Name:
Name: ___________________________________________________ Lab #: ______
Laboratory Safety in the Classroom
The video you are about to view is the most important one you will see all year. It will show you how to prevent accidents in the lab and how to deal with them once they do happen. The Chemistry Lab is a dangerous place and no matter how careful we are, accidents can and will happen. Please watch the video carefully. As you watch, answer the questions below.
Each section of the video is followed by a summary. It would be best for you to watch each segment undistracted, and then answer the questions during the summary. This work will be collected and graded. Anyone who fails to show proficiency in lab safety will not be allowed to work in the lab until they can pass a safety test.
I. HANDLING CHEMICALS
1. The boy in the video spilled the chemical because the container was too _________________.
2. When diluting acids, you should always add _______________ to ______________, not the reverse.
3. List 3 precautions that will keep YOU from getting contaminated or burned by chemicals.
a. _______________________________________________________________________
b. _______________________________________________________________________
c. _______________________________________________________________________
4. When working with volatile liquids, you should work under a _________________________.
5. To prevent contamination of chemicals you should
a. _______________________________________________________________________
b. _______________________________________________________________________
6. If a chemical is spilled, you should ______________________________________________.
7. When you are done working, you should clean up by ________________________________
______________________________________________________________________________.
II. BURNERS AND GLASSWARE
8. Before lighting a burner, you should check the hose for ____________ and ____________.
9. A cool flame has a ____________ color; a hot flame has a ____________ color.
10. Immediately turn off a burner if:
a. ___________________________________________________________________
b. ____________________________________________________________________
c. ____________________________________________________________________
11. List 3 precautions for heating glassware.
a. _____________________________________________________________________
b. ______________________________________________________________________
c. _____________________________________________________________________
III. THERMOMETER SAFETY
12. List 2 safety rules when working with thermometers.
a. ______________________________________________________________________
b. _______________________________________________________________________
IV. CLOTHING
13. List 7 items to remember for dressing safely in the lab.
a. _______________________________________________________________________
b. _______________________________________________________________________
c. _______________________________________________________________________
d. _______________________________________________________________________
e. _______________________________________________________________________
f. _______________________________________________________________________
V. BEHAVIOR IN LAB
14. Correct student behavior includes:
a. _____________________________________________________________________
b. ______________________________________________________________________
c. ______________________________________________________________________
VI. EMERGENCY EQUIPMENT
15. If you get a cut in lab, you should:
a. _______________________________________________________________________
b. _______________________________________________________________________
c. _______________________________________________________________________
17. If chemicals get in your eyes you should use the ____________________ for __________ minutes.
18. If chemicals get on your skin, you should ________________________________________
___________________________________________________________________________.
19. If chemicals, equipment or the room begin to burn you can:
a. _______________________________________________________________________
b. _______________________________________________________________________
20. Clothing and students that are burning can be extinguished by:
a. ________________________________________________________________________
b. ________________________________________________________________________
Name: _________________________________________________________ Lab #: _______
Bunsen Burner Lab
Purpose: To learn and practice how to light and adjust a Bunsen burner.
Materials: Bunsen burner, matches, ring stand, ring clamp, wire mesh, 250mL beaker, graduated cylinder, 100mL water, safety goggles, hot hands and colored pencils.
Procedure: Carefully read the directions on how to light a Bunsen burner below. Color the diagram to show how a correctly adjusted flame should look. Set up the ring stand and heat the water until small bubbles begin to form. Shut down the burner, allow it to cool and clean up your station.
Almost all laboratory burners used today are modifications of a design by the German chemist Robert Bunsen. In Bunsen’s fundamental design, also widely used in domestic and industrial gas burners, gas and air are premixed by admitting the gas at relatively high velocity from a jet in the base of the burner. This rapidly moving stream of gas causes air to be drawn into the barrel from aide ports and to mix with the gas before entering the combustion zone at the top of the burner.
The burner is connected to a gas cock by a short length of rubber or plastic tubing. With some burners, the gas cock is turned to the fully on position when the burner is in use, and the amount of gas admitted to the burner is controlled by adjusting a needle valve in the base of the burner. In burners that do not have this needle valve, the gas flow is regulated by partly opening or closing the gas cock. With either type of burner the gas should always be turned off at the gas cock when the burner is not in use (to avoid possible dangerous leakage).
Examine the construction of your burner (Figure 1.1) and familiarize yourself with its operation. A burner is usually lighted with the air inlet ports nearly closed. The ports are closed by rotating the barrel of the burner in a clockwise direction. After the gas has been turned on and lighted, the size and quality of the flame is adjusted by admitting air and regulating the flow of gas. Air is admitted by rotating the barrel; gas is regulated with the needle valve, if present, or the gas cock. Insufficient air will cause a luminous yellow, smoky flame; too much air will cause the flame to be noisy and possibly blow out. A Bunsen burner flame that is satisfactory for most purposes is shown in Figure 1.2; such a flame is said to be “nonluminous.” Note that the hottest region is immediately above the bright blue cone of a well-adjusted flame.
[pic] [pic] [pic]
Figure 1.1 Figure 1.2 Figure 1.6
Questions and Problems:
1. Why is it necessary to turn off the gas with the gas cock rather than with the valve on the burner?
2. Why is air mixed with gas in the barrel of the burner before the gas is burned?
3. How would you adjust a burner which:
a) Has a yellow and smoky flame?
b) Is noisy with a tendency to blow itself out?
Name: _________________________________ Lab # _______
Observation Laboratory Activity
Purpose: To demonstrate how to make scientific observations and collect qualitative data.
[pic]
• What is an Observation: __________________________________________________________________________________ ____________________________________________________________________________________________________________________________________________________________________
• What in an Inference: __________________________________________________________________________________ ____________________________________________________________________________________________________________________________________________________________________Materials:
Package of Mentos
Bottle of Diet Coke
Procedure:
1. Observe the Mentos and diet coke as we open the bottle. Record the sound, color, smell, or other characteristics after opening the bottle.
2. Add the Mentos to the diet coke and watch the experiment take place.
3. Write down at least three observations and during the activity and then three more after the geyser settles.
|Object |Observation 1 |Observation 2 |Observation 3 |Observation 4 |
|During | | | | |
|After | | | | |
Discussion:
1. Why is it important to take observations while conducting an experiment?
2. Did you make qualitative observations or quantitative observations? How do you know?
3. Other than watching, what are other ways you can make observations?
4. How can making observations help you improve your scientific experiment?
Name: _________________________________________________________________ Lab #: _______
Density, Measurement and Significant Figures
During the experimentation process, data is collected, compiled, analyzed and utilized. In this exercise, students will collect data using equipment, which varies in accuracy and will be expected to recognize its limit of measurement using metric system units. Data will be used to calculate the density of different substances, and then students should comment on the reasons for this error. ALL MEASUREMENTS MUST BE TAKEN USING THE CORRECT NUMBER OF SIGNIFICANT FIGURES. BE SURE TO LABEL ALL MEASUREMENT WITH THE CORRECT UNITS.
Procedure:
I. DISTANCE
1) Using a 30 cm ruler, measure and label the length and width of a piece of paper. Make all measurements in cm and to the greatest degree of accuracy that the ruler will allow. Record the results.
2) On the piece of paper, draw a diagonal from one corner to the other. Measure and record this value in cm.
|Length of Paper |Width of Paper |Length of Diagonal |
| | | |
3) Measure the length and width of a student desk using a 30 cm ruler. Measure it again using a meter stick. Record results.
