Chemistry Lab Report Format Lab reports may be handwritten on notebook paper or typed. Pre-labs are due the day we perform the lab. Post-

labs are due one week after we perform the lab (you can always turn them in sooner). Every section should be written in complete sentences. Every section should also have a heading in bold, caps, or underlining to distinguish between sections.

“Prelab”

o Heading- upper left hand corner (include name, name of lab and date(s) the lab will be performed o The purpose - written in one or two sentences. This is a statement of the problem to be investigated. It provides

the overall direction for the laboratory investigation and must be addressed in the conclusion. o Safety warnings - written out in one or two sentences. After reading the procedure, list all safety issues that need

to be kept in mind while performing the lab. This includes but is not limited to acids, bases, other dangerous chemicals, hot objects, open flames, and broken glass.

o All pre-lab questions must be answered in complete sentences. Calculations for math problems must be shown. You do not have to write the question but number them like they are in the lab for easy reference.

o The procedure should be read and then summarized in paragraph form. You do not have to rewrite the entire thing word-for-word. SUMMARIZE AND PARAPHRASE! Copying word-for-word will result in a grade penalty.

*All of the above is to be turned in the day you start a lab. If you do not have the pre-lab to turn in on the day of the lab, you will not do the lab in class.* “Postlab”

o Data and observations: This section will include all data that was measured and observations that were made. Data should be neatly organized in a data table, and graphs should be included when applicable. Tables and graphs should be drawn with a ruler and include titles/labels.

o Calculations/Results should be included here. This section may include a “results table.” If you did multiple trials, always show calculations for each individual trial, and you include any averages at the end if applicable.

o Post-Lab Questions should be answered. Calculations for math problems must be shown. o Conclusion – Very briefly restate the purpose of the lab, then discuss how you addressed the purpose. State and

discuss results. A good conclusion for a lab is between 5 & 7 sentences. This section is not a summary of the lab procedure. It is a place for you to discuss the main point of the lab, what was learned, surprising results, and anything else relevant to the purpose of the lab.

o Error Analysis: Discuss possible sources of error in the lab. Never cite reasons such as “we measured wrong” or “we calculated wrong.” If you did a math-based lab and have not already calculated percent error or percent yield, you must do so here.

See next page for example of lab report formatting.

Katy Lyles Flame Test Lab 8/23/2016 I. Purpose The purpose of this lab is to….. II. Safety In today’s lab, students should be aware that… III. Pre-lab Questions

1. Answers…

2. Answers… IV. Procedure In this section you will summarize and paraphrase the lab procedure. Do not copy word-for-word! (all of the above sections are considered the “pre-lab” and should be completed before the lab is done) V. Data and Observations

Chemistry Is So Awesome

VI. Calculations & Results

1. Calculate the mass of …..

87 g – 82 g = 5 g VII. Post-lab Questions 1. Answers… VIII. Conclusion Here you will write your organized and well thought-out conclusion. IX. Error Analysis Discuss sources of error and calculate % error or % yield if not done in a previous section.

AP Chemistry Lab – Observations of a Burning Candle

Background

All scientists use a variety of different measurements. There are QUALITATIVE and

QUANTITATIVE measurements. Qualitative measurements refer to properties or qualities of

something that cannot be measured, i.e. blue color, heat produced, and silvery-shiny metal.

Qualitative measurements are descriptive in nature: they describe. Quantitative measurements are

measurements that are made with the aid of an instrument, and therefore, can be reproduced by

another person. For example: 2.56 grams and 4.676 liters. Each type of measurement is a valuable

piece of information that may help you to understand and identify what is happening in the lab.

Objective

To distinguish between qualitative and quantitative measurements and recognize the

significance of both.

Prelab Questions

1. Why is it important to include both qualitative and quantitative measurements in a lab report?

2. Give five quantitative measurements about yourself.

3. Give five qualitative measurements about yourself.

4. Provide five quantitative measurements about your classroom.

Equipment and Chemicals

candle glass plate ruler balance matches

Procedure

1. Make all necessary observations of the candle while it is unlit.

2. Light the candle and observe it. Record your observations.

Data and Observations

Create an organized method of reporting your data. It should include both quantitative and

qualitative data from before, during, and after the match was burning.

Postlab Questions

None

Conclusion

See your lab report guide for how to write a stellar conclusion!

AP Chemistry Lab – Formula of a Hydrate (Inquiry)

Background Some chemical compounds, especially inorganic salts, incorporate water into their

crystalline structures. These are called hydrates. Heat can be used to dehydrate a hydrated salt

causing the water molecules to evaporate and produce an anhydrous salt which will often appear

different from its hydrate. Formulas for hydrates are written using a dot convention: a dot is used to

separate the formula of the salt from the formula of the water of hydration. A numerical coefficient

gives the molar amount of water included in the hydrate.

For example: NiCl2·6H2O is the formula for nickel(II) chloride hexahydrate. This formula

indicates that for every 1 mole of NiCl2 there are 7 moles of water present.

One key point: the dot is not a multiplication sign. When calculating the molar mass you add

the molar mass of water (multiplied by the coefficient). An everyday example of hydration is

concrete. Concrete is made by mixing Portland cement with water and aggregate materials. The

aggregate materials are the gravel and sand that add strength to the final concrete. The Portland

cement is a mixture of calcium silicates, calcium aluminate, calcium aluminoferrite and gypsum.

All of these chemicals absorb water by hydration. This means that concrete does not ‘dry’ in a

conventional sense. Instead the water mixed with the concrete combines chemically with the

materials in the cement and the resulting hydrates form a strong matrix that holds the concrete

together and makes it strong. Another interesting example of the value of hydration is the

incorporation of hydrated building materials (such as concrete, gypsum wall board and plaster). The

building materials will not rise above the 100°C boiling point of water until all of the water of

hydration has been driven off. This can help keep damage to a minimum until the fire can be put

out. In this experiment you will investigate some of the properties of hydrates and you will

determine the number of water molecules present in a formula unit of an unknown hydrate.

Objective • To identify some properties of hydrates.

• To determine the number of waters of hydration in an unknown hydrate.

Prelab Questions 1. The mineral “gypsum” is one of the most important hydrates, calcium sulfate dihydrate. It

is hard and dense. When it is heated and the molecules of water are driven off, the

anhydrous calcium sulfate that remains is known as plaster of Paris. Rehydrating the plaster

of Paris creates a highly interlocked crystalline substance known as plaster.

Write a balanced equation showing gypsum heated to form plaster of Paris and water.

2. A 2.123 g sample of washing soda, Na2CO3∙xH2O, was heated and all the water of

hydration was lost. 0.787 g of the anhydrous sodium carbonate was left behind. Calculate

the value of x.

Equipment and Chemicals Ring stand ring clay triangle crucible and cover

Bunsen burner crucible tongs analytical balance

unknown hydrate

Procedure You must write your own procedure. It must be detailed enough to indicate everything you will do

in the lab. If it is not detailed enough or is incorrect, you will not be permitted to do the experiment.

Copying a procedure (in part or in whole) from the internet or another source is plagiarism.

Data and Observations • Create an organized method of reporting your data. It should include both quantitative and

qualitative data.

Calculations The compound you used was either a copper (II) sulfate hydrate, CuSO4∙xH2O or MgSO4∙xH2O

1. Calculate the mass percent of water in your sample.

2. Show your work to calculate the integer of hydration for your sample, and write the complete

formula of the hydrate.

3. Give the full name of your sample.

Postlab Questions 1. If your sample was heated for too short a time, will your value for the mass percent for water

be too high or too low? Explain your reasoning.

2. If your crucible was still hot when you took the final mass, will your value for the mass

percent of water be too high or too low? Explain your reasoning.

Conclusion See your lab report guide for how to write a stellar conclusion!

Error Analysis Don’t forget to do this section.

AP Chemistry Lab – Determining the Percent Copper in Brass

Background Brass is a generic term for alloys of copper and zinc. In addition to these metals, brass may also contain small amounts

of iron, lead, aluminum, and tin. More than 300 different brass alloys are known, with uses ranging from decorative

hardware to architectural construction, musical instruments, and electrical switches. The amount of copper in brass

affects its color, hardness, ductility, strength, conductivity, corrosiveness, etc. Visible spectroscopy provides a simple

tool for determining the percent copper in brass.

Spectroscopy involves the interaction of electromagnetic radiation and matter. A visible spectrophotometer is used to

measure the absorption of visible light through a solution – it is particularly helpful with transition metals which have

characteristic and beautiful colors. In general, absorbance is proportional to concentration, a concept that is expressed

mathematically through Beer’s Law.

A = abc

In the equation above, A is absorbance, c is concentration in molarity, b is the path length (width of the solution’s

container), and a is known as the molar absorptivity coefficient.

Beer’s law is helpful to determine the “unknown” concentration of a metal ion in solution if absorbance is measured.

Plotting absorbance vs. concentration produces a “calibration curve” – a straight line that passes through the origin.

Objective

To use the following concepts to determine the percent copper in brass:

A serial dilution

Beer’s Law

Visible spectroscopy

Prelab Questions

1. Dissolving brass requires an oxidizing acid such as nitric acid. Nitrogen dioxide is produced

as a by-product in this reaction. Write a balanced chemical equation for the reaction of

copper metal with concentrated nitric acid to produce copper (II) nitrate, nitrogen dioxide,

and water.

2. Nitrogen dioxide is a toxic reddish-brown gas. What safety precautions are required in this

lab to protect against this hazard, as well as the hazards due to the use of concentrated acid?

