Types of Chemical Reactions and Equations
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Objectives: You will be able to identify a reaction according to the chemical changes that occur. You will be
able to write balanced chemical equations given the reactants. You will be able to write complete ionic
equations and net ionic equations when appropriate. Finally you will be able to identify the element oxidized,
the element reduced, the oxidizing reactant and the reducing reactant for oxidation-reduction reactions where
appropriate.
Background: Chemical reactions lead to changes in the composition of matter at the atomic and molecular
level. Chemical changes are typically signaled by one or more of the following observations: gas evolution (p.
161 in Tro 2nd Edition, color changes, formation of a solid precipitate (Chapter 4 Section 6 in Tro 2nd Edition),
release or absorption of energy as light or heat (Chapter 6 Section 6 in Tro 2nd Edition).
Reactions may be described in many ways. One way is to describe the reaction by what physically occurs
during the reaction. Four terms used are Double Displacement, Single Displacement, Synthesis, and
Decomposition. Yet another way is to describe the reaction according to the chemical process it undergoes.
Three discussed in lecture include Precipitation, Acid / Base Neutralization and Oxidation-Reduction
reactions. We will focus on describing the reaction according to what physically occurs during the reaction in
this laboratory but references are made to the way by which these two reaction types overlap in the background
reading.
Synthesis reactions are the result of more than one reactant combining to form a single product. In the
diagram below the reactant bird is wasting away. When the bird eats the worm, hence combines with the worm,
you get a fat bird. The product bird is therefore the result of the skinny bird combining with the worm.
Synthesis reactions are often a type of redox reaction.
The most common way to illustrate this type of reaction
is via the general reaction:
A + B à AB
Examples of four different combination reactions
involving elements are shown on p. 162 of Tro 2nd
Edition.
Decomposition reactions occur when one reactant gives off two or more products. In the diagram below the
reactant egg hatches to give two products, the egg shell plus turtle.
This reaction type can be
generalized as:
AB à A + B
Examples of decomposition
reactions can be seen on p. 219
for the depletion of ozone and
260 for the decomposition of methane in Tro 2nd Edition.
Double Displacement reactions, or metathesis reactions, are ones in which the cation swaps the anion it is
attached to. In the picture below the man is the cation and the hat the
anion. Notice that as the reaction proceeds from left to right, the
cation simply swapped its anion.
Chemical Reactions and Equations p.1
Acid / Base Neutralization reactions (Chapter 4 Section 8 in Tro 2nd Edition) and Precipitation reactions
(Chapter 4 Section 6 in Tro 2nd Edition) involving two ionic solutions fall into this category. More commonly
this reaction is illustrated using the general formula: AB + CD à AD + CB
Single Displacement reactions are reactions in which one elemental component of a compound gains or
loses electrons to become a neutral element while the neutral element in the reaction dissolves to form its
corresponding anion or cation. In the diagram below the man in the red shirt represents the neutral element on
the reactant side. As the reaction proceeds the man in the red shirt dissolves taking on the anion. As this occurs
the other male changes from being a cation as a reactant to being
neutral as a product.
Precipitation reactions in which a metal plates out of solution is an
example of a single displacement reaction. More commonly this
reaction is illustrated using the general formula:
A + BY à AY + B
One example of such reactions can be seen on p. 166
Example 4.17. Appendix II Table D is a useful tool for predicting
elements capable of displacing other less-active elements in such
reactions as will be discussed extensively in College Chemistry II,
Chapter 18.
In addition to being able to recognize the type of reaction according to the physical changes the reactants
undergo, in this lab you will practice writing balanced chemical equations by predicting the products formed
from the identity of the reactants.
Balanced Chemical Equation (Molecular Equation):
(1) Identify the type of reaction: Single Displacement, Double Displacement, etc.
(2) Double Displacement: (Example 4.10 and 4.11 in Tro 2nd Edition)
a. Identify the cation and anion in each of the reactants e.g.
CaCl2 – the cation is Ca2+ and the anion is Cl- not Cl2b.
c.
d.
e.
f.
Swap the anion partner as you move toward the products.
Determine the formula of the compounds.
Write the equation with reactants and products.
Determine the state of matter of the compounds.
Balance the equation.
Example:
Reactants:
a.
b.
c.
d.
e.