4) Measure the length and width of a tile in the hallway using a 30 cm ruler. Measure it again using a meter stick. Record results.
| |30 cm Ruler |Meter Stick |
|Length of student desk | | |
|Width of student desk | | |
|Length of tile | | |
|Width of tile | | |
II. DENSITY
1) Measure the length, width and height of a wood block using the 30 cm ruler. (Note: Volume = LxWxH) Using the scale, find the mass of that same wood block. Record your results.
|Length of wood block |Width of wood block |Height of wood block |
| | | |
|Volume of wood block |Mass of wood block |Density of wood block |
| | | |
2) Find the mass of a dry 50 mL graduated cylinder. Add water to the cylinder, about halfway, and read the exact volume. Then, find the mass of the cylinder with the water. Record your results.
|Mass of dry 50 mL grad cyl. |Mass of grad cyl. + water |Mass of water |
| | | |
|Volume of water |Density of water | |
| | | |
3) Find the mass of a copper cylinder. Add about 25 mL of water to a 50 mL graduated cylinder. Record the exact volume of the water. Then, carefully drop the copper cylinder in the graduated cylinder and record the new volume.
|Mass of copper cylinder |Mass of 50 mL grad cyl. |Mass of 50 mL grad cyl. + water |
| | | |
|Mass of water |Volume of water |Density of water |
| | | |
|Volume of water + copper cylinder |Volume of copper cylinder |Density of copper cylinder |
| | | |
III. VOLUME OF ONE DROP
1) Place exactly 5 mL of water in a 10 mL graduated cylinder. Using the dropper, count the number of drops needed to bring the meniscus level to exactly 6 mL.
Number of drops of water in 1 mL: _________________________
Analysis
Calculations:
1) Calculate all the densities required to complete the data table.
2) Calculate the percent error for the following:
Density of water: 1.0 g/ml
3) Using the meter stick as the accepted length of the student desk, calculate the percent error of the measurement taken with the 30 cm ruler.
4) Calculate the volume of one drop of water.
Questions:
1) Why is the meter stick the more accurate of the two devices used to measure the student desk?
2) 100.0 milliliters of gas has a mass of 0.2898 grams. What element is the gas? (Hint: use table S)
Name: __________________________________________________________ Lab #: _______
Separation of a Mixture
Purpose: To separate the components of a mixture using the techniques of decanting, filtration and evaporation.
Materials: Salt, sand and iron mixture and any lab equipment necessary to carry out your approved procedure.
Procedure:
1) You will be given a beaker filled with sand, salt and iron filings.
2) Design a procedure to separate what you have into its three separate components.
3) Write out all the steps of the procedure on the lines provided below and show them to the teacher before starting.
4) Gather all the materials for your approved procedure.
5) Show the teacher the three substances when you have separated them and have the teacher sign on the line below to indicate your procedure was successful. An unsuccessful procedure will not receive full credit!!!!!
__________________________________________________________________________________
__________________________________________________________________________________
__________________________________________________________________________________
__________________________________________________________________________________
__________________________________________________________________________________
Teacher Signature: ___________________________________
Questions/Conclusions:
1) On a separate sheet of paper, draw a picture of the filtration set-up.
2) What is a mixture?
3) Were the contents of the beaker a mixture? Why?
4) Were the contents of the beaker homogeneous or heterogeneous?
5) Was anything in the beaker an element? If so, what?
6) Was anything in the beaker a compound? If so, what?
7) Define evaporation.
8) Describe how you separated the iron from the mixture.
9) Describe how you separated the sand from the mixture.
10) Describe how you separated the salt from the mixture.
Name: ____________________________________________________________ Lab #: _______
Probability of an Electron
Purpose: What is the “probability” model for atomic structure?
Pre-Lab: Define electron cloud.
Procedure:
1. Put on goggles.
2. Set up Target as instructed, on cardboard and on the floor.
3. Aim and drop a dart toward the target (on the floor) according to instructions.
4. Repeat procedure #3 for a total of 100 drops.
5. Clean up
Data: Record your results on the following table.
|Region # |# of Hits |
|1 | |
|2 | |
|3 | |
|4 | |
|5 | |
Conclusion: Don’t forget the purpose question.
1. According to your graph, in what region will a dart most likely hit? If you dropped a dart one more time, can you predict EXACTLY where it will hit? Explain.
2. How are the answers to question #1 analogous to the probability of locating an electron in an atom?
3. If the dart hit represents the location for a single electron at a specific moment in time, what does the dark dot in region 1 represent AND what does the region most frequently hit represent?
Name: ________________________________________________________ Lab #: _______
Emission Spectroscopy
Objectives:
1. Demonstrate an understanding of the relationship between atomic structure and emission spectroscopy.
2. Observe continuous and bright-line spectra.
3. Identify an unknown metal by means of a flame test and its bright line spectrum.
Discussion:
All atoms give off electromagnetic radiation if their gases or ions are energized by heating or by a high voltage electric discharge. If the light emitted by a gas is passed through a spectroscope, a pattern of narrow lines of light is produced. Each element produces its own distinct pattern that differs from the pattern of every other element. The particular pattern of frequencies of light emitted by an atom is referred to as its emission spectrum, or bright-line spectra. The emission spectrum of an element can be used as a means of identification, just as fingerprints or DNA can be used to identify a human being.
The unique pattern of frequencies of light emitted by an atom corresponds to the set energies given off as electrons drop from higher allowed energy levels to lower energy levels closer to the nucleus. An electron can be raised from its lowest allowed energy level (ground level) to other higher allowed energy levels by absorbing certain set amounts of energy. This “excited” electron cannot remain at any of these higher energy levels if there are unoccupied lower energy levels closer to the positively charged nucleus. The electron is attracted back to one of the lower allowed energy levels. As the electron drops back down it emits a photon, a set amount of energy, in the form of electromagnetic radiation. As you observe the spectra of elements, you will notice that some lines are brighter than others. This is because more electrons are taking that particular quantum leap.
In this experiment, you will use a spectroscope to examine continuous spectra from light sources and the bright-line spectra of some gaseous elements. You will observe the color that various metallic salts impart to a Bunsen burner flame and then examine the bright line spectra that results when the light from the colored flame is passed through a spectroscope. You can use the information you obtain to determine the identity of an unknown salt.
Pre-Lab Questions:
1. According to the modern theory of the atom, where may an atom’s electrons be found? ____________
____________________________________________________________________________________
2. How do electrons become “excited”? ____________________________________________________
____________________________________________________________________________________
3. What form of energy emission accompanies the return of excited electrons to the ground state?
____________________________________________________________________________________
4. How should the burner flame be adjusted for the flame test? __________________________________
Data and Observations:
Part 1: Using colored pencils, draw the spectral lines that you see when each heated gas is viewed through the spectroscope.
|Gas Tube |VIOLET |BLUE |GREEN |YELLOW |ORANGE |RED |
|Fluorescent | |
|Bulb | |
| | |
| | |
| | |
| | |
|Argon | |
|Ar | |
| | |
| | |
| | |
| | |
|Helium | |
|He | |
| | |
| | |
| | |
| | |
|Deuterium | |
|2H | |
| | |
| | |
| | |
| | |
|Mercury | |
|Hg | |
| | |
| | |
| | |
| | |
|Krypton | |
|Kr | |
| | |
| | |
| | |
| | |
Part 2: Flame Test
For each substance burned in the Bunsen burner flame, record the color of the flame and any other observations you may make.
|Substance |Observations |
| | |
| | |
| | |
| | |
| | |
Analysis and Conclusions:
1. Compare the spectra produced by Mercury and fluorescent sources.
_____________________________________________________________________________________
_____________________________________________________________________________________
2. How can the difference in the brightness of spectral lines be explained? _________________________
_____________________________________________________________________________________
3. How can we identity an unknown salt? ___________________________________________________
_____________________________________________________________________________________
4. Prior to its discovery on Earth, the existence of Helium was first confirmed in the sun. Explain how this can be possible. _______________________________________________________________________
_____________________________________________________________________________________
5. Astronomers study the universe by analyzing electromagnetic radiation of all frequencies from microwaves to X-rays. What kinds of information may be obtained in this way?
_____________________________________________________________________________________
_____________________________________________________________________________________
_____________________________________________________________________________________
Name: ____________________________________________________________ Lab#: _______
Isotopes of “Pennium”
Introduction: Unless you are a coin collector you probably think all United Stated pennies are pretty much the same. To the casual observer, all the pennies in circulation do seem to be identical in size, thickness and composition. But just as elements have one or more isotopes with different masses, the pennies in circulation have different masses. In this investigation, you are going to use pennies with different masses to represent different “isotopes” of an imaginary element called Pennium, or Pe. Remember that chemical isotopes are atoms that have the same number of protons, but different numbers of neutrons. Thus, chemical isotopes have nearly identical chemical properties, but some different physical properties.
