Medgar Evers College Preparatory School
1186 Carroll Street
Brooklyn, New York 11225
Dr. Michael Wiltshire Genice Reid
Principal Assistant Principal, Supervision
Chemistry II
Laboratory Manual
Table of Contents
Lab # 17 – Charles’ Law 36
Lab # 18 – Balancing Equations Using Molecular Models 38
Lab # 19 – Mole Lab 42
Lab # 20 – Percentage of Water in Popcorn 46
Lab # 21 – Periodic Properties Part B – Solubility of Salts of Group 2 Elements 48
Lab # 22 – Solubility Curve of KNO3 50
Lab # 23 – Precipitates and Solubility Rules 54
Lab # 24 – Electrolytes 58
Lab # 25 – Properties of Acids and Bases 62
Lab # 26 – Titration 66
Lab # 27 – Rates of Reaction with Alka Seltzer 70
Lab # 28 – Collision model 72
Lab # 29 – Reactions of Acids and Metals 74
Lab # 30 – A Redox Reaction 76
Lab # 31 – Organic Chemistry I – Hydrocarbons 78
Lab # 32 – Organic Chemistry II – Functional Groups 80
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Name: ______Date: ______
Teacher: ______Period: ______OT / L
Lab # 17 – Charles’ Law
Pre-Lab Questions:
Read through the problem, introduction and procedure. Then answer the following questions on a separate sheet of paper.
1. Write a hypothesis about the problem question.
2. a. What is a phase?
b. How is the gaseous phase of matter characterized?
c. What types of movements do gas molecules possess?
3. What is the phase of a substance dependent on?
4. Compare and contrast the Celsius (°C) and the Kelvin (K) or absolute temperature scale. Include boiling point and melting point of water, as well as absolute zero.
5. State Charles’ Law.
6. Write Charles’ Law as a mathematical formula.
7. What is the equation used to convert between degrees Celsius and Kelvin?
8. In which table in your Reference Tables do you find this equation?
9. Based on your experimental procedure, should your graph be connected to the origin (point 0,0). Explain your answer.
Introduction:
The temperature dependence of the volume of a fixed quantity of gas at constant pressure was first reported by Jacques Alexandre Cesar Charles (1746–1823) in 1787. This work was repeated by Joseph Louis Gay-Lussac and is attributed at times to him as well. He found that when he plotted a graph of volume vs. temperature, it gave a linear relationship. The temperature at the zero volume intercept is called absolute zero. This is the lowest possible temperature.
In this lab, we will heat a gas contained in a flask and allow it to displace water contained in a second flask. By measuring the water displaced, we can approximate the change in volume of the gas.
Problem: What is the relationship between the temperature and volume of a gas at a constant pressure?
Materials:
ring stand 2 Erlenmeyer flasks with stoppers
thermometer hot plate
rubber tubing large graduated cylinder
Procedure:
1. Set up the Charles’ Law apparatus as shown by your teacher.
2. Create a Data Table to measure Charles’ Law.
3. Measure the starting temperature and initial volume of water.
4. Begin heating the flask. Measure the temperature and volume at fixed intervals. Obtain at least 8 data points.
5. Use your data table to prepare a graph of volume vs. temperature. Be sure to label your axes and give your graph a title.
Discussion: Answer the following questions on a separate sheet of paper.
1. As the temperature of a quantity of gas increases, explain what is happening to the molecules inside a piston in terms of the Kinetic Molecular Theory.
2. If an increase in temperature were applied to a sealed container of gas with a fixed volume, what would happen to the pressure inside the container?
3. Sketch the following graphs on loose-leaf paper. Label your axes.
- Pressure vs. Temperature
- Pressure vs. Volume
- Volume vs. Temperature
4. For each of the graphs in question 3, identify the relationship as either direct or inverse.
5. The volume of a gas is 4.00 liters at 293 K and constant pressure. For the volume of the gas to become 3.00 liters, what must the Kelvin temperature be equal to?