CaCl2
+
AgNO3
Ca2+ Cl+
Ag+ + NO32+
Ca NO3
+
Ag+ + ClCa(NO3)2
+
AgCl
CaCl2 + AgNO3 à Ca(NO3)2 + AgCl
Since these are ionic compounds you will use your solubility rules to determine the state
of matter. If acid/base neutralization reaction, then the acid and base should be aqueous
(aq), the water will be a liquid (l), and you will determine the state of matter for the ionic
salt using the solubility rules.
Chemical Reactions and Equations p.2
The compound will be
Soluble (aq)
Soluble (aq)
Soluble (aq)
Soluble (aq)
Insoluble (s)
Insoluble (s)
If it contains
Group 1, NH4+
NO3-, C2H3O2-, ClO4-, HCO3Halides (F-, Cl-, Br-, I-)
SO42O2-, OHAll other anions
Unless combined with
----------------------------Ag+, Pb2+, Hg22+
Ag+, Pb2+, Hg22+, Ca2+, Sr2+, Ba2+
Ca2+, Sr2+, Ba2+
----------------
CaCl2 (aq) + AgNO3 (aq) à Ca(NO3)2 (aq) + AgCl (s)
CaCl2 (aq) + 2 AgNO3 (aq) à Ca(NO3)2 (aq) + 2 AgCl (s)
f.
(3) Single Displacement Reactions
a. Identify the cation and anion of the ionic compound e.g.
HCl – the cation is H+ and the anion is Clb. Give the anion to the reactant that was a neutral element as you move toward the products allowing
the original cation to become a neutral element.
c. Determine the formula of the resultant ionic compound. (The reactant that was a neutral element will
typically take on its favorite charge.)
d. Write the equation with reactants and products.
e. Determine the state of matter of the compounds.
f. Balance the equation.
Example:
Reactants:
a.
b.
c.
d.
e.
f.
HCl
+
Al
+
H Cl
+
Al
H2
+
Al + ClH2
+
AlCl3 because the charge on an Al cation is +3
HCl + Al à H2 +
AlCl3
Use solubility rules if the compound is an ionic compound. Acids and bases are aqueous
(aq), pure elements will be either solid (s), liquid (l), or gas (g) depending on the element.
HCl (aq) + Al (s) à H2 (g) + AlCl3 (aq)
6 HCl (aq) + 2 Al (s) à 3 H2 (g) + 2 AlCl3 (aq)
(4) Synthesis and Decomposition reactions
There is less of a pattern. Some of the above information may help so keep it in mind. Common types of
decomposition reactions include: dehydration of hydrates, decomposition of carbonates into ionic oxides
and carbon dioxide gas.
Chemical Reactions and Equations p.3
This lab has also been designed to give you practice with the development of complete ionic equations and net
ionic equations where appropriate (Chapter 4 Section 7 in Tro 2nd Edition). A complete ionic equation shows
all substances as they truly appear in solution. To write a complete ionic equation it is essential that you be able
to identify substances as strong electrolytes, weak electrolytes, or non-electrolytes.
Strong Electrolytes – compounds that dissociate entirely into their ions in solution
Weak Electrolytes – compounds that only partially dissociate into their ions in solution
Non-electrolytes – compounds that do not dissociate in solution
Strong Electrolytes
Soluble Ionic Compounds (Solubility rules)
Strong Acids
Strong Bases
Weak Electrolytes
Insoluble Ionic Compounds (solubility rules)
Weak Acids
Weak Bases
Materials that are not strong or weak electrolytes are classified as non-electrolytes. Now, why is it essential that
you be able to classify compounds into one of these categories? In going from the balanced chemical equation
to the complete ionic equation, you only break apart the strong electrolytes into their ions.
Complete ionic equation:
(1) Write the balanced chemical equation.
(2) Determine the ions for all strong electrolytes. (Since they are ions include the charge.)
(3) Balance the equation so as to not lose any ions between steps 1 and 2.