In this investigation, you will determine the relative abundance of the isotopes of Pennium and the masses of each isotope. You will then use this information to determine the atomic masses of the isotopes of the element. This average is based on both the mass and the relative abundance of each isotope as it occurs in nature.
Pre-lab:
1. What do the 20 pennies in this investigation represent? _____________________________________
___________________________________________________________________________________
2. What do the different masses of the pennies represent? _____________________________________
___________________________________________________________________________________
3. What information do you need to calculate the average atomic mass for an element?
___________________________________________________________________________________
Problem: What are the masses and relative abundances of isotopes of Pennium and what is the atomic mass of the element?
Materials: balance, 20 pennies
Procedure:
1. Remove the pennies from the resealable bag and count them to make sure that there are 20 Determine and record the combined mass of your 20 pennies.
2. Find the mass of each penny separately. In the data table, record the year the penny was minted and its mass to the nearest 0.001 g
3. Place the 20 pennies in the resealable bag and return the pennies and the balance. Clean up your work area and wash your hands.
Data Table:
|Penny |Year |Mass |
|1 | | |
|2 | | |
|3 | | |
|4 | | |
|5 | | |
|6 | | |
|7 | | |
|8 | | |
|9 | | |
|10 | | |
|11 | | |
|12 | | |
|13 | | |
|14 | | |
|15 | | |
|16 | | |
|17 | | |
|18 | | |
|19 | | |
|20 | | |
Calculations:
1. Combined mass of all 20 pennies __________________________
2. Calculate the percent abundance of each isotope in your sample.
(Percent abundance = total mass of pennies for each isotope x 100%)
Combine mass of all 20 pennies
3. Calculate the average atomic mass of each isotope.
(average mass = total mass of pennies of each isotope)
number of pennies of that isotope
4. Using the decimal amount and the average atomic mass of each isotope, calculate the atomic mass of Pe.
(avg. mass of isotope 1)x(% abundance) + (avg. mass of isotope 2)x(% abundance)…………
Conclusions:
1. Was the mass of 20 pennies equal to 20 times the mass of one penny? Explain. ___________________
_____________________________________________________________________________________
2. In what year(s) did the mass of Pe change? How could you tell? _______________________________
_____________________________________________________________________________________
3. How can you explain the fact that there are different “isotopes” of pennium? _____________________
_____________________________________________________________________________________
4. Why are the atomic masses for most elements not whole numbers? _____________________________
_____________________________________________________________________________________
5. How are the three isotopes of hydrogen (hydrogen-1, hydrogen-2, hydrogen-3) alike? How are they different? ____________________________________________________________________________
_____________________________________________________________________________________
_____________________________________________________________________________________
6. Copper has two isotopes, copper-63 and copper-65. The relative abundance or copper-63 is 69.1% and copper-65 is 30.9%. Calculate the average atomic mass of copper.
Name: ______________________________________________________________ Lab #: _______
Top Secret Periodic Table
Because of the exceptional skill you have demonstrated in the organization of the Periodic Table, you have been chosen for a top-secret mission! The mission, should you decide to accept it (and you will), is to work with the sketches of the characters given to you. These represent families of secret agents, BUT the most important member has never been sketched.
You are to organize the pictures and sketch the missing secret agent!
Clue 1: You could begin by grouping the people by similar characteristics OR you could sequence the pictures, then you could make shorter rows and create columns.
Clue 2: Each secret agent is different from every other one in TWO of the properties. No two sketches have the same amount or kind of these properties. If you can find one of these two, it will be possible to sequence the sketches correctly.
Clue 3: All members of a period will have something in common and all members of a group will have something in common.
Assignment
1. Once you have organized the groups and periods, glue the secret agents on paper to hand it in.
2. Answer the following questions:
a. In what TWO ways are all the secret agents different
i. _______________________________________________________________________
ii. _______________________________________________________________________
b. What do the agents in a PERIOD have in common?
____________________________________________________________________________
c. What do the agents in a GROUP have in common?
____________________________________________________________________________
3. Draw the missing agent in his/her spot on the table.
4. Relate FIVE characteristics of the agents to properties of elements on the Periodic Table.
a. _________________________________________________________________________
b. _________________________________________________________________________
c. __________________________________________________________________________
d. __________________________________________________________________________
e. __________________________________________________________________________
Name: __________________________________________________________ Lab #: _______
Trends of the Periodic Table
Purpose: To plot several properties of elements and observe the trend as you move across a period or down a group within the periodic table.
Materials: Reference Tables for Chemistry, pencil
Procedure:
1) On the graphs provided, plot the properties listed in #2 for the elements in Group 16 and then again for those in Period 3. YOU WILL NEED TO REFER TO TABLE S FOR THE INFORMATION!
2) Graphs to plot: a) Ionization energy
b) Electronegativity
c) Atomic radius
3) For each graph, connect the dots to make as smooth a line as possible.
Questions & Conclusions:
1) State the Periodic Law
2) Looking at your graphs explain the change in each property as you go down a group and across a period.
3) Define the following terms:
a) Ionization energy:
b) Atomic radius:
c) Electronegativity:
4) Explain why you skipped the transition metals when plotting the trends for elements in Period 3.
Name: __________________________________________________________ Lab #: _______
Heating Curve of Water
Introduction: The law of conservation of energy states that energy cannot be created or destroyed, however energy may be changed fro on form to another. Phase changes in matter occur because of heat lost or absorbed in the substance. Heat energy is converted or absorbed in order to increase the kinetic and potential energy of molecules. As a sample receives energy, the system changes, for example there may be an increase in temperature or a phase change that occurs because of the change in heat. A plot of the temperature versus time for a phase change is called a heating curve.
Purpose: To create a heating curve for H2O starting from its melting temperature and ending at its boiling temperature.
Materials: Safety goggles, stopwatch, hot plate with stirrer function, 250 mL beaker, ice thermometer, rubber stopper, clamp, graph paper and ring stand.
Procedure:
1) Put on safety goggles.
2) Set up thermometer stand. Attach the clamp to the ring stand. Place the thermometer inside the thermometer camp.
3) Fill the 250 mL beaker to the 150 mL mark with ice. Then fill up the container now containing ice with water to the 150 mL mark.
4) Place the beaker with ice under the ring stand setup and lower the thermometer into the ice/water mixture.
5) Record the temperature of the ice/water mixture as the starting temperature (time 0:00) in the data table.
6) Place a hot plate under the set up for the thermometer. Place the beaker of ice on the hot plate. Place the thermometer back into the beaker. Make sure that the thermometer does not touch the bottom of the beaker.
7) Turn on the hot plate to 10, turn on the stirrer and start the timer. After 1 minute, record the temperature of the ice/water mixture in the beaker.
8) Using the chart provided, record the temperature of the ice/water in the beaker after each minute.
9) Continue to record the temperature on the chart every minute until the water has been boiling for a total of 3 minutes.
10) Place an asterisk ( * ) on the data table at the time and temperature when all of the ice has melted and then again when the water boils.
11) Turn off the hot plate after you are finished recording the temperatures. Graph your data and label the graph following the directions provided. Then answer all conclusion questions.
Data:
|Time |Temperature (°C) |Time |Temperature (°C) |
|0:00 | |11:00 | |
|1:00 | |12:00 | |
|2:00 | |13:00 | |
|3:00 | |14:00 | |
|4:00 | |15:00 | |
|5:00 | |16:00 | |
|6:00 | |17:00 | |
|7:00 | |18:00 | |
|8:00 | |19:00 | |
|9:00 | |20:00 | |
|10:00 | |21:00 | |
If extra space is needed please continue the table on a piece of paper.
Graph Directions & Labels:
1) Label the x-axis on your graph as time. Label the y-axis as temperature. Make an appropriate scale for both axes.
2) Plot the points from your data table on the graph. Connect the points.
3) Trace with colored pencils the following parts of your line graph:
a. Trace the section in blue where the potential energy is increasing.
b. Trace the section in green where the kinetic energy is increasing.
4) Label the phase(s) of matter present at each section of your graph that is highlighted.