3. Copper (II) ions appear blue in aqueous solution. This is the transmitted color. The

wavelengths of light that are NOT absorbed give rise to the perceived or transmitted color of

a substance. Based on the principle of complementary colors, which colors of wavelengths

of light would you expect to be most strongly absorbed by Cu2+ ions?

Equipment and Chemicals

Procedure

Part 1 – Work as a Class

1. Obtain 50 mL of the 0.4 M stock solution of Cu(NO3)2.

2. Using several 100-mL volumetric flasks, use a serial dilution procedure to create 0.2M,

0.1M, and 0.05M solutions.

3. Label the solutions, cover them with parafilm, and place them on a table that can be easily

accessed by all groups.

Part 2A – Work in your Small Group

1. Add two brass shot pieces (about 1 g) to a tared weighing dish and determine the mass.

2. Concentrated nitric acid is required to dissolve the brass shot. You will need an excess of the

acid to completely dissolve the shot. Since the reaction produces a toxic gas, NO2 (g), a

fume hood is required for this step in the procedure. Carefully measure about 6mL of the

nitric acid into a 10 mL graduated cylinder, place it in the fume hood, tilt it to the side, and

carefully drop the two pieces of brass into the cylinder. Close the fume hood.

3. Wait until tomorrow!

Part 2B – Work in your Small Group

4. Set up your colorimeter with the LabQuest. Set the wavelength to 635 nm. {If 635 isn’t

available, choose the closest number} Place a cuvette with distilled water in the colorimeter

and press the “Cal” button – it should eventually read the absorbance as zero.

5. Beginning with the most dilute solution from the serial dilution, test and record the

absorbance of the 0.05M, 0.1M, 0.2M, and 0.4M Cu2+ solutions.

6. Collect your graduated cylinder from the fume hood with your new copper solution.

Transfer it to a 100 mL graduated cylinder.

7. Using the original small graduated cylinder, add about 10 mL of water to the cylinder then

pour it into the larger cylinder. Repeat this two times. You are getting all of the copper ions

out of the graduated cylinder into the larger cylinder.

8. Record the exact volume of the cylinder in your data table.

9. Analyze the absorbance of the solution you created.

Data and Observations

Data Table

Solution of Cu2+ Absorbance Qualitative Data

Describe the solutions.

Compare the Solution from

Brass to the Solutions from

the Serial Dilution

0.05 M

0.10 M

0.20 M

0.40 M

Solution created from brass

(exact volume after dilution)

Exact volume of brass solution ____________

Postlab Questions

1. Create a graph showing the relationship between Cu2+ concentration and absorbance from

your data. Include the data point 0,0. Your graph should clearly show data points as small

circles or squares (they should not blend into the line) and a best-fit line.

2. Based on the absorbance of your brass solution, what is the final concentration of Cu+2 ions

in the brass solution?

3. Using the molarity and volume of the Cu+2 in the brass solution, determine the mass of

copper in the original brass sample.

4. Determine the percent copper in the original brass sample.

5. Why did you include the data point 0,0 in your graph? Give a full explanation.

Conclusion

Don’t forget this part!

Error Analysis

You can still do error analysis, but an actual percentage error cannot be determined. Discuss w

AP Chemistry Lab – Verifying the Stoichiometry of a Reaction

Background

Analytical chemists use a variety of methods to investigate the stoichiometry of a reaction. One of

the methods is to use a constant amount of substance A, while varying the amount of substance B.

The amount of product formed is measured, and from the maximum amount of product formed, the

stoichiometry of the reaction can be found. The results of the data are graphed. The top, flat part of

the graph will yield the maximum amount of product.

The reaction between silver nitrate and sodium chloride solutions is studied. If 10 ml of 1M solution

of sodium chloride is used, while a variable amount of 1M solution of silver nitrate is used, a data

table and graph might look like this.

The value that yields the maximum amount of product is 10 mmol of silver nitrate. Since 10 mmol

of sodium chloride reacts with 10 mmol of silver nitrate, the stoichiometric ratio is 1:1. The

products are either sodium nitrate or silver chloride. Since 10 mmol of sodium nitrate is 0.85 grams,

10 mmol of silver chloride is 1.42 grams, the product that is soluble is sodium nitrate, while silver

chloride is insoluble and is your product.

The “amount” of product can be measured in a variety of ways – mass of precipitate, height of

precipitate in identical containers, amount of heat released or absorbed, etc.

Objective

To verify the mole ratio of a chemical reaction by determining which ratio of compounds

produces the most product.

Prelab Questions

1. The graph to the right shows the results of a lab in

which 5 mL of 1.0 M lead (II) nitrate react with

varying amounts of 1.0 M KI. Write a balanced

equation.

2. Does the data on the graph support the ratio that

you wrote in Part 1? Fully explain your reasoning.

Equipment and Chemicals

Test tube rack 5 identical medium test tubes

volumetric pipette and pump

0.1 M Na3PO4 0.1 M Cobalt Chloride, CoClx ( I am deliberately not telling you which

cobalt ion)

Procedure

1. Set up five medium (identical!) test tubes in a test tube rack, number them one through five.

2. Use a volumetric pipette to precisely measure the volumes of solutions shown below into the test

tubes.

Test Tube 1 Test Tube 2 Test Tube 3 Test Tube 4 Test Tube 5

mL of 0.1 M

Na3PO4

8 mL 8 mL 8 mL 8 mL 8 mL

mL of 0.1 M

CoClx

4 mL 8 mL 12 mL 16 mL 20 mL

3. Allow the mixtures to settle for ten minutes.

4. Use a ruler to measure the height of the precipitate in each test tube and record into your data table.

5. Record all available qualitative data into your data table.

Data and Observations

Test Tube Height of Precipitate Qualitative Data

1

2

3

4

5

Graph your data in a manner similar to the graph shown in the prelab. The varying volume of one solution

should be on the x-axis and the measured data will be on the y-axis.

Postlab Questions

1. What mole ratio corresponds to the stoichiometry of the reaction?

2. Did you have Cobalt (II) chloride or Cobalt (III) chloride? Write a balanced equation to

support your decision.

3. Explain why this method allows you to find the mole ratio of reactants.

4. What is the precipitate formed in the reaction? Support your answer with experimental

evidence. How does this relate to the solubility table in your book?

5. How would the lab results be different if a student used 10. mL of sodium phosphate? (all

else remains constant)

6. How would the lab results be different if a student used 1.0 mL of sodium phosphate? (all

else remains constant)

Conclusion and Error Analysis

See your lab report guide for how to write a stellar conclusion! Don’t forget to include error

analysis.

Determination of Enthalpy Change Associated with a Reaction

As you have seen in previous experiments, a great deal can be learned by conducting an acid-base

reaction as a titration. In addition, acid-base reactions can be observed and measured

thermodynamically. In this case, the reaction is carried out in a calorimeter. If the temperature of the

reaction is measured precisely, the enthalpy of neutralization of an acid by a base (or vice versa) can

be determined. In this experiment, you will react phosphoric acid with sodium hydroxide in an open

container, which means the pressure is held constant. At constant pressure ΔHrxn = q = mcpΔT.

You will use two polystyrene cups nested in a beaker as a calorimeter, as shown in Figure 1. For

purposes

of this experiment, you may assume that the heat loss to the calorimeter and the surrounding air is

negligible. Phosphoric acid will be the limiting reactant in this experiment, and you will accordingly

be determining the enthalpy of neutralization ΔHneut, of the acid. Selecting a limiting reactant helps

ensure that the temperature measurements and subsequent calculations are as precise as possible.

OBJECTIVES

In this experiment, you will

Measure the temperature change of the reaction between solutions of sodium hydroxide and

phosphoric acid.

Calculate the enthalpy of neutralization ΔHneut, of phosphoric acid.

Compare your calculated enthalpy of neutralization with the accepted value.

Figure 1

MATERIALS

Data Collection Mechanism 0.60 M phosphoric acid solution

computer or handheld 1.85 M sodium hydroxide solution

Temperature Probe two 50 mL graduated cylinders

Polystyrene cup calorimeter (2 cups per group) ring stand

two 250 mL beakers utility clamp

PROCEDURE

1. Obtain and wear goggles. Conduct this experiment in a well-ventilated room.

2. Turn on the LabQuest Stream. Make sure Bluetooth is blue so it will connect with your iPad.

Open the Vernier graphical analysis software on your iPad. Tap “Create a new experiment”.

Tap “LabQuest Stream” and Select an Accessory. Select “Time Based Collection”. Make sure

the proper probe (temperature) is plugged into CH 1. Tap “Collect” when you are ready to start

data collection. When finished collecting data tap “Stop”. Share the data with other members

in your lab group.

3. Nest two polystyrene cups inside each other and place them into a 250 mL beaker as shown in

Figure 1. Measure out 50.0 mL of 0.60 M H3PO4 solution into the foam cup. CAUTION:

Handle the phosphoric acid with care. It can cause painful burns if it comes in contact with the

skin.

4. Use a utility clamp to suspend the Temperature Probe from a ring stand (see Figure 1). Lower

the Temperature Probe into the phosphoric acid solution.

5. Measure out 50.0 mL of 1.85 M NaOH solution and transfer it to a 250 mL beaker. CAUTION:

Sodium hydroxide solution is caustic. Avoid spilling it on your skin or clothing.

6. Start the data collection. Let the program gather and graph a few initial temperature readings,

then add the NaOH solution. Stir the reaction mixture gently and thoroughly. Review the graph

to confirm and record the initial temperature and the highest temperature of the reaction.

