page ap chemistry 2: matter and intermolecular forces · 10/2/2019 · page 8 of 43 cw 1:...

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Class: AP Chemistry

2: Matter and Intermolecular Forces Topics/ Daily Outline: Day A B Content: TEXT CW HW

1 10/17 10/18 Intermolecular Forces 10.1 1 1

2 10/21 10/22 Solids, Liquids, and Gases 10.1 – 10.7 2 -- 3 10/23 10/24 Kinetic Molecular Theory and Gas Laws 5.1 – 5.9 3 2

4 10/25 10/28 Separation of a Solution -- 4 3

5 10/29 10/30 Separation of a Solution -- 4 -- 6 10/31 11/1 Quarterly Assessment -- -- --

7 11/4 11/5 Spectroscopy and the EM Spectrum 7.1 5 --

8 11/6 11/7 Visible Spectroscopy A18 6 -- 9 11/8 11/11 Percent by Mass of Copper in Brass -- 7 4

10 11/12 11/13 Percent by Mass of Copper in Brass -- 7 --

11 11/14 11/15 Unit Test -- -- -- Important Due Dates:

• Analysis of Hydrogen Peroxide Lab Report, 10/21 (A Day) and 10/22 (B Day)

• Separation of a Solution Lab Report, 11/6 (A Day) and 11/7 (B Day)

• Unit 3 Progress Check (AP Classroom), 11/18 (A Day) and 11/19 (B Day)

• Percent by Mass of Copper in Brass, 11/22 (A Day) and 11/25 (B Day)

For tutorials and additional resources: www.leffellabs.com

If you are absent, use this sheet to determine what you missed and collect the appropriate materials from your teacher. Get help from a friend, the link above, or the instructor.

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HW Assignments

HW 1: Intermolecular Forces Due: 10/21 (A) & 10/22 (B) Complete HW 1 through OWL.

HW 2: Gas Behavior Due: 10/25 (A) & 10/26 (B) Complete HW 2 through OWL.

HW 3: Review for Quarterly Assessment Due: 10/29 (A) & 10/30 (B) Complete HW 3 through AP Classroom.

HW 4: Review for Unit test Due: 11/14 (A) & 11/15 (B) Complete HW 4 through AP Classroom.

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Drills

Date Outcome

Date Outcome

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Date Outcome

Date Outcome

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CW 1: Intermolecular Forces

Directions: 1. Go to LEFFELlabs and download the Intermolecular Forces Presentation from the Unit 2

page. 2. View as a slideshow and progress through the presentation with a partner. Take time to

fully engage with discussion questions with your partner. 3. Once you have completed the presentation, organize your understanding into Cornell

style notes. To do this, review each topic below, and write a question that is answered by the content on the slides. At minimum you should have one question per topic, but you might want to break a topic up into two or more questions if that suits your learning preferences.

a. Polar and nonpolar electron clouds b. London dispersion forces c. Dipole-dipole forces d. Hydrogen bonds e. Other types of intermolecular forces f. Predicting properties using intermolecular forces g. Intermolecular forces in biomolecules

4. One you have formulated your questions, answer them, organizing your responses using the Cornell Notes template of your choice from the options provided by the instructor.

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CW 2: Solids, Liquids, and Gases

States of Matter 1. Go to this website: http://mw.concord.org/nextgen/interactives/. Click on “Phase

Change” and view the simulations below. Experiment with the simulation, making sure to mark two atoms, trace a random atom, and that show attractions is checked on.

a. Molecular View of a Gas b. Molecular View of a Liquid c. Molecular View of a Solid d. Intermolecular Attractions and State of Matter

2. Draw three diagrams that show 5 water molecules in each state of matter. Be sure to

show the relative strength of intermolecular forces occurring in each state.

3. How do you change between different states of matter? For example, if you wanted to

change solid water into water vapor, what would you do?

4. Explain how the addition of heat impacts the kinetic energy of the particles and the impact this has on the strength of the intermolecular forces.

5. Use intermolecular forces to explain why gases take the shape of their container but liquids and solids do not.

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6. How is the motion of particles in gases and liquids different from the motion of particles in solids? How does this relate to the strength of intermolecular attractions in each state?

7. One of the assumptions of the ideal gas law is that gas particles do not attract to each other. Based on your observations, is this true?

8. Why is this assumption acceptable for gases, but not liquids or solids?

9. When water is placed into a graduated cylinder, it forms a concave meniscus. When mercury is poured into a graduated cylinder, it forms a convex meniscus. Use the image below and intermolecular forces to explain why.

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10. For a gas to behave more ideally, should it be heated or cooled? Explain.

11. What factors affect the arrangement and movement of particles in all states of matter?

12. What factors affect the collision frequency and spacing of gas particles?

Structure and Types of Solids Solids can be classified into two categories: crystalline solids, which have a highly regular arrangement of their components, and amorphous solids, which have considerable disorder within their structures. Amorphous solids are composed of atoms or molecules that display only short-range order, due to chemical bonds that hold the solid together. These solids tend to “soften” when heated. When broken, they produce irregular curved surfaces. Almost any substance can solidify in amorphous form if the liquid phase is cooled rapidly enough.

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The positions of the components of a crystalline solid are usually represented by a lattice, a three-dimensional system of points designating the positions of the components (atoms, ions, or molecules) that make up the substance. The smallest repeating unit of the lattice is called the unit cell.