5) Label the melting and boiling temperatures.
Questions and Conclusions:
1) Explain what is happening, in terms of energy, to the water molecules in the parts of your graph that you traced in blue.
2) Explain what is happening, in terms of energy, to the water molecules in the parts of your graph that are traced in green.
3) What phase changes are endothermic processes?
4) What phase changes are exothermic processes?
5) Does the graph you made in this lab represent an endothermic or exothermic heating curve?
6) Explain how this lab illustrates the law of conservation of energy.
Graph:
______________________________________
| | | | |
|Shape | | | |
|Volume | | | |
Name: __________________________________________________________ Lab #: _______
The Burning Cheese Puff
Purpose: How many calories are there in a cheese puff?
Materials: Goggles, beaker, balance, 100 ml graduated cylinder, matches, cheese puffs, cheese puff apparatus, ring stand, iron rings, wire gauze, water and thermometer.
Procedure:
1) Put on your goggles.
2) Randomly select two samples of cheese puffs (do not eat).
3) Weigh the beaker. Record data in data table below.
4) Measure 100 ml of water using the graduated cylinder. Pour the 100 ml of water into the beaker and weigh it again. Record.
5) Record the initial temperature of the water.
6) Place the cheese puffs on the paper clip as demonstrated by your teacher and determine their mass. Record. Place this close under the beaker.
7) Light the cheese puffs with a match. Try to get them to burn completely. WHILE THEY ARE BURNING, gently stir the water with the thermometer and record the highest temperature it reaches.
8) Weigh the cheese puff + apparatus again. Record.
9) Wash your hands and clean up.
Data:
|Mass of empty beaker |Initial Temp of water |Initial mass of cheese puffs |
| | | |
|Mass of beaker + water |Finial Temp of water |Final mass of cheese puffs |
| | | |
Calculations: SHOW ALL WORK!
1) Calculate the mass of the water being used.
2) Calculate the heat absorbed by the water.
3) Calculate the mass of the cheese puff that burned.
4) Calculate the heat per gram of cheese puff. (Hint: “per” means…?)
5) List the class data. What is the average heat per gram from the class data?
6) Convert your answer from #5 to cal/g (1 cal = 4.18J).
7) The package says 150kcal/30 g. Convert this to kcal/1g. Compare your experimental value (#6) to the given on the package.
Conclusion:
1) Define heat and temperature.
2) What causes inaccuracy in your answer?
Name: ____________________________________________________________ Lab #: _______
Bonding and Chemical Nomenclature A
Purpose:
( To investigate how elements will bond together in specific ratios to form a compound
( To learn how to properly name compounds based on the elements that they contain
Materials: Scissors, sheet of various ions, glue, paper.
Procedure:
1) Use a pair of scissors to cut out the ions on the sheet provided.
2) Use the ions to determine the formulas for the compounds listed in Data Table #1.
3) Write the formulas on Data Table #1 and glue the ions to the space provided for each compound.
Data Table #1
|Compound |Formula | Ions Used |
|Tin (II) fluoride | | |
| | | |
| | | |
| | | |
| | | |
| | | |
| | | |
|Tin (IV) oxide | | |
| | | |
| | | |
| | | |
| | | |
| | | |
| | | |
| | | |
|Iron (III) oxide | | |
| | | |
| | | |
| | | |
| | | |
| | | |
| | | |
| | | |
|Magnesium hydroxide | | |
| | | |
| | | |
| | | |
|Sodium hydrogen carbonate | | |
| | | |
| | | |
| | | |
| | | |
| | | |
| | | |
|Copper (II) hypochlorite | | |
Data Table #2
|Compound |Name |Compound |Name |
|KF | | |Dinitrogen Hexafluoride |
|CaI2 | | | Sodium Sulfide |
|AlN | | | Mercury (II) Oxide |
|HCl | | | Lead (II) Nitrite |
|H2S | | | Aluminum Carbonate |
|N4O8 | | |Copper (II) Phosphate |
|LiMnO4 | | | Potassium Chromate |
Discussion
1. Are all the compounds in parts 1 and 2 ionic or covalent? How do you know?
2. What is the charge of all the cations and what are the charges of all the anions?
3. List 3 polyatomic Ions used in the chemicals in parts 1 or 2.
4. What if we didn’t have the stock system, would we be able to name compounds with different oxidation states? Why?
Ions
Name: ____________________________________________________________ Lab #: _______
Bonding and Chemical Nomenclature B
Purpose:
( To investigate how elements will bond together in specific ratios to form a compound
( To learn how to properly name compounds based on the elements that they contain
Materials: Scissors, sheet of various ions, glue, paper.
Procedure:
1) Use a pair of scissors to cut out the ions on the sheet provided.
2) Use the ions to determine the formulas for the compounds listed in Data Table #1. Keep in mind that the ion charges should add up to zero for a neutral compound.
3) Write the formulas on Data Table #1 and glue the ions to the space provided for each compound.
4) Use Table E in your reference tables to help you name the compounds in Data Table #2.
Data Table #1
|Compound |Formula | Ions Used |
|Tin (II) fluoride | | |
| | | |
| | | |
| | | |
| | | |
| | | |
|Tin (IV) oxide | | |
| | | |
| | | |
| | | |
| | | |
| | | |
| | | |
|Iron (III) oxide | | |
| | | |
| | | |
| | | |
| | | |
| | | |
| | | |
|Magnesium hydroxide | | |
| | | |
| | | |
| | | |
| | | |
|Zinc Oxide | | |
| | | |
| |
| | | |
| | | |
|Copper (II) hypochlorite | | |
Data Table #2
|Compound |Name |Compound |Name |
|KF | | |Magnesium Chloride |
|CaI2 | | | Sodium Sulfide |
|AlN | | | Mercury (II) Oxide |
|HCl | | | Lead (II) Nitrate |
|H2S | | | Aluminum Carbonate |
|NO | | |Copper (II) Phosphate |
|LiMnO4 | | | Dihydrogen monoxide |
Discussion
1. Are all the compounds in parts 1 and 2 ionic or covalent? How do you know?
2. What is the charge on a metal ion? Nonmetal ion?
3. List 3 polyatomic Ions used in the chemicals in parts 1 or 2.
4. What if we didn’t have the stock system, would we be able to name compounds with different oxidation states? Why?
Ions
Name: ____________________________________________________________ Lab #: _______
Bonding and Chemical Nomenclature C
Purpose:
( To investigate how elements will bond together in specific ratios to form a compound
( To learn how to properly name compounds based on the elements that they contain
Materials: Scissors, sheet of various ions, glue, paper.
Procedure:
1) Use a pair of scissors to cut out the ions on the sheet provided.
2) Use the ions to determine the formulas for the compounds listed in Data Table #1. Keep in mind that the ion charges should add up to zero for a neutral compound.
3) Write the formulas on Data Table #1 and glue the ions to the space provided for each compound.
4) Use Table E in your reference tables to help name the compounds in Data Table #2. Remember the roman numerals in the stock system show the charge, not the amount of the ions.
Data Table #1
|Compound |Formula | Ions Used |
|Tin (II) fluoride | | |
| | | |
| | | |
| | | |
| | | |
| | | |
| | | |
|Tin (IV) oxide | | |
| | | |
| | | |
| | | |
| | | |
| | | |
| | | |
| | | |
|Iron (III) oxide | | |
| | | |
| | | |
| | | |
| | | |
| | | |
| | | |
| | | |
|Magnesium hydroxide | | |
| | | |
| | | |
|Zinc Oxide | | |
| | | |
| | | |
| | | |
|Sodium hydrogen carbonate | | |
| | | |
| | | |
| | | |
| | | |
Data Table #2
|Compound |Name |Compound |Name |
|KF | | |Magnesium Chloride |
|CaI2 | | | Sodium Sulfide |
|AlN | | | Mercury (II) Oxide |
|HCl | | | Calcium Sulfate |
|H2S | | | Aluminum Carbonate |
|NO | | |Copper (II) Phosphate |
|LiNO3 | | | Dihydrogen Monoxide |
Discussion
1. Are all the compounds in parts 1 and 2 ionic or covalent? How do you know?
2. When a chemical name has a roman numeral in it, what does that number mean?
3. List 3 polyatomic Ions used in the chemicals in parts 1 or 2.
4. What is the charge on a metal ion? Nonmetal ion?
Ions
Name: _____________________________________________________________ Lab #:_______
Covalent Bonding Lab A
Purpose: What is the value of structural formulas?