7. Rinse and dry the Temperature Probe, polystyrene cup, and stirring rod. Dispose of the solution

as directed.

8. Repeat the necessary steps to conduct a second reaction with new portions of the same

solutions. If directed, conduct a third reaction. Print a copy of the graph of the second trial to

include with your data and analysis.

DATA TABLE

Trial 1 Trial 2 Trial 3

Maximum temperature (°C)

Initial temperature (°C)

Temperature change (∆T)

PRE LAB QUESTIONS

1. Define ΔHrxn. What units are associated with ΔHrxn?

2. Define specific heat. What units are associated with specific heat?

3. Write the balanced molecular equation for the reaction of phosphoric acid and sodium

hydroxide.

4. A student conducts an experiment to determine the enthalpy of neutralization for phosphoric

acid reacting with sodium hydroxide. The student combines 25.0 mL of equimolar solutions

of acid and base, both having an initial temperature of 22.5 °C, in an open polystyrene

calorimeter. Each solution has a density of 1.00 g/mL and a specific heat of 4.18 J

g • °C.

The student collects data until a maximum temperature of 26.4 °C is reached.

(a) What assumption is the student allowed to make?

(b) Explain how the student knows data collection is complete.

(c) Calculate the qrxn for the neutralization reaction.

(d) Is the reaction exothermic or endothermic? Justify your answer.

(e) Calculate the ΔHneut if the initial molarities of both solutions is equal to 0.80 M. Be

sure to report your answer with the correct sign and units.

POST LAB QUESTIONS AND DATA ANALYSIS

1. What measurements did you make during this experiment?

2. Use your average temperature change to calculate the amount of heat energy, q, produced in the

reaction. In determining the mass, m, of the solution use 1.11 g/mL for the density. Use

4.18 J/(g•°C) as the specific heat, cp, of the solution.

3. Use the heat energy that you calculated in question 2 above to determine the enthalpy change,

∆Hneut,

for the reaction in units of kJ/mol of phosphoric acid. This is your experimental value of ∆Hneut.

4. Write the balanced net ionic equation for the complete neutralization of H3PO4 with NaOH.

5. Use a table of standard thermodynamic data below to calculate the ∆Hneut of neutralization for

the net ionic equation of phosphoric acid reacting completely with sodium hydroxide. Consider

this the accepted value of ∆Hneut.

Standard Enthalpy of Formation

ΔHf ° (kJ/mol)

OH− (aq) −230

H+ (aq) 0

H2O () −286

6. Calculate the percent discrepancy between the calculated (accepted) value of the ∆H of

neutralization of H3PO4 and your experimental value.

7. A student conducts two trials of this experiment, but neglects to rinse the calorimeter cup

between trials. What effect, if any, would this error have on the calculated value of the ∆H of

neutralization?

Determination of Molar Mass by Vapor Density

One of the properties that helps characterize a substance is its molar mass. If the substance in

question is a volatile liquid, a common method to determine its molar mass is to vaporize it and

apply the ideal gas law, PV = nRT to the data collected. Because the liquid is volatile, it can easily

be converted to a vapor. Volatile substances are usually composed of nonpolar molecules. As a

result the molecules have primarily London dispersion forces and very little thermal energy is

required to overcome these attractive forces since the molecules are relatively small. Therefore, the

liquid vaporizes easily and quickly at temperatures less than 100°C. While the substance is in the

vapor phase, you can measure its volume, pressure, and temperature. You can then use the ideal gas

law to calculate the number of moles of the substance. Finally, you can use the number of moles of

the gas to calculate molar mass.

OBJECTIVES

In this experiment, you will

Evaporate a sample of a liquid substance and measure certain physical properties of the substance as it condenses.

Determine the molar mass of an unknown liquid.

Figure 1

MATERIALS

Data Collection Mechanism unknown volatile liquid Temperature Probe fume hood one Gas Pressure Sensor or barometer test tube, 13 100 mm, and holder ring stand two 400 mL beakers two utility clamps hot plate aluminum foil analytical balance ice needle tissues or paper towels

Best results are

obtained when the

test tube containing

the sample is

submerged in the

water bath to just

below the foil cap.

PROCEDURE Before beginning the experiment, make sure that you have a means of measuring the barometric

pressure in the room. A conventional barometer or a Gas Pressure Sensor may be used.

1. Obtain and wear goggles. Conduct this experiment in a fume hood or well-ventilated area.

2. Trim a piece of aluminum foil so that it just covers the top of a small, 13 100 mm, test tube.

Use a needle to make a small hole in the center of the foil. Measure and record the mass of the

test tube and foil.

3. Prepare a hot-water bath by warming about 300 mL of tap water in a 400 mL beaker. Keep the

beaker on a hot plate once the water is warm. You want it to boil by Step 8.

4. Use a second 400 mL beaker to prepare an ice-water bath.

5. Set up the data collection system.

a. Connect a Temperature Probe to the interface.

b. Start the data collection program. Be sure the program shows the correct readings for the

probe.

c. There is no need to store and graph the data for this experiment.

6. Obtain a liquid sample of an unknown volatile compound. Pour about 0.5 mL of the liquid into

the test tube and quickly cover the test tube with the aluminum foil. Use your fingernail to make

an air-tight seal with the foil just under the lip of the test tube. Place the test tube in the hot

water bath. Make sure that the foil is above the water level, but submerge your test tube as far as

possible without making contact with the bottom of the beaker (see Figure 1).

7. Immerse the Temperature Probe in the hot-water bath (see Figure 1). Do not allow the tip of the

probe to touch the beaker.

8. Heat the hot-water bath to boiling and maintain the boiling as your sample of liquid vaporizes.

Note that some of your sample will escape the test tube through the needle hole in the foil. This

process also serves to flush the air out of the test tube.

9. Keep the test tube in the boiling-water bath for at least three minutes after all of the liquid in the

test tube has vaporized. Watch the temperature readings and record the temperature of the

boiling water bath, which will be used in the ideal gas law calculations.

10. Use a test-tube holder to quickly transfer the test tube to the ice water bath. Cool the test tube for

about one minute, then remove it and dry it completely. Measure and record the mass of the test

tube and the aluminum foil top.

11. Record the room’s barometric pressure.

12. Rinse out the test tube and fill it to the top with tap water. Cover the test tube with the aluminum

foil. Measure and record the mass of the test tube, water, and foil.

DATA TABLE

Trial 1 Trial 2

Mass of test tube and foil cover (g)

Temperature of hot water bath (ºC)

Mass of test tube and foil and gas sample (g)

Barometric pressure (atm)

Mass of test tube and foil and water (g)

PRE-LAB QUESTIONS

1. What is the difference between a vapor and a gas?

2. A student performs an experiment designed to determine the molar mass of a sample of an

unknown volatile liquid. The following data was collected:

Trial 1

Mass of test tube and foil cover (g) 7.5228 g

Temperature of water bath (ºC) 99°C

Mass of test tube and foil and condensed gas (g) 7.5387 g

Barometric pressure (atm) 0.987 atm

Mass of test tube and foil and water (g) 16.1228 g

(a) Determine the mass of the condensed unknown.

(b) Assuming the density of water to be 1.00 g/mL, determine the total volume of the test tube.

(c) Starting with the ideal gas law, substitute the fact that n is equal to grams/molar mass and

derive an expression for calculating the experimental molar mass of the student’s volatile

liquid sample.

(d) Calculate the molar mass for this student’s sample.

(e) Calculate the percent error for this student’s experimental determination of the molar mass

assuming the unknown liquid was methyl alcohol (CH3OH).

POST-LAB QUESTIONS AND DATA ANALYSIS

1. Determine the mass of the condensed unknown.

2. Use the mass and density of the water in the test tube from Step 12 of the procedure to calculate

the volume of the test tube.

3. Use the expression you derived for Pre-Lab Question # 2 part (c) along with the data collected

to calculate the molar mass of your unknown compound.

4. Use your experimentally determined molar mass and reference material to identify the unknown

volatile liquid you tested.

5. A student failed to vaporize the entire sample prior to placing the test tube in the ice bath. How

did this error affect the calculated molar mass? Justify your answer using calculations.

6. A different student failed to dry the outside of the test tube prior to massing it in Step 12. How

did this error affect the calculated molar mass? Justify your answer using calculations.

7. How would your calculated molar mass have been affected if you had used twice the initial

amount of the unknown compound?

Determination of the Rate of a Reaction, Its Order, and Its Activation Energy

Reaction kinetics is defined as the study of the rates of chemical reactions and their

mechanisms. Reaction rate is simply defined as a change in a measurable quantity divided by the

change in time. In chemistry, the “measurable quantity” is usually molar concentration or

absorbance. Consider the generalized chemical reaction equation A + B C + D. Symbolically it

can be represented in multiple ways:

Note that the units on rate are always M/time which can also be expressed as M time−1. The

negative sign on the first expression indicates that the molar concentration of reactant A will

decrease as time goes by. The second expression is simply the differential rate law expression

where the rate constant k, and the order of reactant A (the exponent m) must be experimentally

determined. Never, ever forget that the value of k is temperature dependent. Since two reactants

are present in our example reaction we can write comparable expressions for reactant B, but beware

that the order of B will not necessarily be the same as the order for A, so we often use a different

variable, such as n, for the exponent on B. The differential rate law can be integrated to link changes

in concentration with time as opposed to rate. This sounds way more complicated that it really is!