Ionic solids have ions at the points of the lattice that describes the structure of the solid, represented by NaCl. Ionic solids are highly stable substances held together by the strong electrostatic forces that exists between oppositely charged ions. These strong attractions are hard to break, explain why ionic solids have such high melting points and hardness. The structure of ionic solids is best explained by the closest packing of spheres, in which atoms are packed in the most efficient use of space. This maximizes the electrostatic attractions between oppositely charged ions and minimizes repulsions amongst like charged ions.

Molecular solids have discrete covalently bonded molecules at each of its lattice points, represented by sucrose. Ice is a molecular solid that has an H2O molecule at each point. These substances are characterized by strong covalent bonding within the molecules but relatively weak forces between the molecules. The strength of these intermolecular forces depends on the molecules themselves. Large molecules are more easily polarizable, meaning they are more likely to form dipoles, and have stronger intermolecular forces. Therefore, elements like phosphorus (P4) and sulfur (S8) are solids at room temperature, rather than liquids. These compounds are typically soft with low melting points.

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Atomic solids have atoms at their lattice points. These include solid elements like carbon (graphite, diamond, fullerenes), boron, silicon, and metals.

A metallic crystal may be pictured as containing spherical metal atoms packed together and bonded to each other equally in all directions. This structure is modeled in the closest packing arrangement, in which atoms are packed in the most efficient use of space. The electron sea model predicts a regular array of metal cations in a sea of shared valence electrons, meaning that the metal is held together covalently by delocalized electrons. The strength of this attraction dictates the properties of the metal, which is why metals have a large range of melting points. The mobile electrons can conduct heat and electricity, and the metal cations can be easily moved around as the metal is hammered into a sheet or pulled into a wire. Metals can be made into alloys by adding other elements. In the case of a substitutional alloy, some of the host metal atoms are replaced by other metal atoms of a similar size. An interstitial alloy is formed when the holes of the closest packed metal structure are occupied by small atoms.

Another type of atomic solids are network solids. These compounds are held together by covalent bonds, and typically have high melting points and are poor conductors. Examples include carbon and silicon. There are many applications of these solids through ceramics and semiconductors.

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13. Complete the table below comparing types of solids.

Solid Type Amorphous Atomic Metallic

Atomic Network

Molecular Ionic

Structure

Type of Bonding

Conductivity

Relative Melting Points

Malleability

Examples

14. Using the table, classify each of the following substances according to the type of solid it

forms. a. Gold b. Carbon dioxide

c. Lithium fluoride d. Krypton

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CW 3: Gas Behavior

Ideal Gas Law The ideal gas law relates several variables that affect gases in one expression:

𝑃𝑉 = 𝑛𝑅𝑇 R is the gas law constant, which has many possible values. The correct R value is selected by considering the units that need to cancel out to find the answer in the desired units.

Common Gas Law Constant Values

𝑅 = 8.31 𝐿 ∙ 𝑘𝑃𝑎

𝑚𝑜𝑙 ∙ 𝐾 𝑅 =

0.0821 𝐿 ∙ 𝑎𝑡𝑚


62.364 𝐿 ∙ 𝑚𝑚𝐻𝑔

𝑚𝑜𝑙 ∙ 𝐾

Partial Pressures According to Dalton’s Law of Partial Pressures, the total pressure of a mixture of gases is equal to the sum of the partial pressures of the gases.

𝑃𝑔𝑎𝑠 𝐴 = 𝑃𝑡𝑜𝑡𝑎𝑙 × (𝑚𝑜𝑙𝑒𝑠 𝑔𝑎𝑠 𝐴

𝑡𝑜𝑡𝑎𝑙 𝑚𝑜𝑙𝑒𝑠 𝑜𝑓 𝑔𝑎𝑠)

𝑃𝑡𝑜𝑡𝑎𝑙 = 𝑃𝑔𝑎𝑠 𝐴 + 𝑃𝑔𝑎𝑠 𝐵 + 𝑃𝑔𝑎𝑠 𝐶

1. Consider the decomposition of 33 g of potassium bicarbonate at 520 °C and 880 mmHg.

2KHCO3(s) → K2CO3(s) + CO2(g) + H2O(g) a. How many moles of each gas will be created?

b. What volume will the gases occupy?

c. What is the partial pressure of each gas?

2. Calculate the mass of KClO3 required to produce 29.5 L of O2 at 127°C and 760 mmHg. 2KClO3(s) → 2KCl(s) + 3O2(g)

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Kinetic Molecular Theory A model is considered successful if it explains the observed behavior in question and it correctly predicts the results of future experiments. No model can be proven correct, as all models are approximations based on experimental results. Simple models are used to predict approximate behavior, while more complex models are used to account very precisely for observed quantitative behavior. The kinetic molecular theory is a simple model used to explain the properties of ideal gases. This model is based on the following postulates about the behavior of individual gas particles:

A. Gas particles are so small compared with the distances between them that the volume of the individual particles can be assumed to be zero.

B. Particles are assumed to exert no forces on each other; they are assumed neither to attract nor repeal each other.

C. The particles are in constant, random motion. Collisions between the particles and the walls of the container exert gas pressure. Collisions are elastic (no energy lost).

D. The average kinetic energy of a collection of gas particles is assumed to be directly proportional to the Kelvin temperature of the gas.

3. Each of the following are behaviors of gases that contradict the postulates of KMT.

Match each to its postulate.