Procedure: Fill in the table below and use it to construct molecular models.
|Color of ball |Group # |# of valence e- |Total # needed for “Octet” |# of bonding sites |
|Yellow |1 | | | |
|Black |14 | | | |
|Red |16 | | | |
|Green, violet, orange |17 | | | |
|Light blue |15 | | | |
1. Make and draw models of compounds using the elements in data table 1.
2. Make and draw a model of CH3OH (methanol) in data table 2.
3. Make and draw models of compounds that contain multiple bonds in data table 3.
4. Make and draw models of compounds with isomers in data table 4.
Date Table 1: Build each molecule using the model kits. Draw the molecule you built in the space below.
a. H2O b. CH4
_______________ _______________
c. Cl2 d. CCl4
_______________ _______________
e. NH3 f. H2
_______________ _______________
g. CF4 h. HCl
_______________ _______________
i. H2S j. PCl3
_______________ _______________
Data Table 2:
CH3OH
Data Table 3:
a. O2 b. N2
c. C2H4
Data Table 4:
a. C4H10 b. C5H12
Conclusion:
1) What is C2H4 molecule and bond type?
2) What shape is CO2?
3) List 3 examples not listed above that are non polar covalent bonds?
Name: _____________________________________________________________ Lab #:_______
Covalent Bonding Lab B
Purpose: What is the value of structural formulas?
Procedure: Fill in the table below and use it to construct molecular models.
|Color of ball |Group # |# of valence e- |Total # needed for “Octet” |# of bonding sites |
|Yellow = H |1 | | | |
|Black = C |14 | | | |
|Red = O, S |16 | | | |
|Green, violet, orange = |17 | | | |
|F, Cl, Br, I | | | | |
|Light blue = N, P |15 | | | |
1. Make and draw models of compounds using the elements in data table 1.
2. Make and draw a model of CH3OH (methanol) in data table 2.
3. Make and draw models of compounds that contain multiple bonds in data table 3.
4. Make and draw models of compounds with isomers in data table 4.
Date Table 1: Build each molecule using the model kits. Draw the molecule you built in the space below. a. H2O b. CH4
_______________ _______________
c. Cl2 d. CCl4
_______________ _______________
e. NH3 f. H2
_______________ _______________
g. CF4 h. HCl
_______________ _______________
i. H2S j. PCl3
_______________ _______________
Data Table 2:
CH3OH
Data Table 3:
a. O2 b. N2
c. C2H4
Data Table 4:
a. C4H10 b. C5H12
Conclusion:
1) What is Cl2 molecule and bond type?
2) What shape is CO2?
3) List 3 examples from above that are non polar covalent bonds?
Name: _____________________________________________________________ Lab #:_______
Covalent Bonding Lab C
Purpose: What is the value of structural formulas?
Procedure: Fill in the table below and use it to construct molecular models.
|Color of ball |Group # |# of valence e- |Total # needed for “Octet” or “Duet” |# of bonding sites |
|Yellow = H |1 | | | |
|Black = C |14 | | | |
|Red = O, S |16 | | | |
|Green, violet, orange = |17 | | | |
|F, Cl, Br, I | | | | |
|Light blue = N, P |15 | | | |
1. Make and draw models of compounds using the elements in data table 1.
2. Make and draw a model of CH3OH (methanol) in data table 2.
3. Make and draw models of compounds that contain multiple bonds in data table 3.
4. Make and draw models of compounds with isomers in data table 4.
Date Table 1: Build each molecule using the model kits. Draw the molecule you built in the space below.
a. H2O b. CH4
_______________ _______________
c. Cl2 d. CCl4
_______________ _______________
e. NH3 f. H2
_______________ _______________
g. CF4 h. HCl
_______________ _______________
i. H2S j. PCl3
_______________ _______________
Data Table 2:
CH3OH
Data Table 3: These contain double and triple bonds. Use the springs instead of the sticks.
a. O2 b. N2
c. C2H4
Data Table 4:
a. C4H10 b. C5H12
Conclusion:
1) What is H2O molecule and bond type?
2) What shape is CO2?
3) List 3 examples from above that are non polar covalent bonds?
Name ________________________________________________ Lab #: _______
Types of Reactions A
Objective: To observe some examples of different types of chemical reactions and to write equations for the reactions.
Introduction: In Inorganic Chemistry there are four main types of reactions: synthesis, decomposition, single replacement and double replacement. Synthesis is a reaction that starts with two smaller chemicals and results in one larger product. Decomposition, the opposite of synthesis, starts with one larger reactant and results in two or more smaller products. Single replacement, can be characterized by one element replacing another in a chemical reaction. Double replacement, can be characterized by two elements changing what they are bonded to in a chemical reaction
In this lab, you will carry out procedures and make observations. Then you will determine which reactions fall under each type and write the symbols for it.
Materials: Iron filings, Magnesium pieces, Zinc pieces, Copper sulfate solution (CuSO4), Sodium hydrogen carbonate (NaHCO3), Sodium iodide solution (NaI), Lead (II) nitrate solution (Pb(NO3)2), test tube clamp, ring stand, Bunsen burner, wood splint, test tubes, beaker, goggles, test tube rack, tongs and wax pencil.
Safety precautions: Wear your safety goggles at all times. If you get any chemicals on you, wash them off immediately. Never leave your Bunsen burner unattended.
Procedure:
Part A
1) Set up your Bunsen burner and light it.
2) Get a small strip of Magnesium ribbon.
3) Using tongs to hold the ribbon, place the ribbon into the flame of the Bunsen burner and observe what happens.
4) Record your observations and reactants below.
5) Name the reaction and write a complete reaction in the boxes below.
|Observations | Reactants | Type of Reaction |
| | | |
| Balanced Equation | |
Part B
1) Place a small amount (1 scoopful) of sodium hydrogen carbonate (NaHCO3) into a test tube. Use the test tube clamp to attach it in a horizontal position to a ring stand. Spread out the sodium hydrogen carbonate in the bottom half of the test tube. See Figure 1.
2) Heat the sodium hydrogen carbonate with a burner for 3 minutes.
3) Light a wood splint.
4) While continuing to heat the sodium hydrogen carbonate, insert the flaming splint into the mouth of the test tube. See Figure 2. Note any deposits on the side of the test tube.
5) Record your observations and reactants below.
6) Name the reaction and write a complete balanced equation below.
[pic]
|Observations | Reactants | Type of Reaction |
| | | |
| Balanced Equation | |
Part C
1) Place a small amount of copper sulfate (CuSO4) solution into each of two test tubes. Label these 1 and 2 and place them back in the test tube rack.
2) Place a small amount of iron filings (Fe) in test tube #1.
3) Place a small piece of zinc metal (Zn) in test tube #2.
4) Observe any reactions and record your observations on the table below.
5) Name the reactions and write complete balanced equations.
|Observations | Reactants | Type of Reaction |
| | | |
| Balanced Equation (1) | |
|Observations | Reactants | Type of Reaction |
| | | |
| Balanced Equation (2) | |
Part D
1) Add a few milliliters of lead (II) nitrate (Pb(NO3)2) to a clean test tube. Place it in the test tube rack.
2) Add a few drops of sodium iodide (NaI) to the lead nitrate.
3) Record any observations you observe on the table below.
4) Name the reaction and write a complete balanced equation below.
|Observations | Reactants | Type of Reaction |
| | | |
| Balanced Equation | |
Conclusion Questions:
1) In the decomposition of sodium hydrogen carbonate, the products are sodium carbonate, carbon dioxide gas and water. Based on your observations, how is a flaming splint used to test for the presence of carbon dioxide?
2) If we were to add Au to the CuSO4, would a reaction happen? Why or why not?
3) Based on your observations, how can we determine if a double replacement reaction has taken place?
4) Compare synthesis with decomposition.