Example:
(1) CaCl2 (aq) + 2 AgNO3 (aq) à Ca(NO3)2 (aq) + 2 AgCl (s)
(2) Ca2+ (aq) + Cl- (aq) + Ag+ (aq) + NO3- (aq) à Ca2+ (aq) + NO3- (aq) + AgCl (s)
(3) Ca2+ (aq) + 2 Cl- (aq) + 2 Ag+ (aq) + 2 NO3- (aq) à Ca2+ (aq) + 2 NO3- (aq) + 2 AgCl (s)
Example:
(1) 6 HCl (aq) + 2 Al (s) à 3 H2 (g) +
2 AlCl3 (aq)
+
(2) H (aq) + Cl (aq) + Al (s) à H2 (g) + Al3+ (aq) + Cl- (aq)
(3) 6 H+ (aq) + 6 Cl- (aq) + 2 Al (s) à 3 H2 (g) + 2 Al3+ (aq) + 6 Cl- (aq)
A Net Ionic Equation (See. Sec. 4.3) shows only the substances that undergo a change during the reaction
process. This can be a change in oxidation state or state of matter. To get from the complete ionic equation to
the net ionic equation, you will remove all spectator ions. Spectator Ions are ions that appear on both reactant
and product side of the equation. They have the same charge and the same state of matter.
Net Ionic Equation
(1) Begin with the complete ionic equation.
(2) Remove all spectator ions
Example:
(1) Ca2+ (aq) + 2 Cl- (aq) + 2 Ag+ (aq) + 2 NO3- (aq) à Ca2+ (aq) + 2 NO3- (aq) + 2 AgCl (s)
(2) 2 Cl- (aq) + 2 Ag+ (aq) à 2 AgCl (s)
Example:
(1) 6 H+ (aq) + 6 Cl- (aq) + 2 Al (s) à 3 H2 (g) + 2 Al3+ (aq) + 6 Cl- (aq)
(2) 6 H+ (aq) + 2 Al (s) à 3 H2 (g) + 2 Al3+ (aq)
Chemical Reactions and Equations p.4
Finally, this lab is designed to give the student practice with oxidation / reduction reactions. You will use rules
for determining the oxidation state of the elements and ions in a reaction. From that information you will be
able to identify the element oxidized, the element reduced, the oxidizing reactant and the reducing reactant.
Rules for Determining Oxidation States of Elements (Chapter 4 Section 9 in Tro 2nd Edition)
1. Elements in their standard state have an oxidation state of 0.
2. Ions of elements have an oxidation state equivalent to the ions charge.
3. Hydrogen has an oxidation state of +1 unless the only other element present in the compound is a metal,
in which case hydrogen has an oxidation state of -1.
4. Oxygen has an oxidation state of -2 unless it is part of the peroxide polyatomic ion, in which case
oxygen has a -1 oxidation state.
5. Halogens have a -1 oxidation state unless paired with oxygen or another halogen.
a. If paired with oxygen then the oxidation state of the halogen must be calculated
b. If paired with another halogen, then the smaller of the halogens will have a -1 oxidation state and
the oxidation state for the other halogen must be calculated.
6. The unknown oxidation state of an element is calculated the same as the charge on a transition metal in
an ionic compound with the calculation set equal to zero. The only difference occurs when the
compound is itself a charged ion, in which case the calculation is set equal to the charge.
The Element Oxidized is the element that becomes more
positive (losing electrons) as it moves from reactant to product
side of the equation.
The Element Reduced is the element that becomes more
negative (gaining electrons) as it moves from reactant to
product side of the equation.
The Oxidizing Reactant (reductant, reducing agent) is the reactant that contains the element that was reduced.
It is called the oxidizing reactant because of the effect it has on another element within the reaction; it causes
another element to be oxidized.
The Reducing Reactant (oxidant, oxidizing agent) is the reactant that contains the element that was oxidized.
It is called the reducing reactant because of the effect it has on another element within the reaction; it causes
another element to be reduced.
These four concepts can be illustrated with a single displacement reaction or with a synthesis reaction that starts
with its elements.
6HCl(aq) + 2Al(s) à 3H2(g) + 2AlCl3(aq)
+1 -­‐1
0
0
+3 -­‐1
e-­‐ lost, oxidation
e-­‐ gain, reduction
HCl is the oxidizing reactant because it caused Al to be oxidized.
Al is the reducing reactant because it caused H to be reduced.
8Fe(s) + S8(s) à 8FeS(s)
0
0
+2 -­‐2
e-­‐ gain, reduction
e-­‐ loss, oxidation
Fe is the reducing reactant because it caused S to be reduced.
S8 is the oxidizing reactant because it caused Fe to be oxidized.
Chemical Reactions and Equations p.5
Reactions: Procedures / Data and Analysis
I.