Behavior Postulate

Water vapor condenses into liquid water.

Gas particles take up space.

Gases can be compressed greatly, but not completely.

Kinetic energy may be lost as heat after gas particles collide.

4. What is the benefit in using KMT to describe gases when there are several gas behaviors

that contradict its postulates?

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The Meaning of Temperature The Kelvin temperature is an index of the random motions of the particles of a gas, with higher temperature meaning greater motion. This relationship is described by:

(𝐾𝐸)𝑎𝑣𝑔 =3

2𝑅𝑇

Where: KE = kinetic energy (J) R = 8.3145 J/mol∙K (gas law constant) T = temperature (K)

5. Consider four samples of gases:

a. How does the average kinetic energy of two different gases at the same

temperature compare?

b. What is the relationship between the average kinetic energy and the temperature of a gas?

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Kinetic Energy and Root Mean Square Velocity When examining the gas molecules individually, the molecules at a given temperature do not all move at the same speed. Each gas molecule has slightly different kinetic energy. To calculate the kinetic energy of a gas particle, we use an average speed of the gas, called the root mean square speed (urms).

𝐾𝐸𝑔𝑎𝑠 𝑝𝑎𝑟𝑡𝑖𝑐𝑙𝑒 =1

2𝑚𝑢𝑟𝑚𝑠

2

Where: KE = kinetic energy (J) m = mass of the molecule (kg) urms = root mean square speed (m/s)

𝑢𝑟𝑚𝑠 = √3𝑅𝑇

𝑀

Where: R = 8.3145 J/mol∙K (gas law constant) T = temperature (K) M = molar mass (kg/mol)

6. Consider three identical flasks filled with different gases.

Flask A: CO (molar mass 28 g/mol) at 760 torr and 0°C Flask B: N2 (molar mass 28 g/mol) at 250 torr and 0°C Flask C: H2 (molar mass 2 g/mol) at 100 torr and 0°C

a. In which gas will the molecules have the greatest average kinetic energy? Explain.

b. In which gas will have the molecule with the greatest kinetic energy? Explain. (HINT: not asking about the average!)

c. In which gas will the molecules have the greatest average velocity? Explain.

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Gases do experience attractions and repulsions (Postulate A)

Gas particles take up space (Postulate B)

Maxwell-Boltzmann Distribution The graph shows the velocity distribution of N2 gas molecules at three temperatures, known as a Maxwell-Boltzmann distribution. The peak of the curve reflects the most probable velocity.

7. Explain what happens to the peak and to the range of velocities as the temperature of a gas is increased.

8. What is the relationship between temperature, kinetic energy, and velocity of a gas?

Nonideal Gas Behavior While the ideal gas law does approximate the behavior of gases, it fails to do so over a wide range of temperatures and pressures. This is because the assumptions of the kinetic molecular theory are only accurate within certain temperature and pressure ranges. Kinetic molecular theory states that gas particles are assumed to take up no volume and that gas particles exert no attractive or repulsive forces on each other. These postulates apply to the P and V terms in the ideal gas law.

𝑃𝑉 = 𝑛𝑅𝑇

9. Make a prediction about how the pressure and volume terms may change from their values at ideal behavior.

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The Van der Waals Equation To come up with such a simple equation to explain complex gas behavior, several assumptions were made about gases that aren’t true. For example, according to ideal gas law theory, a gas can never condense into a liquid, because gas particles do no experience attraction for each other. The ideal has law also treats gas molecules are point particles with perfectly elastic collisions. While this works well for dilute gases, a slightly more complex formula corrects for some of these assumptions. The real gas law (also known as the van der Waals Equation) contains two corrective factors.

(𝑃 +𝑛2𝑎

𝑉2) (𝑉 − 𝑛𝑏) = 𝑛𝑅𝑇

The values of 𝑎 and 𝑏 generally increase with the size and complexity of the molecule, and 𝑏 is generally much smaller in magnitude than 𝑎.

10. Calculate the pressure exerted by 0.500 mol of N2 in a 1.00 L container at 25°C: a. using the ideal gas law.

b. using the van der Waals equation.

Alters the pressure in the ideal gas equation to account for attractive intermolecular forces between gas

molecules. The magnitude indicates the strength of the intermolecular forces. The

constant, 𝑎, has units of 𝐿2𝑎𝑡𝑚

𝑚𝑜𝑙2.

Accounts for the volume occupied by the gas molecules. The constant, 𝑏, has units

of 𝐿

𝑚𝑜𝑙. Since 𝑏 corresponds to the total

volume occupied by gas molecules per mole, it closely corresponds to the

volume per mole of the liquid state, whose molecules are closely layered.

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11. Calculate the pressure exerted by 0.500 mol of N2 in a 10.00 L container at 25°C: c. using the ideal gas law.

d. using the van der Waals equation.

12. Compare the results of Question 10 and Question 11. How does increasing the volume affect the interactions between gas particles? Why?

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Effusion and Diffusion Diffusion is used to describe the mixing of gases. Effusion is used to describe how quickly a gas travels through a pinhole into an evacuated chamber. The rate of effusion depends directly on the velocity (and thus the temperature). This allows us to derive Graham’s Law of Effusion, which states that the rate of effusion of a gas is inversely proportional to the square root of the mass of its particles.