5) Based on you knowledge of how a double replacement reaction works, could we use a different chemical besides NaI to create a precipitate reaction? Why?
6) Write an equation (other than the ones used in this lab) showing an example of each of the 4 types of reactions.
Name ________________________________________________ Lab #: _______
Types of Reactions B
Objective: To observe some examples of different types of chemical reactions and to write equations for the reactions.
Introduction: In Inorganic Chemistry there are four main types of reactions: synthesis, decomposition, single replacement and double replacement. Synthesis is a reaction that starts with two smaller chemicals and results in one larger product. Decomposition, the opposite of synthesis, starts with one larger reactant and results in two or more smaller products. Single replacement, can be characterized by one element replacing another in a chemical reaction. Double replacement, can be characterized by two elements changing what they are bonded to in a chemical reaction
In this lab, you will carry out procedures and make observations. Then you will determine which reactions fall under each type and write the symbols for it.
Materials: Iron filings, Magnesium pieces, Zinc pieces, Copper sulfate solution (CuSO4), Sodium hydrogen carbonate (NaHCO3), Sodium iodide solution (NaI), Lead (II) nitrate solution (Pb(NO3)2), test tube clamp, ring stand, Bunsen burner, wood splint, test tubes, beaker, goggles, test tube rack, tongs and wax pencil.
Safety precautions: Wear your safety goggles at all times. If you get any chemicals on you, wash them off immediately. Never leave your Bunsen burner unattended.
Procedure:
Part A
1) Set up your Bunsen burner and light it.
2) Get a small strip of Magnesium ribbon.
3) Using tongs to hold the ribbon, place the ribbon into the flame of the Bunsen burner and observe what happens.
4) Record your observations and reactants below.
5) Name the reaction and write a complete reaction in the boxes below.
|Observations | Reactants | Type of Reaction |
| | | |
| Balanced Equation | |
Part B
1) Place a small amount (1 scoopful) of sodium hydrogen carbonate (NaHCO3) into a test tube. Use the test tube clamp to attach it in a horizontal position to a ring stand. Spread out the sodium hydrogen carbonate in the bottom half of the test tube. See Figure 1.
2) Heat the sodium hydrogen carbonate with a burner for 3 minutes.
3) Light a wood splint.
4) While continuing to heat the sodium hydrogen carbonate, insert the flaming splint into the mouth of the test tube. See Figure 2. Note any deposits on the side of the test tube.
5) Record your observations and reactants below.
6) Name the reaction and write a complete balanced equation below.
[pic]
|Observations | Reactants | Type of Reaction |
| | | |
| Balanced Equation | |
Part C
1) Place a small amount of copper sulfate (CuSO4) solution into each of two test tubes. Label these 1 and 2 and place them back in the test tube rack.
2) Place a small amount of iron filings (Fe) in test tube #1.
3) Place a small piece of zinc metal (Zn) in test tube #2.
4) Observe any reactions and record your observations on the table below.
5) Name the reactions and write complete balanced equations.
|Observations | Reactants | Type of Reaction |
| | | |
| Balanced Equation (1) | |
|Observations | Reactants | Type of Reaction |
| | | |
| Balanced Equation (2) | |
Part D
1) Add a few milliliters of lead (II) nitrate (Pb(NO3)2) to a clean test tube. Place it in the test tube rack.
2) Add a few drops of sodium iodide (NaI) to the lead nitrate.
3) Record any observations you observe on the table below.
4) Name the reaction and write a complete balanced equation below.
|Observations | Reactants | Type of Reaction |
| | | |
| Balanced Equation | |
Conclusion Questions:
1) In the decomposition of sodium hydrogen carbonate, the products are sodium carbonate, carbon dioxide gas and water. Based on your observations, how is a flaming splint used to test for the presence of carbon dioxide?
2) How can we predict if a single replacement reaction will take place?
3) Based on your observations, what kinds of products result from a double replacement reaction?
4) Compare synthesis with decomposition.
5) How is single replacement different from double replacement?
6) Write 4 equations (other than the ones used in this lab) and label them to with the appropriate type of reaction.
Name ________________________________________________ Lab #: _______
Types of Reactions C
Objective: To observe some examples of different types of chemical reactions and to write equations for the reactions.
Introduction: In Inorganic Chemistry there are four main types of reactions: synthesis, decomposition, single replacement and double replacement. Synthesis is a reaction that starts with two smaller chemicals and results in one larger product. Decomposition, the opposite of synthesis, starts with one larger reactant and results in two or more smaller products. Single replacement, can be characterized by one element replacing another in a chemical reaction. Double replacement, can be characterized by two elements changing what they are bonded to in a chemical reaction
In this lab, you will carry out procedures and make observations. Then you will determine which reactions fall under each type and write the symbols for it.
Materials: Iron filings, Magnesium pieces, Zinc pieces, Copper sulfate solution (CuSO4), Sodium hydrogen carbonate (NaHCO3), Sodium iodide solution (NaI), Lead (II) nitrate solution (Pb(NO3)2), test tube clamp, ring stand, Bunsen burner, wood splint, test tubes, beaker, goggles, test tube rack, tongs and wax pencil.
Safety precautions: Wear your safety goggles at all times. If you get any chemicals on you, wash them off immediately. Never leave your Bunsen burner unattended.
Procedure:
Part A
1) Set up your Bunsen burner and light it.
2) Get a small strip of Magnesium ribbon.
3) Using tongs to hold the ribbon, place the ribbon into the flame of the Bunsen burner and observe what happens.
4) Record your observations and reactants below.
5) Name the reaction and write a complete reaction in the boxes below.
|Observations | Reactants | Type of Reaction |
| | Magnesium and Oxygen | |
| Balanced Equation | |
Part B
1) Place a small amount (1 scoopful) of sodium hydrogen carbonate (NaHCO3) into a test tube. Use the test tube clamp to attach it in a horizontal position to a ring stand. Spread out the sodium hydrogen carbonate in the bottom half of the test tube. See Figure 1.
2) Heat the sodium hydrogen carbonate with a burner for 3 minutes.
3) Light a wood splint.
4) While continuing to heat the sodium hydrogen carbonate, insert the flaming splint into the mouth of the test tube. See Figure 2. Note any deposits on the side of the test tube.
5) Record your observations and reactants below.
6) Name the reaction and write a complete balanced equation below.
[pic]
|Observations | Reactants | Type of Reaction |
| |NaHCO3 | |
| Balanced Equation: This compound will break into sodium carbonate, carbon |___NaHCO3 ( ___Na2CO3 + ____CO2 + ____H2O |
|dioxide gas and water. | |
Part C
1) Place a small amount of copper sulfate (CuSO4) solution into each of two test tubes. Label these 1 and 2 and place them back in the test tube rack.
2) Place a small amount of iron filings (Fe) in test tube #1.
3) Place a small piece of zinc metal (Zn) in test tube #2.
4) Observe any reactions and record your observations on the table below.
5) Name the reactions and write complete balanced equations.
|Observations | Reactants | Type of Reaction |
| |Iron and copper sulfate | |
| Balanced Equation (1) | |
|Observations | Reactants | Type of Reaction |
| |Zinc and copper sulfate | |
| Balanced Equation (2) | |
Part D
1) Add a few milliliters of lead (II) nitrate (Pb(NO3)2) to a clean test tube. Place it in the test tube rack.
2) Add a few drops of sodium iodide (NaI) to the lead nitrate.
3) Record any observations you observe on the table below.
4) Name the reaction and write a complete balanced equation below.
|Observations | Reactants | Type of Reaction |
| | Pb(NO3)2 and NaI | |
| Balanced Equation | |
Conclusion Questions:
1) What other substance besides Magnesium is used to create the synthesis reaction in our first experiment?
2) In the decomposition of sodium hydrogen carbonate, the products are sodium carbonate, carbon dioxide gas and water. Based on your observations, how is a flaming splint used to test for the presence of oxygen gas?
3) How can you predict if a single replacement reaction will happen based on its reactants?
4) Based on your observations, what kinds of products result from a double replacement reaction?
5) How is single replacement different from double replacement?
6) Write 4 equations (other than the ones used in this lab) and label them to with the appropriate type of reaction.