(“Integrated” is a Calculus term you need not worry about in this course—we will linearlize the data

to avoid Calculus since it is not a prerequisite for AP Chemistry.) In this experiment you will

investigate the reaction of crystal violet with sodium hydroxide. Crystal violet, in aqueous solution,

is often used as an indicator in biochemical testing. The reaction of this organic molecule with

sodium hydroxide can be simplified by abbreviating the chemical formula for crystal violet as CV.

CV+(aq) + OH–(aq) → CVOH(aq)

As the reaction proceeds, the violet-colored CV+ reactant will slowly change to a colorless

product, following the typical behavior of an indicator. The color change will be precisely measured

by a colorimeter (see Figure 1) or spectrophotometer set at 565 nm (green) wavelength. You can

assume that absorbance is directly proportional to the molar concentration of crystal violet

according to Beer’s law.

Figure 1

The rate law for this reaction is in the form: rate = k[CV+]m[OH−]n, where k is the rate

constant for the reaction, m is the order with respect to crystal violet (CV+), and n is the order with

respect to the hydroxide ion. Since the hydroxide ion concentration is much more than the

concentration of crystal violet, [OH−] will not change appreciably during this experiment. This

technique is often referred to as “swamping”. Thus, you will find the order with respect to crystal

violet (m), but not the order with respect to hydroxide (n). Therefore, the rate constant you will

determine is a pseudo rate constant.

A[A]mRate k

time

You will use integrated rate law methods to determine the order m and the value of the rate

constant by graphing the absorbance and time data that you collect. Set up your axes so that time is

always on the x-axis. Plot the absorbance of CV+ on the y-axis of the first graph. Plot the natural

log of the absorbance of CV+ (ln [CV+], NOT log[CV+]) on the y-axis of the second graph and the

reciprocal of the absorbance of CV+ on the y-axis of the third graph. You are in search of the best

linear fit. Here comes the elegant part… If you do the set of graphs in this order with the y-axes

being “concentration”, “natural log of concentration” and “reciprocal concentration”, the

alphabetical order of the y-axis variables leads to orders of 0, 1 and 2 respectively for CV+. You

can then quickly derive the integrated rate law equations using y = mx + b.

Another important part of the kinetic analysis of a chemical reaction is to determine the

activation energy, Ea. Activation energy can be defined as the energy necessary to initiate an

otherwise spontaneous chemical reaction so that it will continue to react without the need for

additional energy. An example of activation energy is the combustion of paper. The reaction of

cellulose and oxygen is spontaneous, but you need to initiate the combustion by adding activation

energy from a lit match.

We can use a different graphical analysis method to easily determine the activation energy of

a chemical reaction. Each laboratory group will simply repeat the reaction between crystal violet

and sodium hydroxide at a temperature other than room temperature, while keeping the initial

concentrations of the reactants the same for each trial. Recall that the value of k is temperature

dependent. Class data will be collected, graphed and analyzed as follows:

Zero order

k = negative slope

First order

k = negative slope

Second order

k = the slope

ln k

Ea = −R × slope

OBJECTIVES

In this experiment, you will

React solutions of crystal violet and sodium hydroxide at different temperatures.

Graph the concentration-time data and use integrated rate law methods to determine the order

of CV

and the value of a pseudo rate constant, k, for the reaction.

Measure and record the effect of temperature on the reaction rate and rate constant.

Calculate the activation energy, Ea, for the reaction.

Figure 1

MATERIALS

Data collection device 0.10 M NaOH solution

computer or handheld 2.5 × 10–5 M crystal violet solution

colorimeter or spectrophotometer plastic cuvettes

temperature probe or thermometer Beral pipettes

cups or beakers

ice or hot water bath

PROCEDURE

1. Obtain and wear goggles.

2. Set up the data collection system.

Calibrate your spectrophotometer or your colorimeter. We will be collecting data

using the 565 nm (Green) setting.

d. Connect a temperature probe to your device.

e. If using a spectrophotometer, you will need to manually record data every 5

seconds. If using a colorimeter set the program to generate a time graph with 3

seconds between samples for 60 samples.

Do NOT start data collection until Step 3 b.

3. This first trial will be performed at room temperature. You have been given a pipette of 0.10 M

NaOH (colorless solution) and a pipette containing an equal quantity of 2.5 × 10–5 M crystal

violet (purple solution).

a. Simultaneously squirt both solutions into a beaker or cup. Use the tip of the temperature

probe to stir the mixture. Record the temperature of the mixture.

b. Rinse the cuvette with the mixture, discard the rinse into the sink, refill the cuvette at least

2/3 full and place it correctly in the colorimeter and start data collection for the first trial.

c. Once the trial is finished, discard the reaction mixture into the sink.

4. Analyze your data using graphical methods as explained in the introduction. Determine the

order of CV+ and value of the rate constant, k. Record the value of the rate constant for this trial

before proceeding to the next step.

5. Repeat Steps 3-4, using the second set of pipettes that have been sitting in the water bath at your

station.

6. Record your data in the class data table that your teacher has displayed. For your own lab report,

mark your group number with an asterisk *.

CLASS DATA TABLE

Trial 1

Room Temperature (°C)

Rate constant, k (supply appropriate units)

Trial 2 Temperature

(°C)

Rate constant, k (supply appropriate units)

Group 1

Group 2

Group 3

Group 4

Group 5

Group 6

Group 7

Group 8

Group 9

Group 10

Group 11

Group 12

Group 13

Group 14

Group 15

Group 16

PRE-LAB QUESTIONS

1. Refer back to generalized chemical reaction presented in the introduction and write two

comparable rate expressions for reactant B.

2. Compare the molar concentration of the crystal violet solution to that of the sodium hydroxide

solution. Approximately how much more concentrated is the sodium hydroxide? Justify your

answer.

3. Why do we set the spectrophotometer or colorimeter to a wavelength of 565 nm or “green” in

order to measure the absorbance of crystal violet?

4. A student mixes 3.00 mL of 2.27 × 10−5 M crystal violet solution with 3.00 mL 0.1 M sodium

hydroxide both at 24.5 °C and collects the following data:

(a) Describe the graphical analysis steps the student should perform in order to determine

the

(i) Order of the reaction with respect to crystal violet

(ii) Value of the rate constant, k

(b) Use a graphing calculator or computer graphing software to determine the order of the

reaction with respect to crystal violet. Justify your answer.

(c) Write the law expression for this reaction. Justify your answer.

(d) Determine the value of k including its units. Justify your answer.

(e) Calculate the half –life of the reaction. Include units with your answer.

(f) Determine the absorbance of crystal violet at 3.00 minutes. Justify your answer.

(g) Determine the time at which the absorbance of crystal violet is equal to 0.060.

5. The student repeats the experiment at 32.5 °C with the same initial quantities and molarities of

crystal violet and sodium hydroxide. Predict whether the value of the rate constant k will be

increase, decrease or remain unchanged. Justify your answer.

Time

(min) Absorbance

1.49 0.206

4.17 0.157

6.22 0.131

8.15 0.108

10.10 0.093

13.00 0.072

14.02 0.066

16.18 0.055

19.00 0.044

POST-LAB QUESTIONS AND DATA ANALYSIS

1. Graph the class data and calculate the activation energy, Ea, for the crystal violet and sodium

hydroxide reaction.

2. Extrapolate your graph to predict the value of the rate constant k for this reaction at 40 °C.

3. A well-known approximation in chemistry states that the rate of a reaction often doubles for

every

10°C increase in temperature. Use your data to verify or refute this approximation.

4. A student failed to fill the cuvette 2/3 full with the reaction mixture. What effect does this error

have on the measured absorbance values?

AP Chem – Determining the Ka of a Weak Acid – Inquiry Lab

Background Information

, which we will do in the

next chapter.

Prelab Questions

1.

2.

3.

Partial Procedure:

1. Obtain a sample of acetic acid, approximately 125 mL.

2. Use a 25 mL graduated cylinder to place 25 mL samples into four 100 or 250 mL

Erlenmeyer flasks.

You will use these four flasks to run two trials of creating a half-neutralized acetic acid

solution. You have phenolphthalein, dropper bottles of sodium hydroxide, and pH meters

available. You are not allowed to use any equipment to stir, such as stir bars or magnetic

stirrers.

Write the rest of your procedure to create the half-neutralized solutions. Failure to write the

procedure will result in you sitting out for the lab and receiving a zero.

Safety

Use your procedure to determine safety hazards, and look up safety information regarding all

chemicals.

Data

Create a neat and organized data table. Data tables should only include measurements, not

calculations.

Postlab Questions/Calculations

1. What is the pKa of acetic acid?

2. What is the Ka of acetic acid?

3. Calculate your percent error, based on the literature (Accepted) value of Ka for acetic acid.

Use any acceptable reference material to determine the literature value.

4. How would your Ka value differ if you over-titrated with sodium hydroxide? Thoroughly

explain your reasoning.

5. How would your Ka value differ if you added too much phenolphthalein? Thoroughly

explain your reasoning.

6. How is the strength of an acid compare to the Ka value?

Be sure to include conclusion and sources of error.

AP Chemistry Lab 9 – Buffers

Background

Objective

The purpose of this experiment is to study the properties of buffer solutions. An ideal buffer

solution consisting of a weak acid and its conjugate base will be made. The initial pH of the buffer

will be determined, then you will add either a strong acid or strong base to the buffer in a series of

steps. The pH will be measured after each addition. The measured pH values will be compared to

the calculated pH values.

You will also create a buffer of a very specific pH.