𝐸𝑓𝑓𝑢𝑠𝑖𝑜𝑛 𝑟𝑎𝑡𝑒 𝑓𝑜𝑟 𝑔𝑎𝑠 1

𝐸𝑓𝑓𝑢𝑠𝑖𝑜𝑛 𝑟𝑎𝑡𝑒 𝑓𝑜𝑟 𝑔𝑎𝑠 2=

𝑢𝑟𝑚𝑠𝑓𝑜𝑟 𝑔𝑎𝑠 1

𝑢𝑟𝑚𝑠 𝑓𝑜𝑟 𝑔𝑎𝑠 2=

√3𝑅𝑇𝑀1

√3𝑅𝑇𝑀2

=√𝑀2

√𝑀1

13. The effusion rate of an unknown gas is found to be 31.50 mL/min. Under identical

experimental conditions, the effusion rate of O2(g) is found to be 30.50 mL/min. Which of the following could be the unknown gas: CH4, CO, NO, CO2, or NO2? Explain.

14. One way to separate oxygen isotopes is by gaseous diffusion of carbon monoxide, a process that behaves as effusion. Predict the relative rates of effusion of the following isotopes.

Isotope Mass Number 12C16O 28 12C17O 29 12C18O 30

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CW 4: Separation by Chromatography

Challenge You are working for a crime lab and a dye residue linked to a crime has been turned in for analysis. To identify the dyes in the residue, each will need to be separated from the mixture. Another lab technologist has tried to separate the molecules but was only successful in getting one dye to separate out of the mixture. They have turned to you for help in improving their separation method. You will first attempt their method and then propose and attempt a modification.

Background There are two phases in paper chromatography, a stationary phase (the paper) and a mobile phase (the solvent). A molecule can have a greater affinity for either the paper or for the solvent. For example, consider the separation of molecules using chromatography paper in water. The paper is made of cellulose, a polymer with exposed hydroxyl (OH–) groups that will attract water molecules. This interaction makes a thin layer of water on the paper that competes for the attraction of the molecules being separated. Alternately, the molecule can be attracted to the solvent and travel with the solvent up the paper. When doing chromatography, a small amount of solvent is placed in a sealed container. The container must be sealed so the solvent saturates the paper and does not evaporate first. The mixture being separated is put on a piece of paper, the starting point is marked, and the paper is put into the solvent. The level of separation is measured by a ratio that compares the distance that the molecule travels to the distance the solvent travels, called the Rf value. To find the Rf value, the experimenter must measure the distance that the solvent traveled on the paper and the distance that each molecule traveled. It is best to run the test more than once to reach the best separation values possible. The Rf value is a ratio of the distance of the molecule divided by the distance of the solvent. The greater the distance the molecule travels, the greater its affinity for the solvent and the greater the Rf value. A mixture of solvents is often used to fine tune the separation of a mixture of molecules. The most common solvents and their structures are given in the diagram.

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Introductory Experiment The following procedure was given to you by your colleague at the lab.

1. Obtain and wear goggles. 2. Lightly draw a line with a pencil on a thin layer chromatography slide about 1 cm from

bottom, being careful not to scratch or touch the surface of the plate. These slides are covered with a thin layer of silica, which is an extremely polar surface.

3. Place a drop of mixture in the middle of the line. 4. Label a tall beaker with the name of the solvent to be used: chromatography solution

(petroleum ether and acetone mixture) 5. Place slide inside the beaker, leaning it upright against the side of the glass. 6. Add enough solvent until it rises to just below (but not touching) your dot and the line

drawn on paper. 7. Place a watch glass on top. Leave in container until solvent line has reached near the top

of the plate. 8. Remove and gently mark solvent line. Allow to dry. 9. Measure and record solvent distance, molecule distance to calculate Rf values.

Data Table

Solvent distance

Molecule A distance

Molecule B distance

Molecule C distance

Detailed Observations

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Guided Inquiry Design and Procedure In your groups, discuss the following questions, recording your thinking on a sheet of newsprint.

1. While your initial chromatography experiment runs, watch the video posted here: http://bit.ly/2k8wSc2 using the x1.25 or x1.5 options. Then answer the questions below.

a. What are the three general ideas covered by the video? b. Explain the difference between molecules A, B, C, and D? c. Why is molecule D most attracted to a polar surface? d. Why is cellulose (paper) a polar surface? e. Explain what the mobile phase is. f. Why did A travel the farthest when a non-polar mobile phase was used? g. What two factors influence the distance traveled by each component? h. What data needs to be collected to solve for the retention factor, Rf? i. Where do we measure for each spot? j. What do low and high Rf values mean respectively? k. What items are required to properly report a chromatograph? l. Why might a chemist create a special blend of solvents for a given separation? m. Explain why the second trial of the chromatography simulation failed to achieve

an acceptable separation. 2. Once your first separation is complete, observe the results. Did you achieve a good

separation? What do the results tell you about the affinity of the dyes for the solvent? For the silica (very polar) in the thin layer plate?

3. Write a detailed, step-by-step procedure for separating a mixture of Blue 1, Red 40, and Yellow 5 that improves on the separation in the Introductory Experiment. You should use a blend of the available solvents in a simple ratio. Include all the materials, glassware and equipment that will be needed, and safety precautions that must be followed. You will be given a dye sample (dissolved in water) to analyze.

4. Review additional variables that may affect the reproducibility or accuracy of the

experiment and how these variables will be controlled.