6. A gas occupies a volume of 40.0 milliliters at 20°C. If the volume is increased to 80.0 milliliters at constant pressure, what will the resulting temperature be equal to?
Conclusion: Answer the problem question in complete sentences on a separate sheet of paper.
Name: ______Date: ______
Teacher: ______Period: ______OT / L
Lab # 18 – Balancing Equations Using Molecular Models
Pre-Lab Questions:
Read through the problem, introduction and procedure. Then answer the following questions on a separate sheet of paper.
- Write a hypothesis about the problem question.
- Differentiate between an atom and a molecule.
- Which color represents oxygen atoms? chlorine atoms? sodium atoms?
- State the Law of Conservation of Matter.
- What is a molecule?
- a. What is a coefficient?
b. What number is represented when there is no coefficient written?
- What is a subscript?
- What is a diatomic molecule?
- What phase are most diatomic molecules found in?
- What is a chemical equation?
- How can you tell whether an equation is balanced?
- What can an equation tell us about bonding?
Problem: How can we use molecular models to help us balance equations?
Introduction: To balance equations, we follow these four rules:
1. Equations must be balanced so that the number of atoms of each element is equal on the left side (reactants) and on the right side (products) of the reaction.
2. We MUST NOT change the subscripts of any of the reactants or products; if we did that, we would be changing the very nature of the substances, and often inventing substances that don’t exist.
3. We can change the coefficients of each substance. These coefficients represent the number of moles (amount) of each substance.
4. Equations should always be balanced using the smallest whole number coefficients.
Materials: Model Kits containing 16 of each type of atom and 24 bonds:
White = Hydrogen Green = Chlorine, Sodium, or Potassium
Red = Oxygen Black = Carbon, Nitrogen or Aluminum
Procedure and Observations:
1. Construct models of each of the following molecules:
a. Hydrogen gas (H2) b. Chlorine gas (Cl2) c. Hydrochloric acid (HCl)
2. Using the molecules constructed in Step 1, determine how many molecules of each substance would be needed to balance the equation for the formation of hydrochloric acid:
__H2(g) + __Cl2(g) à __HCl(g)
REMINDER: You cannot modify the molecules themselves. All you can do is build more molecules of the same three substances.
· Sketch models of all of the molecules you constructed:
Hydrogen gas + Chlorine Gas à Hydrochloric Acid
Check your work. When the equation is balanced:
How many atoms of hydrogen are on the left side of the equation? _____
How many atoms of hydrogen are on the right side of the equation? _____
How many atoms of chlorine are on the left side of the equation? _____
How many atoms of chlorine are on the right side of the equation? _____
3. Repeat the procedure step 2 for the following reaction:
___NaN3 à ___Na + ___N2
· Sketch models of all of the molecules you constructed:
Sodium Azide à Sodium metal + Nitrogen Gas
Check your work. When the equation is balanced:
How many atoms of sodium are on the left side of the equation? _____
How many atoms of sodium are on the right side of the equation? _____
How many atoms of nitrogen are on the left side of the equation? _____
How many atoms of nitrogen are on the right side of the equation? _____
4. Using models, balance the following equations:
a. ___N2 + ___H2 à ___NH3
b. ___CH4 + ___O2 à ___CO2 + ___H2O
c. ___Al + ___O2 à ___Al2O3
5. Given the equation below, construct models to represent the equation as written.
4K + 4H2O à 4KOH + 2H2
a. Does this equation below obey the law of conservation of matter?
______
______
b. Explain why this equation is not considered “balanced” according to the rules given in the introduction.
______
Discussion: Answer the following questions on a separate sheet of paper.
1. Why is it necessary for an equation to be balanced?
2. Why can’t the subscripts be changed in a compound?
3. Based on the equation in Procedure Step 4c, answer the following questions. Show all work.
- How many moles of O2 would be needed to produce 6 moles of Al2O3?
- How many moles of Al would be needed to produce 2 moles of Al2O3?
- How many moles of O2 would be needed to react completely with 6 moles of Al?