Prelab Questions

3. Calculate the ratio of 0.1 M sodium acetate solution and 0.1 M acetic acid solution needed

to make 50 mL of a buffer solution with a pH value of 5.00. Use this ratio to determine the

correct volumes of each solution needed to prepare the solution. {Hint – the ratio of volumes

if equal to the ratio you solved for}

Equipment and Chemicals

0.1 M acetic acid 0.2 M HCl or NaOH beaker graduated cylinder

0.1 M sodium acetate pH probe Magnetic stirrer

1 mL beryl pipette

Procedure

Part 1

1. Obtain 25.0 mL of the 0.1 M acetic acid solution in a 25-mL graduated cylinder. Transfer

the acetic acid solution to a 100-mL beaker.

2. Rinse the 25-mL graduated cylinder with deionized water. Obtain 25.0 mL of the 0.1 M

sodium acetate solution with the 25-mL graduated cylinder.

3. Transfer the sodium acetate solution to the beaker containing 25 mL of the acetic acid

solution. Stir the solution.

4. Set-up a pH meter and place the probe in the beaker. Record the pH of the buffer in your

data table.

5. Set the 100-mL beaker containing the acetic acid–acetate buffer solution on a magnetic

stirrer, if one is available. Add a stir bar to the solution. Gently stir the buffer solution. If a

magnetic stirrer is not available, gently swirl the solution.

6. Obtain approximately 30 mL of 0.2 M HCl or NaOH (as instructed by your teacher) solution

in a clean 50-mL beaker and label the beaker. The lab procedure will henceforth refer to this

as the “strong solution.”

7. Note in your data table which strong solution you are using.

8. Using a graduated Beral-type pipet, transfer 1.0 mL of the 0.2 M strong solution to the acetic

acid–acetate buffer solution.

9. Record the pH of the solution in the Data Table. (1 mL of 0.2 M strong solution added)

10. Using the same graduated pipet, transfer another 1.0 mL of the 0.2 M strong solution

solution to the acetic acid–acetate buffer solution. Record the pH of the solution in the Data

Table.

11. Repeat steps 10, recording the pH in the Data Table after each 1.0 mL of strong solution is

added, until a total of 10.0 mL of strong solution has been added to the solution.

12. Remove the pH probe from the solution, and place it in the rinse beaker.

13. Rinse all solutions down the sink.

Part 2

1. Mix the volumes of acetic acid and sodium acetate that you calculated in Prelab Question 3.

The total should be 50 mL. Record the pH in your data table.

Data and Observations

Part 1

My strong solution is 0.2 M _____________

mL of 0.2 M strong solution

added

Measured pH Calculated pH (you will do this

in postlab)

0

1

2

3

4

5

6

7

8

9

10

Part 2

Measured pH of buffer

Calculations

1. Calculate the theoretical starting pH value of your buffer from Part 1 and enter it into the

Data Table. Show work.

2. Calculate the theoretical pH value after each 1-mL addition of strong solution and enter

them into the Data Table. Show work.

Postlab Questions

1. Consider the series of pH values that you measured and calculated.

a. Did your pH values increase or decrease?

b. How would the series be different if you used the other strong solution?

c. How would the series be different if you added the 1-mL increments to pure water?

d. How would the series be different if the buffer was made with 1.0 M solutions?

Conclusion

See your lab report guide for how to write a stellar conclusion!

Error Analysis

Don’t forget to do this section.

AP Chem lab- Acid/Base titration PURPOSES

Part 1: o Standardize a solution of NaOH using KHP.

Part 2: o Determine the molar mass of an unknown weak acid. o Determine the Ka of an unknown weak acid.

BACKGROUND Chemists often have to answer the question “how much of something is in a sample?” If the sample is an acid or a base, this question is usually answered by a titration. Acid-base titrations can be used to measure the concentration of an acid or base in a solution, to calculate the molar mass of an unknown, and to determine the equilibrium constant of a weak acid or base (Ka or Kb). Important vocabulary and concepts:

Standard solution – a solution of known concentration.

Equivalence Point – the point when perfect stoichiometric amounts of acid and base have combined. (If it is a 1:1 acid:base ratio, it is when the moles of acid = moles of base)

The concentration of the standard solution in this lab, NaOH, will be measured by titrating it against potassium hydrogen phthalate (KHP). KHP is a weak acid that can be easily dried and is therefore very pure. KHP has one ionizable hydrogen ion (monoprotic acid)

o KHP is KHC8H4O4

To determine the molar mass of the unknown weak acid, use the formula: MM = grams of acid titrated Moles of acid titrated Remember that the moles of acid equal the moles of base at the equivalence point.

To determine the Ka of the weak acid, remember that at the half-neutralization point , the pH = pKa. You will use a titration curve graph to help you determine this.

PRELAB QUESTIONS

1. Calculate the molarity of a solution of sodium hydroxide, NaOH, if 23.64 mL of the solution is needed to neutralize 0.5632 g of KHP.

2. It is found that 24.68 mL of 0.1165 M NaOH is needed to titrate 0.2931 g of an unknown weak acid to the phenolphthalein end point. Calculate the molar mass of the acid.

3. The following data was collected for the titration of 0.145 g of a weak acid with 0.100 M NaOH as the titrant:

Volume of NaOH added, mL pH

0.00 2.88

5.00 4.15

10.00 4.58

12.50 4.76

15.00 4.93

20.00 5.36

24.00 6.14

24.90 7.15

25.00 8.73

26.00 11.29

30.00 11.96

a. Graph the data in a graph with mL of titrant on the x axis and pH on the y axis.

b. What is the pH at the equivalence point?

c. Determine Ka and pKa values for the acid. Show your work or explain your reasoning.

4. The following acid-base indicators are available to follow the titration. Which of them would be most appropriate for signaling the endpoint of the titration? Explain your answer.

Indicator Acid Form Base Form pH Transition Interval Bromphenol blue yellow blue 3.0 – 5.0 Bromthymol blue yellow blue 6.0 – 7.6 Thymol blue yellow blue 8.0 - 9.6

SAFETY PRECAUTIONS Dilute sodium hydroxide solutions are irritating to skin and eyes. Phenolphthalein is an alcohol-based solution and is therefore flammable. It is toxic by ingestion. Avoid contact of all chemicals with eyes and skin and wash hands before leaving the lab. Wear goggles and an apron. PROCEDURE Part 1 – Standardization of the Sodium Hydroxide Solution

1. Obtain a sample of KHP that has been previously dried and stored in a desiccator (keeps it dry). On an analytical balance, weigh between 0.4 and 0.6 g of KHP in a weighing dish. Record the exact mass of KHP.

2. Transfer the KHP into an Erlenmeyer flask – pour all the solid through a funnel and use water from a wash bottle to rinse all the remaining solid in the weighing dish or funnel into the flask as well.

3. Add about 50 mL of distilled water to the flask and swirl until all the KHP is dissolved. 4. Obtain about 150 mL of NaOH solution. Label it with your group name to use for Part 1 AND Part 2. 5. Clean a 50 mL buret, then rinse it with small portions (about 5 mL each) of the NaOH solution. 6. Fill the buret to above the zero mark with the NaOH solution, then drain some of the NaOH into a waste beaker

to remove any air bubbles in the spout. Your NaOH should be between the 0 and 10 mL marks. 7. Measure the precise starting volume of the solution in the buret and record in your Part 1 Data table as the

“initial volume.” *Note* Remember that volumes are read from the top down in a buret. Include the appropriate number of significant figures. Remember to estimate your last digit.

8. Position the buret over the Erlenmeyer flask so that the tip of the buret is just above the flask. 9. Add 3 drops of phenolphthalein solution to the KHP solution in the flask. 10. Begin titrating by adding 1.0 mL of the NaOH solution to the flask, then closing the buret valve and swirl the

flask. 11. Repeat step 12 until 15 mL of the NaOH solution have been added to the flask. Be sure to swirl after each

addition. 12. Reduce the incremental volumes of NaOH solution to 0.5 mL until the pink color starts to persist. Reduce the

rate of addition of Naoh solution to drop by drop until the pink color persists for at least 15 seconds (considered a permanent change). Remember to constantly swirl the flask and rinse the walls of the flask with distilled water if necessary until the endpoint is reached.

13. Measure the volume of NaOH remaining in the buret, estimating to the nearest 0.01 mL. Record this as the “final volume” in the Part 1 Data Table.

Part 2 – Determine the Molar Mass and Ka of a Weak Unknown Acid

1. Set up a pH meter and set the meter in a labeled beaker of distilled water. 2. On an analytical balance, weigh between 0.4 g and 0.6 g of the unknown acid in a weigh dish. 3. Using the same method as in Part 1, transfer the acid to a 250-mL beaker and add approximately 100 mL of

distilled water. Add 3 drops of phenolphthalein. 4. Refill the buret with the now standardized NaOH solution. Record the initial volume as the “initial buret reading”

in the Part 2 data table. 5. Set the beaker containing the unknown acid solution on a magnetic stirrer (hot plate) and drop in a magnetic stir

bar. Place the pH meter in the beaker so it is submerged in the acid solution. Be sure the stir bar will not hit the pH meter. Set the stir bar gently spinning. Do NOT turn on the heat.

6. When the pH reading has stabilized, record the initial pH of the solution in the Part 2 Data Table. 7. Add about 1 mL of sodium hydroxide solution to the beaker. Record the exact buret reading in the Part 2 Data

Table. 8. Record the pH of the solution next to the buret reading in the data table. 9. Record the color of the solution in the data table. 10. Add another 1-mL increment of sodium hydroxide solution. Record the buret reading, color, and pH in the data

table. 11. Continue adding NaOH in 1-mL portions. Record both the buret reading, color, and the pH in the data table. 12. When the pH begins to increase by more than 0.3 pH units after an addition, decrease the portions of sodium

hydroxide added to about 0.2 mL. 13. Continue adding sodium hydroxide in 0.2 mL increments and record the buret reading and the pH reading after

each addition. 14. When the pH change is again about 0.3 pH units, resume adding sodium hydroxide in 1-mL increments.