Name: ____________________________________________________________ Lab #: _______
Precipitation in Double-Replacement Reactions for Solubility Studies
Objectives:
( To observe precipitation reactions by mixing aqueous solutions of cations and anions
( To master the use of table F of the reference tables
Introduction:
In chemistry, the term precipitation does not refer to meteorological phenomena such as rain or snow. But precipitation occurs in solution when two chemicals react to form a product that is insoluble in water and falls out of solution like rain or snow. A precipitate is a solid substance that separates from solution during a chemical reaction. A precipitate can be identified by the cloudy, milky, gelatinous or grainy appearance it gives to the mixture. The solid might even settle to the bottom of the container.
A barium sulfate precipitate can be produced by the reaction of barium chloride and sodium sulfate.
BaCl2(aq) + Na2SO4(aq) ( 2NaCl(aq) + BaSO4(s)
In this reaction, the barium sulfate is the precipitate. It is insoluble in water, thus it precipitates out of solution as a solid (s). These reactions are called double replacement reactions. In this type of reaction, both reactants, swap partners, as seen on the product side.
Purpose:
Lead, silver and calcium commonly undergo reactions that form precipitates. In this lab you will mix solutions of lead, silver and calcium with other compounds in solution. You will observe and describe the precipitates that are formed. Then you will answer the questions on the work sheet. You will be provided with the experimental page that all testing will be performed on. Follow the directions on this page carefully. Especially the clean up!
Data Table: Below in each box, indicate the color precipitate you observed. If there was no precipitate, write the letters NVR (No Visible Reaction).
| |KI |NaCl |NaOH |FeCl3 |Na2SO4 |Na3PO4 |Na2CO3 |CuSO4 |
|AgNO3 | | | | | | | | |
|Pb(NO3)2 | | | | | | | | |
|CaCl2 | | | | | | | | |
Questions:
1. What is a precipitate? What cations commonly form precipitates? What polyatomic ions commonly form precipitates? (Use Table F)
2) How can you tell if a precipitate forms when you mix two solutions?
3) Using table F of your reference table, indicate if the following compounds would be soluble or insoluble in water (note: the presence of a precipitate indicates a compound is insoluble).
|Compound |Soluble/Insoluble |Compound |Soluble/Insoluble |
|NaCl | |CaCO3 | |
|LiI | |Na2CO3 | |
|(NH4)2SO4 | |K2CrO4 | |
|NH4ClO3 | |(NH4)3PO4 | |
|AgCl | |ZnS | |
|PbI2 | |Al(OH)3 | |
4) Predict the products that these reactions will produce. Be sure to label the solutions as aqueous and the precipitates as solids.
A) AgNO3(aq) + KI(aq) (
B) AgNO3(aq) + NaCl (
C) Pb(NO3)2(aq) + KI(aq) (
D) CaCl2(aq) + CuSO4(aq) (
Name: _________________________________________________________ Lab #: _______
Percent Composition of Sugar in Bubble Gum
Pre-Lab: Percent composition is a way to determine the percentage of an element in comparison to the compound it is in.
1. Write the formula for percent composition:
2. Determine the percent composition of each element in the compound CuSO4
Purpose: In this lab, you will find the percent composition of sugar in a piece of bubble gum. You will compare it to the actual percent of sugar in the gum.
Materials: Balance, gum
Procedure: SAVE the wrapper for weighing. DO NOT PUT UNWRAPPED GUM ON THE BALANCE.
1. Zero the balance.
2. Find the mass of the unchewed gum and its wrapper.
3. Find the mass of the wrapper.
4. Chew the gum for 15 minutes.
5. Place the gum and wrapper on the balance.
6. Record all masses in the data table.
7. Determine the amount of sugar in gum and calculate the % composition.
Data:
|Gum 1 |Grams |
|1. Mass of unchewed gum + wrapper | |
|2. Mass of wrapper | |
|3. Mass of unchewed gum (1-2) | |
|4. Mass of chewed gum + wrapper | |
|5. Mass of wrapper | |
|6. Mass of chewed gum | |
|7. Mass of Sugar | |
|Gum 2 |Grams |
|Mass of unchewed gum + wrapper | |
|Mass of wrapper | |
|Mass of unchewed gum | |
|Mass of chewed gum + wrapper | |
|Mass of wrapper | |
|Mass of chewed gum | |
|Mass of Sugar | |
Conclusion:
1. Calculate the % composition of sugar in bubble gum A and B. Show all work.
2. Obtain the accepted amount of sugar or sugar substitute in bubble gum A and B from the teacher.
3. Using the accepted amount of sugar in bubble gum A and B, calculate the accepted percent of sugar in bubble gum A and B. Show all work.
4. Based on your data, which of the two gum samples would be better for a diabetic to eat and why?
Name: _____________________________________________________________ Lab #: _______
Percent Composition of Hugs and Kisses
Purpose:
In this activity you will compare and contrast Hershey Kisses( and Hershey Hugs( using your knowledge of chemistry and the nutritional information provided on the food labels. You will use the percent composition formula to solve the percent by mass problems.
Procedure:
1) The information on the nutritional labels is based on a serving size of nine pieces of candy. Use the labels to determine the following information for a SINGLE Hershey Kiss( and a single Hershey Hug(.
| |Kiss |Hug |
|mass (g) | | |
|calories | | |
|total fat (g) | | |
|total carbohydrates (g) | | |
|protein (g) | | |
2) Calculate the percent mass of carbohydrate in a Hershey Kiss( and a Hershey Hug(. (Show all work!)
Kiss Hug
3) Calculate the percent mass of fat in a Hershey Kiss( and a Hershey Hug(. (Show all work!)
Kiss Hug
4) Calculate the percent mass of protein in a Hershey Kiss( and a Hershey Hug(. (Show all work!)
Kiss Hug
5) C8H10N4O2 is the molecular formula for Caffeine. Caffeine is present in most chocolate products that we consume. Calculate the percent mass of each element in Caffeine.
Kiss Hug
Name:______________________________________________ Lab #: __________
S’more Stoichiometry
Introduction: “Stoichio” means element and “metry” means the process of measuring. The mass and quantity relationships amount reactants and products are found using the process of stochiometry.
Materials: 1 plate, 1 full graham cracker, 2 marshmallows, 4 chocolate nibs, Bunsen burner, stick, goggles, and matches
Procedure:
1. Gather all materials above.
2. Set up Bunsen burner and check hose to make sure there are no holes in it.
3. Light Bunsen burner.
4. Separate graham cracker into two pieces using the lines on the graham cracker as a guide. Designate one as the bottom cracker and one as the top cracker.
5. Place both marshmallows on stick.
6. Hold stick, with marshmallow end out, over the flame of the Bunsen burner. DO NOT put in the flame. Flaming marshmallows will set off fire alarm.
7. As marshmallow is cooking, place the chocolate nibs on the graham cracker half you will designate as the bottom graham cracker.
8. When the marshmallow center is soft, and exterior is golden brown, remove marshmallow and place on top of chocolate.
9. Place graham cracker on top of marshmallow and squish marshmallow to create your s’more.
Data:
|Substance |Symbol |Unit Mass |
|Graham Cracker | | |
| |S |7.00 g |
|Marshmallow | | |
| |Mm |7.10 g |
|Chocolate Piece | | |
| |Or |3.30 g |
|S’more | | |
| |S2Mm2Or4 |g |
Questions:
1. Calculate the unit mass of the S’more (S2Mm2Or4) and place it in the data table above.
2. Using your product from question one, create a balanced chemical equation.
3. What does the equation tell us? What do the coefficients represent?
4. If there are 454g of marshmallows in one bag, how many marshmallows are there in one bag of marshmallows? (Show all work)
5. How many s’mores can you make with one bag of marshmallows? (Show all work)
6. How many graham crackers and chocolate segments are needed to make the amount of s’mores you can make from one bag of marshmallows? (Show all work)
Graham Crackers needed Chocolate segments needed
7. Using the data from table 1, convert your number of graham crackers and chocolate segments into mass (gram) values.
Graham Crackers needed Chocolate segments needed
8. How many boxes of graham crackers do you need to buy and chocolate bars if a box of graham crackers has a mass of 254 grams, and one chocolate bar has a mass of 49.5 grams?