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Handwritten Lab Notebook PRE-LAB PRE-LAB PRE-LAB PRE-LAB PRE-LAB PRE-LAB PRE-LAB

1. Introduction: Explain theory and applications behind the lab, as well as any chemical equations, graphs, diagrams, or mathematical equations that will be required.

o Applications (why do we care?) o Explain how polarity is used to achieve separation of mixtures of molecules.

Discuss how interactions between the molecules and the stationary and mobile phases affect the distance each molecule travels.

o Describe the data that needs to be collected to report a chromatograph, including the calculation of Rf values.

2. Safety Data Sheets: Recreate the table below. Look up the SDS for all chemicals used during the lab and complete the table using the Hazards Identification section.

Chemical Eye Irritant Skin Irritant Respiratory Irritant Other Info

Hexane Ethanol

2-Propanol Acetone

Summary of Safety Precautions:

3. Materials and Methods: Write a detailed step-by-step procedure for the experiment. Include all the materials, glassware and equipment that will be needed, and safety precautions that must be followed.

PRE-LAB PRE-LAB PRE-LAB PRE-LAB PRE-LAB PRE-LAB PRE-LAB

4. Procedural Notes: Record details of the procedure, exact sizes of glassware, concentrations or amounts of chemicals, instructions for using LoggerPro, and any other notes provided by your instructor or group members. Finalize your procedure.

5. Data Collection: Create a data table to neatly record measurements, written with the correct precision based on the equipment, detailed observations, and sources of error.

6. Calculations: Show one neat handwritten worked example for each major required calculation. Include the unrounded answer and the answers with correct sig figs.

o Calculate the Rf for each spot/ dye.

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Typed Lab Report 7. Typed Analysis: Access your class set of data from LEFFELlabs/ Unit 2. Perform all

required calculations using a spreadsheet program such as Excel or Goggle Sheets. Copy and paste into the post lab such that it is readable when the page is printed.

o Calculate the average Rf value for each spot/ dye. o Calculate the standard deviation in the average Rf values. Report your answer as

Rf value ± standard deviation. 8. Typed Conclusion: Include the following in narrative form:

o Discussion of what you did, including the purpose/ goal of the experiment. Describe any changes to the procedure from what you wrote in your lab notebook.

o Describe major findings and make claims with supporting evidence from the lab. Report calculated numbers, averages, and standard deviations as required.

o Comment on error, both inherit error (due to precision of equipment) and experimental/ human error. Remember that “the lab would have been better if we did it correctly” or “human errors were made” are not discussion of error. Errors should be clearly described along with how they impacted the calculations in the data analysis.

o Where you able to improve on the original separation procedure? Provide specific experimental evidence to support your answer.

o Describe another refinement of the procedure that could be attempted next. Justify your refinement using your knowledge of solubilities of each dye and the intermolecular forces present.

9. Typed Cover Sheet: Create a cover sheet using the exact table below. Place cover sheet on top of items 4 to 8 from above.

Report Title: Separation by Chromatography

My Name: Partners:

Class Period:

Due Date:

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CW 5: Spectroscopy and the EM Spectrum

Directions: 1. Go to LEFFELlabs and download the Spectroscopy Presentation from the Unit 2 page. 2. View as a slideshow and progress through the presentation with a partner. Take time to

fully engage with discussion questions with your partner. 3. Once you have completed the presentation, organize your understanding into Cornell

style notes. To do this, review each topic below, and write a question that is answered by the content on the slides. At minimum you should have one question per topic, but you might want to break a topic up into two or more questions if that suits your learning preferences.

a. The flame test lab b. The Bohr model of the atom c. Energy levels are quantized d. The photoelectric effect e. Spectroscopy can be used to determine structure f. Spectroscopy can be used to determine concentration

4. One you have formulated your questions, answer them, organizing your responses using the Cornell Notes template of your choice from the options provided by the instructor.

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CW 6: Visible Spectroscopy

Essential Question How can light be used to study color and determine concentrations of chemical species in solutions?

Background Spectroscopy involves the interaction of electromagnetic radiation with matter. The absorption of electromagnetic radiation results in different types of transitions in a substance, depending on the energy of the radiation. Low energy microwave radiation, for instance, is converted to the energy of molecular rotation. Absorption of infrared radiation excites vibrational frequencies associated with covalent bonds in a molecule. Visible and ultraviolet light cause electron transitions between different electron energy levels in a substance. The energy of electromagnetic radiation is quantized, as are the energy levels associated with various transitions, whether rotational, vibrational or electronic. Furthermore, the energies of these transitions are characteristic of an atom, molecule or compound. As a result, the absorption spectra of substances generally consist of specific lines or bands that can be used as a type of fingerprint to identify a substance. A spectrophotometer is an instrument that uses electromagnetic radiation, such as ultraviolet, visible or infrared light, to analyze the absorption or transmission of radiation by a sample. The basic function of a spectrophotometer is shown below.

In addition to the energy source used in spectrophotometers, a diffraction grating called a monochromator is also incorporated. The monochromator spreads the beam of light into the light’s component wavelengths. The desired wavelength is then focused onto the sample cell to detect any absorption or emission of light by a substance in a sample.

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The electromagnetic spectrum gives the possible wavelengths or frequencies of electromagnetic radiation. In this investigation a visible spectrophotometer will be used - it scans the visible region of the electromagnetic spectrum, from 380 nm to 750 nm. Typical light sources for visible spectrophotometers include xenon and tungsten lamps.