Continue to record the buret reading, color, and pH after each addition. 15. Stop the titration when the pH of the solution is greater than 12. Record the final volume of the solution, the

color, and the final pH. 16. Graph the data, with pH on the y-axis and volume of NaOH on the horizontal axis. Make the graph LARGE

ENOUGH to fill an entire page of graph paper. This will increase the accuracy of your final determinations. You took great care to carefully measure your volumes and pH values – a large graph reflects that care.

Dispose of all solutions down the sink and flush with lots of water. Rinse all glassware with water and return to its original location. DATA TABLE Part 1 – Standardization of Sodium Hydroxide Solution

Mass KHP, g

Initial Buret Volume, mL

Final Buret Volume, mL

Volume of NaOH added, mL

Molarity of NaOH: ________________ M

Part 2 – Molar Mass and Ka

Mass of Unknown Acid

Standard NaOH Concentration

Initial Buret Reading

Initial pH

Buret Reading (mL)

pH Color Buret Reading (mL) Con’t

pH Color

RESULTS, POST-LAB CALCULATIONS AND QUESTIONS

CALCULATIONS 1. From the standardization data in Part 1, calculate the molarity of the sodium hydroxide solution. Enter the

value into the Part 1 Data Table. 2. Consider your Part 2 data. At what buret reading did you reach your equivalence point, according to the

indicator color change? 3. Using your answer from the previous question, calculate the molar mass of the unknown acid. 4. How would the molar mass be different if the unknown acid was diprotic? 5. Why must the KHP and solid unknown acid samples be dried? If they are not dried, how would the results

change (high or low)? 6. Why must NaOH be standardized? Why can’t an exact solution of NaOH be prepared? (You may need to

research NaOH to find out why.) 7. From the graph of pH versus volume of NaOH, determine the pKa of the unknown acid. Convert this value

to Ka.

8. Why is the equivalence point in the titration of the unknown acid with sodium hydroxide not at pH 7?

Virtual Lab: Determining Ksp

1. Go to the following website for a solubility simulation: http://phet.colorado.edu/en/simulation/soluble-salts

a. If you cannot get the link to work, simply Google “PHET”, select the first option, click “Play with Sims”, look under the “Chemistry” tab and find “Salts and Solubilty”.

2. Under “Slightly Soluble Salts” select Silver Bromide. You must use the simulator and create a data table involving at least 3 different volumes of water. 3. Choose a volume of water. Then grab the shaker with the mouse and shake some amount of the AgBr into the container. Make sure to add enough so that there is some undissolved solid remaining. 4. Record the number of dissolved particles of Ag+ and Br- in your data table along with the volume you have chosen. 5. Click the “Reset All” button and repeat steps 3 and 4 using different volumes of water. 6. Use the different volumes of water to determine the solubility of AgBr in each (that is, the number of moles that will dissolve in a particular volume). [You should have 3 different trials at this step.] 7. Use the information you collected in Step 4 to calculate the Ksp value for AgBr of each trial. Then average your results and report this as your Ksp value. (Be aware that the ion measurement numbers in the simulator are for individual particles and you need to be in molarity for Ksp.) 8. Repeat steps 2-7 using strontium phosphate. 9. Under the "Table Salt" tab, choose a single volume quantity and determine the Ksp value for NaCl. 10. Using the "Design a Salt" tab, determine the minimum volume required for 340 particles of cadmium hydroxide (Cd(OH)2) to dissolve. (Ksp is 7.2x10-15) Report: Your grade from this assignment will be based on the completion of the following items.

o Well labeled table of data for Silver Bromide. o Well labeled table of data for Strontium Phosphate. o Calculations of Ksp for both AgBr and Sr3(PO4)2. All steps of the calculations must be o shown. o Data and calculations showing the Ksp value for NaCl.

o Minimum volume for Cd(OH)2 solution as described above. All tables should be made using a computer program. Calculations may either be typed or neatly handwritten under the appropriate data.

AP Chemistry Lab 11 – Redox Titration of Bleach

Background – Read this or you will be LOST

In this lab, a sample of commercial bleach (non-safe bleach) will be diluted, and then titrated with a

solution of sodium thiosulfate to determine the concentration of the bleach. A titration is an

analytical process to determine the concentration of an unknown solution. This is accomplished by

measuring the amount of a solution of known concentration to a known volume of a solution that

has an unknown concentration. Since most bleaches contain the hypochlorite ion (ClO-) in the form

of NaCIO, a strong oxidizing agent, a solution of Kl will be added to allow the hypochlorite ion to

oxidize the iodide ion to iodine. The equation for this reaction is:

The iodine produced in this reaction will be titrated in an acidified solution (1M sulfuric acid) with

sodium thiosulfate.

These two equations give a mole ratio of one mole of hypochlorite ion to two moles of sodium

thiosulfate.

The deep yellow color of iodine will slowly turn a light-yellow color as the iodine molecules are

converted to iodide ions. At this point, the starch indicator is added. Starch turns a deep purple color

in the presence of iodine. The starch solution must be added near the end of the titration, since the

iodine-starch complex will dissociate slowly, and if it is formed too early, the endpoint of the

titration might be missed. The solutions are titrated until the blue color disappears.

The amount of thiosulfate required to reduce the iodine molecules to iodide ions is used to calculate

the amount of bleach in the original sample. An example of the work is shown on the next page to

help you understand.

Example

If 17.55 ml of 0.100 M sodium thiosulfate is required to titrate 25 ml of dilute bleach solution

(prepared by the method explained later), calculate the percent by mass of bleach. Assume the

density of bleach to be equal to 1.06 g/ml.

Since the bleach was diluted, the equation M1V1 = M2V2 , can be used to calculate the molarity of

the concentrated bleach.

Now that we have the concentrated bleach, the mass of NaCIO can be calculated. As stated in

procedure step one, 5 ml of concentrated bleach was originally pipeted to make the aliquot (which is

a diluted solution used for testing that came from a concentrated solution).

For the density, we have the original 5 ml pipet of concentrated bleach that was diluted. From

above, we have the grams of NaCIO that must of been in that pipet. The first step is to calculate the

grams of bleach. Then calculate the percentage.

Objective

To determine the concentration of the hypochlorite ion in commercial bleach.

Prelab Questions

1. What is a titration?

2. Why is the starch indicator added near the end of the titration and not at the beginning?

Use the following information for questions 3-5.

A drain cleaner is to be tested for its acid content (sulfuric acid). A 5.0 ml portion of the drain

cleaner is pipeted into a 100.0 ml volumetric flask and diluted. A 10.0 ml portion of that acid is then

titrated with 0.55 M sodium hydroxide. It is found that 23.75 ml of the sodium hydroxide is

required to neutralize the acid.

3. What is the molarity of the dilute acid solution?

4. Calculate the molarity of the commercial drain cleaner.

5. If the density of the drain cleaner is 1.45 grams/milliliter, calculate the percent, by mass, of

the sulfuric acid.

Equipment and Chemicals

pipets, 5-ml and 25-ml Volumetric flask, 100-ml buret

ring stand pipet bulbs or pumps Erlenmeyer flasks 125 or 250 ml

bleach solution sodium thiosulfate (0.100M) starch indicator solution

1M sulfuric acid solution potassium iodide buret clamp

Procedure

1. Using proper pipetting techniques transfer exactly 5.0 ml of bleach solution to a 100 ml

volumetric flask. Fill the flask with water to the required mark, place a stopper on the flask,

and mix well.

2. Mass 1 gram of solid Kl.

3. Pipet 5 ml of dilute bleach solution, into an Erlenmeyer flask. Place the Kl from step two in

the flask. Add approximately 25 ml of distilled water, and 5 ml of the 1M sulfuric acid. The

solution should be a deep yellow color at this point.

4. Pour the sodium thiosulfate through a funnel into the top of the buret. Record the initial

buret reading. Titrate with the standard 0.1 OOM sodium thiosulfate solution.

5. As the titration is proceeding, the color of the solution in the flask will turn a very light

yellow. At this point, add eight drops of the starch solution. The solution should turn a dark

bluish purple color.

6. Titrate, drop by drop, until all the blue color disappears.

7. When the color just disappears (be sure to check by using a white piece of paper as a

background) record the final buret reading.

8. Refill the buret with more sodium thiosulfate solution. Repeat the titration at least twice

more, using the same experimental setup. Your values of volume of sodium thiosulfate used

should agree to within 5% (average relative deviation) of the mean value. If your figures do

not fall within the required range, continue to titrate until two successive trials do agree to

within 5%.

9. Look at the bleach bottle and record the percentage of NaClO or ClO-, depending on how it

is reported on the label.

Data and Observations

Trial Number Volume of sodium thiosulfate

solution used

1

2

3

4 (if necessary)

5 (if necessary)

Percentage of NaClO OR ClO-

(circle one)

Calculations

1. Calculate the molarity of the diluted bleach in each of the trials.

2. Calculate the molarity of the concentrated bleach in each of the trials.

3. Look at the bleach bottle. If the active ingredients list the percentage of NaCIO, calculate the

percentage of NaCIO in the bleach. If the active ingredients list percentage of CIO-1 then

calculate the percentage of CIO-1 in the bleach solution for each of the trials.