Reacting Mg with HCl Solution
1. Place a small piece (thumb nail length) of magnesium in a small test tube.
2. Cover the magnesium with 0.1M HCl
3. Dispose of mixture according to your instructor’s instructions.
a. Appearance of Mg:
b. Appearance of HCl solution:
c. Evidence that a chemical reaction occurred:
d. Molecular equation: (include states of matter)
e. Complete ionic equation: (include states of matter)
f. Net ionic equation: (include states of matter)
g. Element Oxidized
__________________
h. Element Reduced
__________________
i. Reducing Reactant
__________________
j. Oxidizing Reactant
__________________
k. Type of reaction:
Chemical Reactions and Equations p.6
II.
Reacting Pb(NO3)2 with KI Solution
1. Add 2 to 3 drops of lead (II) nitrate to a small test tube.
2. Add 2 to 3 drops of potassium iodide to the above test tube, note any changes.
3. Dispose of this mixture in the appropriately labeled recovery container.
a. Appearance of Pb(NO3)2 solution:
b. Appearance of KI solution:
c. Appearance of the reaction mixture:
d. Molecular equation: (include states of matter)
e. Complete ionic equation: (include states of matter)
f. Net ionic equation: (include states of matter)
g. Type of reaction:
III.
Heating Magnesium (Mg) and oxygen (O2)
1. Light a Bunsen burner
2. Hold a piece of magnesium (roughly thumbnail length) with a pair of tongs.
3. In the other hand hold a watch glass.
4. Place the magnesium in the flame. (Do not look directly at the magnesium.)
5. When the magnesium catches fire, place it over the watch glass.
*NOTE: This should be done in the hood so as to avoid setting off Fire Alarms. The instructor may
choose to demonstrate this reaction. To perform it in the hood, a propane torch can be used to light
the magnesium.*
a. Appearance of the unreacted Mg:
b. Appearance of the Mg after heating:
c. Molecular equation: (include states of matter)
Chemical Reactions and Equations p.7
d. Element Oxidized
__________________
e. Element Reduced
__________________
f. Reducing Reactant
__________________
g. Oxidizing Reactant
__________________
h. Type of reaction:
IV.
Heating CuSO4 . 5 H2O(s)
1. Light a Bunsen burner, adjust to obtain a moderate flame
2. Place an eraser size amount of copper (II) sulfate pentahydrate solid
into a test tube.
3. While holding the test tube at the top with a test-tube clamp, move
the tube though the flame so as to gently heat the solid. Note any
changes.
4. Dispose of this material according to your instructor’s directions.
a. Appearance of CuSO4 . 5 H2O(s):
b. Appearance of solid after heating:
c. Appearance of the inside wall of the test tube after heating:
d. Molecular equation: (include states of matter)
e. Type of reaction:
Chemical Reactions and Equations p.8
V.
Reacting HCl Solution with NaOH Solution
1.
2.
3.
4.
5.
6.
7.
8.
Using a 10-mL graduated cylinder measure 2.0 mL of 0.1 M NaOH.
Transfer the NaOH to a clean test tube.
Add one drop of phenolphthalein to the test tube.
Set the test tube aside.
Clean the graduated cylinder.
Using the 10-mL graduated cylinder measure out 3.0 mL of 0.1 M HCl.
Add one drop of phenolphthalein.
Holding the bottom of the test tube, carefully pour 1.0 mL of HCl into the test tube containing the
NaOH.
9. Add the remaining 2.0 mL of HCl dropwise to the test tube containing the NaOH until a color change is
observed.
10. This mixture can be rinsed down the drain with lots of water.
a.
Appearance of NaOH solution:
b.
Appearance of NaOH solution with phenolphthalein added:
c.
Appearance of HCl solution:
d.
Appearance of the solution in the graduated cylinder:
e.
Evidence that a chemical reaction occurred:
f.
Molecular equation: (include states of matter)
g.
Complete ionic equation: (include states of matter)
h.
Net ionic equation: (include states of matter)
i.
Type of reaction:
Chemical Reactions and Equations p.9
VI.
Heating Cu with Atmospheric O2
1. Set-up a ring stand with a clay triangle.
2. Place Bunsen burner underneath.
3. Place an eraser size amount of copper in an open crucible and
place the crucible on the clay triangle.