Graham Crackers needed Chocolate segments needed
Now we will transfer this process into the language of chemical reactions.
If we were to add a piece of solid Cu to an aqueous solution of silver nitrate, the silver would be replaced in a single replacement reaction forming aqueous copper (II) nitrate and solid silver.
___Cu + ___AgNO3 ( ___Cu(NO3)2 + ___Ag
How much silver is produced is 15.00 grams of Cu is added to the solution of excess silver nitrate? Show all work.
Step 1: Balance the equation above.
Step 2: Convert grams of Cu to moles of Cu.
Step 3: Using mole ratios, find out how much silver is produced based on the number of moles of Cu that react.
Step 4: Convert moles of Ag produced to grams of Ag produced.
Name: _________________________________________________________ Lab #: _______
Solubility Curve for KNO3
Introduction:
The maximum amount of solute that will dissolve in a given amount of solvent is called its solubility. What factors determine the solubility of a substance? Certainly the identity of the solute affects the amount of the substance that can dissolve. For example, sodium iodide is more soluble than sodium chloride in a given amount of water. The identity of the solvent also affects the solubility of a substance. Sodium chloride is highly soluble in water but not very soluble in ethanol. Temperature of the solvent is another factor affecting solubility. The solubility of most solids varies directly with temperature. In other words, the higher the temperature of the solvent, the more solute will dissolve-that is, the greater the solubility of the solid.
In this investigation, you will study the relationship between the solubility of potassium nitrate (KNO3) and the temperature of the water solvent. Different amounts of KNO3 will be dissolved in a given amount of water at a temperature close to its boiling point. You will observe each solution as it cools and note the temperature at which crystallization occurs. Crystallization indicates a saturated solution, or one that contains the maximum amount of KNO3 in that amount of water. From this solubility data, a solubility curve for KNO3 can be constructed.
Pre-Lab Discussion:
1. Define the following terms:
a) Solubility:
b) Saturated solution:
c) Unsaturated solution:
2. Determine the solubility of sodium sulfate, Na2SO4, in grams per 100 g of water if 0.94 g of Na2SO4 in 20 g of water results in a saturated solution.
3. Why is it necessary to warm the thermometer in Procedure Step 2 before placing it into the solutions?
4. A student stated that the solubility of potassium chloride at 20(C was 36 g of KCl per 100 g of solution. What is wrong with this statement?
5. How do you know when the solid is completely dissolved?
Problem: How does the solubility of potassium nitrate depend on temperature?
Materials:
Goggles Potassium Nitrate (KNO3)
4 test tubes, 18 x 150 mm Graduated cylinder, 10 ml
Marking pen Water
Test tube rack Utility clamp
Beaker, 400 ml Ring Stand
4 Thermometers Weigh boat
Bunsen burner Test tube holder
Balance
Procedure:
1) Put on your goggles.
2) Fill a 400 ml beaker half full of tap water, place thermometer in it, and heat the water over the Bunsen burner until its temperature is about 90(C. While you are waiting for the water to heat, go on to steps 3 and 4.
3) Label four test tubes 1-4 with a marking pen. Place them in a test tube rack.
4) Place the following masses of potassium nitrate (KNO3) into each test tube:
4.0 g in test tube 1
8.0 g in test tube 2
12.0 g in test tube 3
14.0 g in test tube 4
5) Add 10 ml (10 g) of water to each test tube.
6) Stir each of the salt solutions with a clean stirring rod.
7) Place each test tube in the water bath.
8) When all the solute in test tube #1 has dissolved, use a test tube holder to remove it from the water bath. Give it on final stir, then remove the stirring rod and rinse it off. Place the warm thermometer into test tube #1.
9) Watch test tube #1 for the first sign of crystallization (it takes the longest to crystallize) and when it occurs record the temperature in the Data Table. Remove test tube #2 from the hot water bath and repeat steps 8 and 9.
10) Repeat steps 8 and 9 for test tubes 3 and 4.
11) Place all the test tubes back in the hot water bath and re-dissolve the solid. Flush the solutions down the drain with plenty of water. Turn off the Bunsen burner. Clean up your work area and wash your hands before leaving the laboratory.
Observations:
|TT # |Grams solute/10 ml water |Temperature (°C) |Grams solute/100 ml water |
|1 | | | |
|2 | | | |
|3 | | | |
|4 | | | |
Critical Thinking: Analysis and Conclusions:
1. Construct a solubility curve for KNO3 by graphing the mass of KNO3 per 100 g of H2O (y-axis) versus temperature (x-axis). Connect the points in a smooth curve.
| | | | |
| |Vinegar |NaOH |Vinegar |NaOH |Vinegar |NaOH |
|Initial Volume mL | | | | | | |
|Final Volume mL | | | | | | |
|Volume Used mL | | | | | | |
| | | | | | | |
| | | | | | | |
| |Trial 4 |Trial 5 | | |
| |Vinegar |NaOH |Vinegar |NaOH | | |
|Initial Volume mL | | | | | | |
|Final Volume mL | | | | | | |
|Volume Used mL | | | | | | |
Calculations:
1. Calculate the Molarity of NaOH for trial #1. Show your work.
2. Calculate the Molarity of NaOH for trial #2. Show your work.
3. Calculate the Molarity of NaOH for trial #3. Show your work.
4. Calculate the Molarity of NaOH for trial #4. Show your work.
5. Calculate the Molarity of NaOH for trial #5. Show your work.
6. Average the molarities for the five trials. Show your work.
Conclusions: Compare your results to the rest of the class. Do you feel they are accurate? If not, what could be your sources of error? If you think they are accurate, explain why?
Name: _______________________________________________________________________________________ Lab #: _______
Acid/Base Indicators
Introduction: In this activity, you will investigate the colors of six indicators: methyl orange, bromthymol blue, phenolphthalein, litmus, bromcresol green and thymol blue.
Materials: 36-well plate, 36 toothpicks and the following reagents: vinegar, seltzer water, household ammonia, 0.1 M NaHCO3, distilled water, solution of unknown pH and the six indicators listed above.
Procedure:
1. Put on your goggles.
2. Obtain a well plate. Place the sheet of paper (given to you by your teacher) under the well plate. This paper will serve to label the wells to correspond to the data table on the next page. Notice that each box on the data table corresponds to a well on the plate. The solutions with various pH values and unknown values are listed across the top and the indicators are listed along the side.
3. Place several drops of each solution listed at the top of your sheet in the proper well. Then add one drop of methyl orange to each well in the first row. Stir the wells to mix the solutions, using a fresh toothpick in each well. Record the resulting colors in the data table. Continue in the same way with bromthymol blue indicator added to each solution as you proceed across the second row. Record the resulting colors for each well as before. Complete the other rows. Estimate the pH of the unknown solution using the data you have recorded.
Over (
| |Vinegar |Soda |Distilled water |0.1 M NaHCO3 |Ammonia |Unknown |
| |pH = 3 |pH = 5 |pH = 7 |pH = 10 |pH = 12 |pH = ? |
|Methyl Orange | | | | | | |
|Bromthymol Blue | | | | | | |
|Phenolphthalein | | | | | | |
|Litmus | | | | | | |
|Bromcresol Green | | | | | | |
|Thymol Blue | | | | | | |
Conclusions and Questions:
1. What is the approximate pH of your unknown solution? Explain your reasoning.
2. Is your unknown an acidic, basic or neutral solution? For your unknown solution, compare the concentrations of H3O+ and OH- ions.
3. For each pH value, describe the solution as acidic, basic or neutral Next, compare the [H+] to the [OH-], using either =, > or or ................
................
In order to avoid copyright disputes, this page is only a partial summary.
To fulfill the demand for quickly locating and searching documents.
It is intelligent file search solution for home and business.
Related searches
- company name and stock symbol
- why your name is important
- native american name generator
- why is my name important
- why is god s name important
- last name that means hope
- name for significant other
- name synonym list
- me and name or name and i
- name and i vs name and me
- name and i or name and myself
- name and i or name and me