4. Calculate your percent error.

Postlab Questions

1. We used an "aliquot", or a diluted fraction of the initial concentrated solution. What benefit

is there to using an aliquot?

2. How would using 2 grams of Kl affect the results of the concentration of the bleach?

Explain.

3.

Conclusion & Source of Error

See your lab report guide for how to write a stellar conclusion and sources of error!

Photoelectron Spectroscopy (PES)

Unit: Electronic Structure

Knowledge/Understanding:

Photoelectric effect

How photoelectron spectroscopy works

Skills:

Determining electron configuration from photoelectron spectra

The Photoelectric Effect photoelectric effect: the observation that photons short-wave (visible or ultraviolet) light can cause it to emit

electrons. This is an example of ionization, with the photons providing the ionization energy. When visible or ultraviolet light is shone on a substance, the energy from the photons of light excite electrons in the substance. If the energy exceeds the ionization energy of an atom in the substance, the electron is emitted.

The emitted electrons are called photoelectrons.

Photoelectron Spectroscopy photoelectron spectroscopy or photoemission spectroscopy (PES): using the energy from electrons emitted via

the photoelectric effect to gain information about the electronic structure of a substance. The term “photoelectron spectroscopy” is generally used for the technique when applied to gases, and “photoemission spectroscopy” is generally used for electrons emitted from solid surfaces. In photoelectron (or photoemission) spectroscopy, a substance is bombarded with photons, which have a given amount of energy based on their frequency:

hEphoton

By measuring the kinetic energy of the emitted electron, Ek, and predetermining the “work function”, Wo, of the substance (the amount of additional energy it takes to move the delocalized electron to the surface of the material—zero for gases, but nonzero for solids), we can calculate the binding energy, EB, of the electron from Einstein’s equation for the photoelectric effect:

oBk WEhE

The photons used for photoelectron spectroscopy range from ultraviolet light to X-rays. Ultraviolet photoelectron spectroscopy (UPS) and extreme ultraviolet photoelectron spectroscopy (EUPS) are used to study valence electrons and the electrons that participate in chemical bonding. X-ray photoelectron spectroscopy (XPS) is most often used to study core electrons, particularly in solids.

Interpreting Photoelectron Spectra The XPS spectrum for gold looks like the following:

Notice that the peaks for 4s, 4p, 4d, and 4f are different heights. The height of the peak is proportional to the number of electrons in a given sublevel. Analysis of UV photoelectron spectra for a single element with a relatively low atomic number is straightforward. For example, the following is an idealized plot of the photoelectron spectrum for lithium:

The x -axis has units of binding energy of the electron, usually electron-volts (eV). Recall that electrons in higher sublevels have less binding energy, which makes them easier to remove. This means the peak at the left corresponds with the easiest-to-remove electrons (for lithium, this is the 2s electron). The peak at the right corresponds with the electrons that are hardest to remove (for lithium, these are the two 1s electrons). The y -axis indicates the number of photons emitted with that energy. This means the height of each peak is proportional to the number of electrons in the corresponding sub-level. Notice that the peak at 6.26 eV (the 1s electrons) is twice as high as the peak at 0.52 eV (the 2s electron). This means there are twice as many electrons in the 1s sublevel as in the 2s sublevel. The only element that matches this spectrum is lithium (1s2

2s1).

Sample Problem: Q: Identify the element represented by the following photoelectron spectrum:

1s 2s

PES Online Lab:

Answer the following questions on separate sheets of paper or type and print your answers.

Use complete sentences, as if you were answering a real AP FRQ.n

Coulomb’s law states the force of attraction, and thus energy, is directly proportional to the size of the charges and inversely proportional to the distance between the charges.

1. What parts of an atom are considered the “charges” stated in Coulomb’s law and why do they attract each other?

2. What part of the atom is the “distance” stated in Coulomb’s law? 3. Are there parts of an atom that repel each other? Explain your reasoning.

On a Flash capable device, navigate to http://www.chem.arizona.edu/chemt/Flash/photoelectron.html. Click on the periodic table box for hydrogen to get a model photoelectron spectrum for hydrogen to show in the graph area.

4. What is the independent variable, and what should the label be for that variable? 5. What is the dependent variable? 6. What does the number “1.31” represent, specifically? 7. What does the number “1” over the peak represent?

Click on the periodic table box for helium.

8. What does the number “2.37” represent, specifically? 9. What does the number “2” over the peak represent? 10. Does this mean it is easier or harder to remove electrons from helium than from hydrogen? 11. In a well written sentence, explain why it is less energy to remove an electron from hydrogen than from

helium. Note: to correctly answer these types of questions on this assignment be sure to compare and contrast both elements. Again, your answer must mention both elements or it cannot be correct!

Click on the periodic table box for lithium.

12. What is the electron configuration for lithium? 13. How does the electron configuration for lithium compare to the PES graph for lithium? 14. Which electron(s) take the lower amount of energy to remove, the 1s2 or the 2s1? 15. Explain why it takes less energy to remove the 2s1 electron.

Click on the periodic table box for beryllium.

16. Study the new graph carefully and explain what happened. 17. What is the electron configuration for beryllium? 18. What is the numerical amount of the energy associated with the 1s2 electrons? 19. What is the numerical amount of the energy associated with the 2s2 electrons? 20. Thoroughly explain how you used Coulomb’s law to help you answer questions 18 and 19 giving

specific references to the graph. Click on the periodic table boxes for boron, carbon, and nitrogen.

21. What happens to the amount of energy for each electron as you progress from boron to carbon to nitrogen?

22. What is the reason it takes more energy to remove electrons from boron to carbon to nitrogen? Click on the periodic table boxes for boron (again), carbon (again), nitrogen (again), and oxygen.

23. What happens to the energies between nitrogen and oxygen that does not fit the trend? 24. Explain why it is less energy for the 2p4 electrons in oxygen than it was for the 2p3 electrons in

nitrogen when it was expected to increase due to the effective nuclear charge.

Click on the periodic table box for sodium. 25. Study the new graph carefully and explain what happened. 26. Why does sodium easily (really, really easily – which is why it goes boom) lose only one electron to

become Na+? Be sure to specifically reference the PES graph in your answer. 27. What is the electron configuration for sodium? 28. How does the electron configuration show that sodium should easily lose only 1 electron? 29. Sodium can lose all 11 electrons. What would you expect to be the comparative amounts of energy

needed to remove the 1s electrons, 2s electrons, 2p electrons, and the 3 s electrons. State your answer in terms of lowest amount of energy to highest amount of energy.

30. If sodium only naturally loses one electron, what does that show about the amount of energy that can be provided to sodium by its environment?

Click on the periodic table boxes for sodium to magnesium to aluminum to silicon to phosphorus to sulfur to chlorine to argon.

31. What is the one place where the general trend of increasing energy is not followed as the atomic number increases? Explain, again, why the trend is not followed.

32. When a new sublevel is started, where is it found on the graph? Check your answer by clicking from helium to lithium, neon to sodium, and argon to potassium.

Click on the period table boxes for potassium to calcium to scandium.

33. What is unusual about scandium’s d-electron’s peak location? 34. What does this tell you about the distance of the 3d electrons compared to the 4s electrons in

transition metals? Explain your reasoning well using the PES data and Coulomb’s law. 35. If scandium’s last electron is in the 3d sublevel, why does scandium more easily form an oxidation

state of +2? Be sure to explain your answer using the PES data and specifically describe the electrons that get removed to make the +2 oxidation state.

36. After the +2 oxidation state, what would be the next easiest oxidation state to achieve, and as it is not common, what can be inferred about the energy in the environment around scandium?

37. List the sublevels (1s, 2s, 2p, 3s, 3p, 3d, and 4s) for scandium in order from smallest amount of energy to empty to largest amount of energy to empty.

38. Where should the largest jumps in needed energy occur? (For a hint consult the PES graph for scandium)

39. What should be true about the amount of energy to remove electrons from the same energy sublevel, for example, the six electrons in the 3p?

40. Below is an incomplete PES graph for gallium, with the peaks lettered for easy reference. Where would you expect the peak for the 4p1 electron of gallium to appear and what should be the relative size of the height of the peak? Hint: you may wish to use scandium’s PES graph to figure out the identities of the peaks that are shown.

Buffers Lab – AP Chemistry

Disclaimer – you may work in your lab group, but every single answer you write must be your own. If your

answers are too similar to your lab group, it will be considered cheating.

Background

One of the most important applications of acids and bases in chemistry and biology is that of buffers. A

buffer solution resists rapid changes in pH wen acids and bases are added to it. Every living cell contains

natural buffer systems to maintain the constant pH needed for proper cell function. Many consumer products

are also buffered to safeguard their activity.

Objective

The purpose of this experiment is to study the properties of buffer solutions. An ideal buffer solution

consisting of a weak acid and its conjugate base will be made. The initial pH of the buffer will be

determined, then you will add either a strong acid or strong base to the buffer in a series of steps. The pH will

be measured after each addition. The measured pH values will be compared to the calculated pH values.

You will also create a buffer of a very specific pH.

Background Questions

1. How many grams of sodium acetate (molar mass 82 g/mol) must be added to 1.00 L of a 0.200 M

acetic acid solution to form a buffer of 4.20? The Ka of acetic acid is 1.8 x 10-5.

2. Three milliliters of a 2.0 M solution of HCl are added to 1 L of buffer solution containing 0.40 moles

of the weak acid, propanoic acid (Ka = 1.4 x 10-5) and 0.50 moles of sodium propanoate.

a. What is the original pH of the buffer before the strong acid is added?

b. What is the pH of the buffer after the HCl is added? Assume negligible volume change.

