by the class, a discussion may be
helpful. Be sure that everyone was able to deduce the correct answers. Point the
students to the activity series listed in page 2 of the data pack.
42. Sodium is slightly less reactive than calcium. Predict the outcome of the
following reactions. (For some, no reaction will take place)
The following questions should be answered while using the previously
determined activity series.
a. Na + H2O →2Na + 2H2O 2NaOH + H2
b. Na + O2 → 4Na + O2 2Na2O
c. Na + HCl → 2Na + 2HCl 2NaCl + H2
d. Na + Ca2+ → No Reaction
43. For each of the following reactions, indicate which substance is oxidized and
which is reduced. Which substance is the oxidizing agent and which is the
reducing agent?
a. 2Al + 3Cl2 → 2AlCl3
Oxidized Reduced Oxidizing Agent Reducing Agent
Al
Cl2
Cl2
Al
+
2+
3+
2+
b. 8H + MnO4 + 5Fe → 5Fe + Mn + 4H2O
Oxidized Reduced Oxidizing Agent Reducing Agent
Fe2+
MnO4Fe2+
MnO4c. FeS + 3NO3- + 4H+ → 3NO + SO42- + Fe3+ + 2H2O
Oxidized Reduced Oxidizing Agent Reducing Agent
FeS
NO3NO3FeS
d. Zn + 2HCl → ZnCl2 + 2H2
Oxidized Reduced Oxidizing Agent Reducing Agent
Zn
HCl
HCl
Zn
44. Using the information gathered in this unit determine the relative activities of
the metals. If the following redox reactions are found to occur spontaneously,
identify the more active metal in each reaction.
a. 2Li + Cu2+ → 2Li+ + Cu
Li
b. Cr + 3V3+ → 3V2+ + Cr3+
V
c. Cd + 2Ti3+ → 2Ti2+ + Cd2+
Ti