4. Turn on the Bunsen burner.
5. Heat for 2 minutes and turn the burner off.
6. Allow the sample to cool completely to room temperature.
7. Pull the sample of reacted material out of the crucible and crinkle it over a piece of white paper.
8. All of this material can be disposed of in the trash when cool.
a. Appearance of the unreacted Cu mesh:
b. Appearance of the Cu mesh after heating and cooling:
c. Molecular equation: (include states of matter)
d. Element Oxidized
__________________
e. Element Reduced
__________________
f. Reducing Reactant
__________________
g. Oxidizing Reactant
__________________
h. Type of reaction:
VII.
1.
2.
3.
4.
Reacting CuSO4(aq) Solution with Steel Wool (Fe)
Place an eraser amount of steel wool in the bottom of a test tube.
Cover the steel wool with 0.1 M CuSO4 solution.
Let sit 3-5 minutes. (You can continue with next experiment while you wait.
Dispose of this mixture according to your instructor’s directions.
a. Appearance of CuSO4 Solution before reaction:
b. Appearance of the steel wool (Fe) before reaction:
Chemical Reactions and Equations p.10
c. Appearance of the steel wool (Fe) after the reaction has occurred:
d. Appearance of the CuSO4 solution after the reaction has occurred:
e. Molecular equation: (include states of matter)
f. Complete ionic equation: (include states of matter)
g. Net ionic equation: (include states of matter)
h. Element Oxidized
__________________
i. Element Reduced
__________________
j. Reducing Reactant
__________________
k. Oxidizing Reactant
__________________
l. Type of reaction:
VIII. Reacting Fe(NO3)3 Solution with NaOH Solution
1. Add 10 drops of 0.1 M Fe(NO3)3 solution to a small test tube.
2. Add 20 drops of 0.1 M NaOH solution to the test tube.
3. Dispose of this mixture according to your instructor’s directions.
a.
Appearance of the Fe(NO3)3 solution:
b.
Appearance of the NaOH solution:
c.
Appearance of the reaction mixture:
Chemical Reactions and Equations p.11
IX.
d.
Molecular equation: (include states of matter)
e.
Complete ionic equation: (include states of matter)
f.
Net ionic equation: (include states of matter)
g.
Type of reaction:
Heating (NH4)2CO3
1.
2.
3.
4.
Light a Bunsen burner.
Place an eraser amount of (NH4)2CO3 in a small test tube.
Place a piece of red litmus paper on a watch glass and dampen with distilled water
Hold the test tube over the flame with test tube clamps. Simultaneously hold the watch glass with the
litmus paper side facing the tube over the opening of the test tube.
5. Heat until all solid has disappeared.
6. Carefully note the odor of any products by wafting fumes toward you.
a. Appearance of the (NH4)2CO3:
b. Initial appearance of the moist red litmus paper:
c. Description of what happened to the solid (NH4)2CO3 upon heating:
d. Appearance of the inside wall of the test tube after heating:
e. Description of the odor of the fumes from (NH4)2CO3 upon heating:
f. Appearance of the moist red litmus paper after exposure to the fumes:
g. Molecular equation: (include states of matter)
h. Type of reaction:
Chemical Reactions and Equations p.12
Reactions: Pre-Laboratory Assignment
1. Find the figures on p 37 (Question 50) of the text.
a) Which illustrate a physical change? Defend your answer.
b) Which illustrate a chemical change? Defend your answer.
2. Write the balanced chemical equation including states of matter, complete ionic reaction and net ionic
reaction for the reaction between barium nitrate, Ba(NO3)2, and sodium chromate, Na2CrO4.
a. Balanced Chemical Equation
b. Complete Ionic Reaction
c. Spectator Ions
d. Net Ionic Reaction
3. Write a complete ionic reaction and a net ionic reaction for the following balanced chemical equation.
NaOH (aq) + C6H5COOH (aq) à H2O (l) + C6H5COONa (aq)
4. Consult an internet resource for photographs of copper(I) oxide and of copper(II) oxide. Can one
visually distinguish between these two compounds? Explain.
Chemical Reactions and Equations p.13
5. Answer the following questions about the reaction:
S2- + NO3- à SO42- + NO
a. Element oxidized
___________
b. Element reduced
___________
c. Reducing reactant
___________
d. Oxidizing reactant
___________
Chemical Reactions and Equations p.14