Glass or plastic cuvettes may be used in visible spectrophotometers. More specialized spectrophotometers require specific materials. For example, ultraviolet spectrophotometers require quartz cells which do not absorb ultraviolet radiation. Only seven dyes are approved by the Food and Drug Administration for use in foods, drugs and cosmetics. These seven dyes give rise to the entire palette of artificial food colors. The structure of FD&C Blue 1 is shown. Notice the extensive series of alternating single and double bonds (also called conjugated double bonds) in the center of the structure. This feature is characteristic of intensely colored dyes and pigments. Every double bond added to the system reduces the energy difference between energy levels so that the resulting difference in energy corresponds to visible light. A solution containing FD&C Blue 1 appears blue under normal white light because blue is the color of light transmitted by the solution. The colors or wavelengths of light that are absorbed by this solution are complementary to the transmitted color. A color wheel provides a useful tool for identifying the colors or wavelengths of light absorbed by a substance. A blue solution absorbs orange light, so we expect the visible spectrum of FD&C Blue 1 to contain a peak in the 600−640 nm region. The optimum wavelength for spectrophotometric analysis of a dye solution is generally determined from the wavelength of maximum absorbance, called lambda max (λmax). The value of λmax for FD&C Blue 1 is 630 nm.

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Procedure 1. You will be given a solution of known concentration of Blue 1 to explore the relationship

between light transmitted and concentration. 2. Log on to a laptop. Open Logger Pro on the desktop and connect the SpectroVis. 3. Fill a cuvette to the line with the same solvent used in the sample to be analyzed. 4. Wipe the sides of the cuvette clean using a lint

free wipe. Place the cuvette into the SpectroVis according to the diagram.

5. On Logger Pro, select Experiment, select Calibrate, select Spectrometer 1. Allow the lamp to warm up, then select Finish Calibration, then click OK.

6. Fill a cuvette to the line with the sample to be analyzed. Wipe the cuvette clean and place the filled cuvette into the SpectroVis.

7. Click on the green Collect button on the top, right hand corner. You should see a complete spectrum. When you are done, click on the red Stop button.

8. Record the percent transmittance at 630 nm in the data table. 9. Repeat for any additional solutions.

Data Solution

# Dilution Ratio (mL stock/ mL water)

Molar Concentration (μM)

%Transmittance Transmittance as

a Decimal

1 10 mL/ 0 mL

2 8 mL/ 2 mL

3 6 mL/ 4 mL

4 4 mL/ 6 mL

5 3 mL/ 7 mL

6 2 mL/ 8 mL

7 1 mL/ 9 mL

8 0 mL/ 10 mL

9 Sports Drink

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Analysis Questions 1. Open an excel spreadsheet and transcribe the data. Make labeled columns: Solution #,

Molar Concentration, and T (%T as a decimal). 2. Create a graph with concentration (μM) on the x axis and T (%T as a decimal) on the y-

axis. 3. Is there a relationship between concentration and T? Explain.

4. It would be useful to get a straight line that goes through zero. Scientists try to find linear relationships because such relationships make it easier to identify unknowns and predict outcomes of investigations. It is also helpful to have a positive slope. Generate the following graphs.

a. 1/T versus [dye] b. 1×10T versus [dye] c. logT versus [dye] d. -logT versus [dye]

5. Which graph meets the requirements above? Complete a linear fit of this graph and give the linear equation and R2 value.

6. What information does the R2 value give for each linear fit?

7. In your equation, what variable is represented by Y? X?

8. Obtain and sample of blue sports drink and measure its transmittance. Use your transmittance data and linear fit to find the concentration of blue 1 in the sports drink.

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CW 7: Percent by Mass of Copper in Brass

Challenge A crime scene investigator has collected some brass shell casings at a crime scene. You will determine the percent by mass of copper in the brass casings so that they can be matched to a specific manufacturer.

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, mechanical strength, electrical conductivity, corrosion resistance, etc. Visible spectroscopy provides a simple tool for determining the percent copper in brass. In a visible spectrophotometer, the instrument passes a beam of light into a solution and then detects how much of the light comes out of the other side of the solution. When light is absorbed by the solution in the cuvette, the radiant power (P) of the light beam decreases. Transmittance is the fraction of light (P/P0) that passes through the sample.

The relationships between transmittance and percent transmittance (% T) and between transmittance and absorbance (A) are given below.

%𝑇 = 𝑇 × 100 = (𝑃

𝑃0) × 100

𝐴 = 𝑎𝑏𝑠𝑜𝑟𝑏𝑎𝑛𝑐𝑒 = −log (𝑇)

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In general, absorbance is proportional to concentration. The higher the concentration of colored substance in solution, the more intense its color will be, and the greater its absorbance. The linear relationship between absorbance (A) and concentration (c) is expressed in Beer’s law, where b is the path length in cm and ε (epsilon) is a proportionality constant that is characteristic to each substance. If c is given in units of molarity (M = moles/L), then ε is known as the molar absorptivity coefficient, with units M−1 cm−1.

𝐴 = 𝜀𝑏𝑐 Beer’s law can be used to determine the “unknown” concentration of a substance in solution if its absorbance is measured. The most accurate way to do this is to create calibration curve, which plots absorbance versus concentration for a series of solutions of known concentration. This is achieved by measuring the absorbance of the solution at a single selected wavelength. This gives rise to a straight line that passes through the origin, where y = absorbance, m = x = concentration, and b = 0.

