reactions lab v051413 7pm - uca | faculty sites at the...

Chemical Reactions and Equations p.1

Types of Chemical Reactions and Equations v051413_7pm

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.


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 Cl2-

b. Swap the anion partner as you move toward the products. c. Determine the formula of the compounds. d. Write the equation with reactants and products. e. Determine the state of matter of the compounds. f. Balance the equation.

Example: Reactants: CaCl2 + AgNO3 a. Ca2+ Cl- + Ag+ + NO3

- b. Ca2+ NO3

- + Ag+ + Cl- c. Ca(NO3)2 + AgCl d. CaCl2 + AgNO3 à Ca(NO3)2 + AgCl e. 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.


The compound will be If it contains Unless combined with Soluble (aq) Group 1, NH4

+ --------------- Soluble (aq) NO3

-, C2H3O2-, ClO4

-, HCO3- ---------------

Soluble (aq) Halides (F-, Cl-, Br-, I-) Ag+, Pb2+, Hg22+

Soluble (aq) SO42- Ag+, Pb2+, Hg2

2+, Ca2+, Sr2+, Ba2+ Insoluble (s) O2-, OH- Ca2+, Sr2+, Ba2+ Insoluble (s) All other anions ----------------

CaCl2 (aq) + AgNO3 (aq) à Ca(NO3)2 (aq) + AgCl (s) f. CaCl2 (aq) + 2 AgNO3 (aq) à Ca(NO3)2 (aq) + 2 AgCl (s)

(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 Cl-

b. 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: HCl + Al a. H+ Cl- + Al b. H2

+ Al + Cl- c. H2

+ AlCl3 because the charge on an Al cation is +3 d. HCl + Al à H2

+ AlCl3 e. 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) f. 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.


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 Weak Electrolytes Soluble Ionic Compounds (Solubility rules) Insoluble Ionic Compounds (solubility rules) Strong Acids Weak Acids Strong Bases 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)


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.

8Fe(s) + S8(s) à 8FeS(s)0 +20 -‐2

e-‐ gain, reductione-‐ 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.

6HCl(aq) + 2Al(s) à 3H2(g) + 2AlCl3(aq)+1 0-‐1 0 -‐1+3

e-‐ lost, oxidatione-‐ 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.


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:


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:




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)


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:


e. Type of reaction:


V. Reacting HCl Solution with NaOH Solution

1. Using a 10-mL graduated cylinder measure 2.0 mL of 0.1 M NaOH. 2. Transfer the NaOH to a clean test tube. 3. Add one drop of phenolphthalein to the test tube. 4. Set the test tube aside. 5. Clean the graduated cylinder. 6. Using the 10-mL graduated cylinder measure out 3.0 mL of 0.1 M HCl. 7. Add one drop of phenolphthalein. 8. 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:


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 __________________


VII. Reacting CuSO4(aq) Solution with Steel Wool (Fe)

1. Place an eraser amount of steel wool in the bottom of a test tube. 2. Cover the steel wool with 0.1 M CuSO4 solution. 3. Let sit 3-5 minutes. (You can continue with next experiment while you wait. 4. 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:


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:





g. Type of reaction:

IX. Heating (NH4)2CO3

1. Light a Bunsen burner. 2. Place an eraser amount of (NH4)2CO3 in a small test tube. 3. Place a piece of red litmus paper on a watch glass and dampen with distilled water 4. 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)



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.


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 ___________

reactions lab v051413 7pm - uca | faculty sites at the...

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