Conclusion: Answer the problem question in complete sentences on a separate sheet of paper.
Name: ______Date: ______
Teacher: ______Period: ______OT / L
Lab # 19 – Mole Lab
Pre-Lab Questions:
Read through the problem, introduction and procedure. Then answer the following questions on a separate sheet of paper.
1. Write a hypothesis about the problem question.
2. Based on your procedure, how will you know when the reaction has gone to completion?
3. Where do you find the formula for mole calculations on your Reference Table?
4. What is the gram formula mass of sulfuric acid (H2SO4)?
a. What is the mass, in grams, of 2 moles of sulfuric acid?
b. How many moles are contained in a 49 g sample of sulfuric acid?
5. What is the gram formula mass of calcium hydroxide, Ca(OH)2?
a. What is the mass, in grams, of 0.64 moles of calcium hydroxide?
b. How many moles are contained in a 116 g of calcium hydroxide?
Problem: How do we convert between mass and moles?
Introduction: A mole of anything contains the Avogadro’s Number (6.02 x 1023) of that thing. Just as a dozen of anything always contains 12 of that thing. So:
1 dozen eggs = 12 eggs
1 dozen donuts = 12 donuts
1 mole of eggs = 6.02 x 1023 eggs
1 mole of atoms = 6.02 x 1023 atoms
As you have seen from your work in balancing equations, when elements and compounds react they react in definite proportions of atoms and molecules. For example, given the equation:
N2(g) + 3H2(g) à 2NH3(g)
This equation tells you that 1 molecule of nitrogen gas reacts with 3 molecules of hydrogen gas to form 2 molecules of ammonia (NH3) gas. Because a mole is a fixed number of atoms, this also means that 1 mole of nitrogen gas will react with 3 moles of hydrogen gas to form 2 moles of ammonia gas. The reason we use moles is because individual atoms and molecules are too small to see or measure.
When we go into a lab, we cannot directly measure the number of moles of a substance. Instead, we measure the mass of a substance and can then calculate the number of moles using the gram formula mass. The gram formula mass is the mass, in grams, of 1 mole of a substance. Calculating this is made easier by the fact that the gram formula mass, in grams, is numerically equal to the atomic mass of an element, or the sum of the atomic masses of every element in a compound. For example:
1 mole of carbon atoms = 12 g of carbon atoms
1 mole of carbon dioxide (CO2) molecules = 44 g of carbon dioxide (CO2) molecules
In order to use the correct amounts of the substances in a chemical reaction, we must be able to convert between grams and moles.
To convert from moles to mass (grams):
Mass (g) = # of moles x gram formula mass
To convert from mass (grams) to moles:
Moles = mass (g) / gram formula mass
Materials:
Triple-beam balance cups chalk
distilled water crucible baking soda
vinegar
Part A: Determining the number of moles in your full name.
Procedure: Take a piece of chalk (CaCO3) and weigh it. Write the names of everyone in your group in block letters, colored in. Reweigh the chalk. The idea is to use as much chalk as possible to see an appreciable difference when reweighed.
Results:
Initial mass of chalkFinal mass of chalk
Mass of chalk used
Moles of chalk used
Part B: Determining the number of moles in a sip.
Procedure: Weigh an empty cup. Fill the cup halfway with spring water and weight it. Take a sip of water from the cup and reweigh it.
Results:
Initial mass of waterFinal mass of water
Mass of water in sip
Moles of water in sip
Part C: Determining the number of moles of carbon dioxide released in a reaction.
Procedure: Weigh a graduated cylinder. Add 5 mL of vinegar to the graduated cylinder and weigh it. Weigh a crucible. Add a splintful of baking soda to the crucible and weigh it. Pour the vinegar into the crucible and wait until the reaction stops. Reweigh the crucible with the vinegar and baking soda in it.
Results:
Initial mass of vinegarInitial mass of baking soda
Mass of vinegar + baking soda before mixing
Mass after mixing
Mass of CO2 released
Moles of CO2 released
Discussion: Answer the following questions on a separate sheet of paper.