𝑦 = 𝑚𝑥 + 𝑏

Once the calibration curve is generated, the absorbance of a solution of unknown concentration may be measured and used to find its concentration by plugging values into the linear fit of the calibration curve. The optimum wavelength for spectrophotometric analysis of a substance is selected by measuring the entire visible spectrum of the substance, which generates a plot of absorbance (A) versus wavelength (λ, “lambda”). The wavelength that gives the maximum absorbance is selected.

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The color of a solution is an important tool used by scientists to gain information about the composition of the solution. Color is a physical property that is useful for both qualitative and quantitative analysis. Qualitative methods give data about the chemical structure of the components in a solution, as the wavelength of light absorbed by a substance is characteristic of its molecular or electronic structure. The differences in the spectra below are due to structural differences in each molecule, which causes light to interact differently with each molecule. The absorption of visible light by a substance results from electron transitions between energy levels. Both light energy and electron energy levels are quantized, so the specific wavelength of light absorbed by a substance depends on the energy difference between two electron energy levels.

Quantitative methods can be used to determine the concentration of compounds in a solution. The spectra below are both for the same molecule, but one solution is more concentrated. The intensity of light absorbed depends on the amount of the substance in solution. Generally, the more concentrated the solution, the more intense the color will be, and the greater the intensity of light the solution absorbs.

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In this experiment, you will be given a known mass of brass. You will react the brass with concentrated nitric acid, forming copper(II) ions in solution. This reaction must be performed in the fume hood as one of the products, NO(g) is toxic. As the NO(g) is formed, it reacts with oxygen in the air to form the reddish brown NO2(g). The unbalanced ionic equation for the copper reaction is:

Cu(s) + NO3– (aq) → Cu2+(aq) + NO(g) in an acidic solution

The Cu2+ ions in the unknown aqueous solution form the complex ion, [Cu(H2O)6]2+, which causes the blue color. This means that when “white” light (all wavelengths) passes through the solution, the dominant emerging color is blue. A spectrophotometer is used to analyze the color intensity of the copper (II) nitrate solution that forms. For this analysis, it will be necessary to determine which color, with a specific wavelength, will be most strongly absorbed by the copper ions. Once you have selected the wavelength used in your analysis of Cu2+ ions, you will measure the absorbance of several Cu2+ solutions of known concentration and plot the results. By locating the absorbance of the unknown brass solution on the vertical axis of the graph, the corresponding concentration can be found on the horizontal axis. The concentration of the unknown can also be found using the slope of the Beer’s law curve.

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Guided Inquiry Design and Procedure 1. Calculate the %T for a sample with an absorbance of 0.8 and for another sample with

absorbance of 1.6. How do these values compare? What can we say about the amount of light reaching the detector in each case?

2. Explain why absorbance values above 1.5 are unreliable. What should you do if you encounter absorbance values higher than 1.5?

3. A group of students collected the data shown to determine the concentration of a sample. When they measured the absorbance of the sample, they got a value of 1.4. Can they use this data to find the concentration of the sample? Why or why not?

4. When the brass is reacted with nitric acid, nitrogen dioxide (a toxic, reddish-brown gas) is formed as a side product. What safety precautions are needed in this inquiry lab to protect against this hazard, as well as the hazards due to the use of concentrated acid?

5. Examine the spectra given on the next page. a. Based on the spectra, what is absorbing light in solution, the metal ion or the

polyatomic ion? How do you know? b. Do Zn2+ ions absorb visible light? Discuss the answer in terms of (a) the color and

appearance of Zn2+ aqueous solutions and (b) the electronic structure of Zn2+ ions. Hint: See the Background section for information on the electronic transitions in atoms.

c. Identify a suitable wavelength for analysis of Cu2+ ions in aqueous solution. Recall that absorbance measurements are most accurate in the range 0.1 to 1.0.

d. If Cu2+ ions and Fe3+ ions are both suspected of being present in the solution, will Fe3+ ions interfere with the analysis of Cu2+ at the wavelength selected? Why or why not?

6. In order to accommodate the range of possible Cu2+ ion concentrations that may be obtained by dissolving brass, it’s recommended that the calibration curve cover several concentrations between 0.05 M to 0.4 M. Calculate the volumes of 0.4 M Cu(NO3)2 stock solution and water required to prepare 8.0 mL of each standard solution (0.05 M, 0.10 M, 0.20 M, 0.30 M, 0.40 M) for your calibration curve.

7. What is the purpose of “blanking” the spectrophotometer? 8. Write detailed, step-by-step procedures for preparing the standard solutions and

obtaining the calibration curve data. This include dissolving the brass in nitric acid and how to create the standard solutions (See Question 6). Include all the materials, glassware and equipment that will be needed, and safety precautions that must be followed.

9. Review additional variables that may affect the reproducibility or accuracy of the experiment and how these variables will be controlled.

Concentration Absorbance 0.10 M 0.25

0.20 M 0.50

0.30 M 1.00 ? 1.40

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Sample Spectra for Question 5

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Handwritten Lab Notebook PRE-LAB PRE-LAB PRE-LAB PRE-LAB PRE-LAB PRE-LAB PRE-LAB

1. Introduction: Explain theory and applications behind the lab, as well as any chemical equations, graphs, diagrams, or mathematical equations that will be required.

o Applications (why do we care?) o Explain the relationship between transmittance, absorbance, and

concentration o Explain why absorbance values greater than 1.5 are undesirable and what to

do if you encounter high absorbance values. o How to find the concentration of Cu2+ ions using Beer’s law and a calibration

curve, how to make the calibration curve o How qualitative and quantitative spectroscopy data may be used o How the wavelength to measure absorbance at is selected

2. Safety Data Sheets: Recreate the table below. Look up the SDS for all chemicals used during the lab and complete the table using the Hazards Identification section.