Lab Experiment

14. Measure 25 mL of 0.1 M acetic acid solution, transfer it into a 100 mL beaker.

15. Rinse the graduated cylinder (or get a new one) and obtain 25 mL of 0.1 M sodium acetate solution.

Transfer the sodium acetate to the beaker.

Question – what have you created in the beaker? Explain your answer.

Question – if you were to add a strong acid to the beaker, what chemical reaction would take place?

Question – if you were to add a strong base to the beaker, what chemical reaction would take place?

16. Measure the pH of the solution and record your answer here: _______________

Question – what is the theoretical pH of this solution? Show your work below and box your answer.

17. Obtain approximately 30 mL of 0.2 M HCl or NaOH (your choice) solution in a clean 50-

mL beaker and label the beaker. The lab procedure will henceforth refer to this as the

“strong solution.”

Question – which strong solution did you choose? ____________________

18. Using a graduated Beral-type pipet, transfer 1, 2, or 3 mL (your choice) of the 0.2 M strong

solution to the acetic acid–acetate buffer solution.

Question – how many mL of the solution did you add? ____ _mL of __________

I predict that the pH will ________________________ (increase or decrease)

Below, show the work for the theoretical new pH of your buffer. Box your answer.

19. Measure the pH of the solution and record your answer here: _______________

20. Steps 5 & 6 two more times. Add a volume of your strong solution, calculate the theoretical

new pH, then measure the pH.

1st Repeat:

Question – how many mL of the solution did you add? ____ __mL_ of __________

I predict that the pH will ________________________ (increase or decrease)

Below, show the work for the theoretical new pH of your buffer. Box your answer.

Measure the pH of the solution and record your answer here: _______________

2nd Repeat:

Question – how many mL of the solution did you add? ____ __mL_ of __________

I predict that the pH will ________________________ (increase or decrease)

Below, show the work for the theoretical new pH of your buffer. Box your answer.

Measure the pH of the solution and record your answer here: _______________

Postlab Question

1. Calculate the ratio of 0.1 M sodium acetate solution and 0.1 M acetic acid solution needed to

make 50 mL of a buffer solution with a pH value of 5.00. Use this ratio to determine the

correct volumes of each solution needed to prepare the solution. Box your answer. {Hint –

the ratio of volumes if equal to the ratio you solved for}

Virtual Equilibrium Lab

In this experiment you will be introduced to chemical equilibrium. You will then be presented with a number of systems at equilibrium and will be asked to

"stress" these systems by changing the concentration of one of the reactants or products or by changing the temperature of the system.

The experiment is composed of four parts (background, prelab, experiment, postlab) that should be completed in that order.

The experiment can be found at the following link:http://www.harpercollege.edu/tm-ps/chm/100/dgodambe/thedisk/equil/equil.htm

Or at http://goo.gl/DpZZv

Background

Click “background” at the link, and click through each page of the background info. You will keep clicking “continue with…..” until it says “This completes

the background material. Proceed to the prelab.”

This background information is crucial to understanding the experiment. Do not skip it.

Prelab Questions

1. Which describes a solution that contains a system at equilibrium?

circle one (One in which the color of the solution is changing slowly) or (one in which the color is not changing.)

Explain your answer.

2. See the picture and description provided for question two on the website and answer the questions below.

a. What happened to the concentration of each of the ions when the KBr was added?

b. Explain why the solution changed color.

c. Would the tube feel hot or cold when the KBr was added?

3. See the picture and description provided for question three on the website (you need to click “more prelab questions” at the bottom

first”). Suggest a possibility for what chemical could have been added and explain your answer.

4. Methanol has the formula CH3OH and can be produced by the reaction of carbon monoxide with hydrogen gas.

CO + 2 H2 CH3OH + heat

In an attempt to maximize the yield of methanol (amount of methanol produced), a chemist would try to shift the equilibrium as far

to the right as possible. Which of the following would accomplish this? Circle any correct answer.

a. heating the mixture

b. adding an excess of carbon monoxide

c. removing the methanol as it is formed

d. adding a substance that reacts with carbon monoxide

Experiment

In this experiment you will be shown a series of equilibrium systems and asked to apply certain "stresses" to the systems. The results of

applying the stress will then be shown to you. For each system, record:

(a) The chemical reaction (this is given to you)

(b) Your observations for each stressor.This means a color change, a precipitate formed or some like statement. For example, "the red

solution turned green". Do not use this part to state what chemical change occurred. For example, you should not say "more cobalt ion was

formed".

(c) An explanation of what happened chemically. State what chemical reaction occurred or species was formed that resulted in your

observation. For example, "by adding chloride ion to the solution it shifted the equilibrium to the left which caused more CoCl4-2 ion to be

formed. Since this ion is blue, the solution became more blue."

Experiment I

Reaction:

Stressor Observation Chemical Explanation

Experiment 2

Reaction:

Stressor Observation Chemical Explanation

Experiment 3

Reaction:

Stressor Observation Chemical Explanation

Experiment 4

Reaction:

Stressor Observation Chemical Explanation

Experiment 5

Reaction:

Stressor Observation Chemical Explanation

Experiment 6

Reaction:

Stressor Observation Chemical Explanation

Postlab Questions

1. Predict what would happen if the ammonium system described in the experiment was heated.

2. Consider the picture and description given in postlab question 2 on the website.

a. Identify what chemical was added to produce the results shown in the center tube.

b. Identify what chemical may have been added to the center tube to produce the results shown in the right tube.

3. The barium ion is toxic to humans. However, barium sulfate is commonly used as an image enhancer for gastrointestinal x-rays. What

does this imply about the position of the equilibrium shown below?

BaSO4 Ba+2 + SO4-2

4. Hemoglobin (Hb) and oxygen gas form a complex (HbO2) that carries oxygen throughout the human body. Unfortunately, carbon

monoxide also binds to hemoglobin so that an equilibrium is established. Carbon monoxide poisoning occurs when the concentration of HbO2

in the blood is reduced.

HbO2 + CO HbCO + O2

The first aid for a person suffering from carbon monoxide poisoning is to (1) remove them to an area of fresh air, and (2) administer oxygen.

Using the principles of equilibrium, explain how each of these helps to restore the HbO2 concentration.

AP Chemistry Virtual Lab - Colligative Properties

Purpose - to understand the measurements and calculations involved in freezing-point depression and boiling-point elevation labs.

Disclaimer - this is an individual activity. It is not to be done in a group, so your data and answers should all be 100% unique.

Procedure -

Part 1 - Prepare

1. Go to http://goo.gl/uQa8j to access the simulator. If that doesn’t work, google “iowa state colligative properties virtual lab” and click on

the first link called “colligative.”

2. Take a moment to study your options - you can pick different solvents and masses, you can pick different solutes and masses, and you

can select a hot-water bath or cold-water bath, depending on if you will measure boiling point or freezing point.

Part 2 - Freezing Point Sample

3. Select a solvent other than water, and any solute. Change the masses of both substances so that the solvent is greater than 150 g and the

solute is greater than 5 g.

4. Flip the switch at the bottom of the container from hot to cold.

5. Click start. Watch the thermometer reading and watch the sample in the test tube.

6. After a while, the temperature will stabilize. Describe, in your own words, what you see (observation) in the sample and what is actually

happening (explanation).

______________________________________________________________________________________________________

______________________________________________________________________________________________________

7. Using the data shown in the simulator, calculate the change in freezing point (depression) _____________

Part 3 - Boiling Point Sample

7. Select a solvent other than water, and any solute. Change the masses of both substances so that the solvent is greater than 150 g and the

solute is greater than 5 g.

8. Flip the switch at the bottom of the container from cold to hot.

9. Click start. Watch the thermometer reading and watch the sample in the test tube.

10. After a while, the temperature will stabilize. Describe, in your own words, what you see (observation) in the sample and what is actually

happening (explanation).

______________________________________________________________________________________________________

______________________________________________________________________________________________________

11. Using the data shown in the simulator, calculate the change in boiling point (elevation) _____________

Data Collection

Run the simulator six times with six different solvent/solute combinations and random mass combinations. The first three are freezing point,

the last three are boiling point. For each trial, completely fill out the data tables below.

Solvent Mass

Solvent

Solute Mass Solute Normal

FP

Depressed

FP

ΔFP Molality of Solution (show work)

Solvent Mass

Solvent

Solute Mass Solute Normal

BP

Depressed

BP

ΔBP Molality of Solution (show work)

Application

In this final trial, you will be determining the formula of sulfur, which in the simulator is shown as “Sx”. This indicates that sulfur does not

exist as individual atoms but rather as molecules of many sulfur atoms bonded together. By calculating the molar mass of the sulfur molecule,

you can determine the formula.

For this trial, you can either use boiling point or freezing point. You will select a random mass of any solvent (NOT WATER), and a random

mass of sulfur for your solute.

Boiling Point or Freezing Point (circle one)

1. Using the formula for boiling point or freezing point depression, ΔT = kmi, calculate the molality of your solution

2. Calculate the moles of sulfur in your sample

3. Calculate the molar mass of your sample.

4. What is the formula of the sulfur molecules? Round the subscript to the nearest whole number.

Solvent Mass Solvent Solute Mass Solute Normal BP/FP Altered BP/FP ΔBP/ΔFP

Sx

Post-Lab Questions

1. In your own words, write a brief summary of the steps necessary to perform a freezing point depression lab that would gather the data

necessary to calculate the molar mass of an unknown molecular solute. The solute is soluble in water. You can be brief.

2. What measurements must be taken?

3. What calculations must be done?