Chemical Eye Irritant Skin Irritant Respiratory Irritant Other Info

Copper(II) nitrate Nitric acid (15.8 M)

Summary of Safety Precautions:

3. Materials and Methods: Write a detailed step-by-step procedure for the experiment. Include all the materials, glassware and equipment that will be needed, and safety precautions that must be followed. Include:

o General use of the spectrophotometer (cleaning cuvettes, blanking, etc.) o How to create the calibration curve, including preparing all solutions from

0.400 M Cu(NO3)2 stock solution o How to measure the absorbance of the unknown concentration solution

PRE-LAB PRE-LAB PRE-LAB PRE-LAB PRE-LAB PRE-LAB PRE-LAB

4. Procedural Notes: Record details of the procedure, exact sizes of glassware, concentrations or amounts of chemicals, instructions for using LoggerPro, and any other notes provided by your instructor or group members. Finalize your procedure.

5. Data Collection: Create a data table to neatly record measurements, written with the correct precision based on the equipment, detailed observations, and sources of error.

6. Calculations: Show one neat handwritten worked example for each major required calculation. Include the unrounded answer and the answers with correct sig figs.

o Using Beer’s Law/ your linear fit to solve for the concentration of the unknown solution.

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Typed Lab Report 7. Typed Analysis: Using your collected data and a spreadsheet program such as Excel or

Google Sheets to create a calibration curve. Use the linear regression function to generate a linear fit to the data and display Copy and paste into the post lab such that it is readable when the page is printed.

8. Typed Conclusion: Include the following in narrative form: o Discussion of what you did, including the purpose/ goal of the experiment. Describe

any changes to the procedure from what you wrote in your lab notebook. o Describe major findings and make claims with supporting evidence from the lab.

Report calculated numbers, averages, and standard deviations as required. o Comment on error, both inherit error (due to precision of equipment) and

experimental/ human error. Remember that “the lab would have been better if we did it correctly” or “human errors were made” are not discussion of error. Errors should be clearly described along with how they impacted the calculations in the data analysis.

o Describe refinements or future experiments. What can we study next and why? 9. Typed Cover Sheet: Create a cover sheet using the exact table below. Place cover sheet

on top of items 4 to 8 from above.

Report Title: Percent Copper in Brass My Name:

Partners: Class Period:

Due Date:

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Intermolecular and Intramolecular Forces

Intermolecular Intramolecular Dispersion Forces (Induced Dipole)

Dipole-dipole Forces

H-Bonding Covalent Bond Ionic Bond Metallic Bond

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Reference Materials

Polyatomic Ions H3O+ hydronium CrO4

2– chromate

Hg22+ dimercury (I) Cr2O7

2– dichromate

NH4+ ammonium MnO4

– permanganate

C2H3O2–

CH3COO– acetate

NO2– nitrite

NO3– nitrate

C2O42– oxalate O2

2– peroxide

CO32– carbonate OH– hydroxide

HCO3– hydrogen (bi)carbonate CN– cyanide

PO43– Phosphate SCN– thiocyanate

ClO– hypochlorite SO32– sulfite

ClO2– chlorite SO4

2– sulfate

ClO3– chlorate HSO4

– hydrogen sulfate ClO4

– perchlorate S2O32– thiosulfate

Solubility Guidelines for Aqueous Solutions Ions that form Soluble Compounds

Exceptions Ions that form Insoluble Compounds

Exceptions

Group 1 ions (Li+, Na+, etc.)

Carbonate (CO32–)

When combined with Group 1 ions or ammonium

Ammonium (NH4+) Chromate (CrO4

2–) When combined with Group 1 ions, Ca2+, Mg2+ or ammonium

Nitrate (NO3–) Phosphate (PO4

3–) When combined with Group 1 ions or ammonium

Acetate (C2H3O2– or

CH3COOH) Sulfide (S2–)

When combined with Group 1 ions or ammonium

Hydrogen carbonate (HCO3

–) Hydroxide (OH–)

When combined with Group 1 ions, Ca2+, Ba2+, Sr2+ or ammonium

Chlorate (ClO3–)

Perchlorate (ClO4–)

Halides (Cl–, Br–, I–) When combined with Ag+, Pb2+, and Hg2

2+

Sulfates (SO42–)

When combined with Ag+, Ca2+, Sr2+, Ba2+, and Pb2+

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Strong Acids and Bases Acids Bases

o HCl o HBr o HI o HNO3 o HClO4 o H2SO4 o HClO3

o Group 1 hydroxides o NaOH o KOH

Pressure Conversion Factors 1 atm = 760 mmHg = 101.3 kPa

Temperature Conversions 𝐾 = ℃ + 273.15 ℃ =

5

9(℉ − 32)

Van der Waals Constants for Common Gases

Gas a (𝑎𝑡𝑚∙𝐿2

𝑚𝑜𝑙2 ) b (𝐿

𝑚𝑜𝑙)

He 0.0341 0.0237

Ne 0.211 0.0171 Ar 1.35 0.0322

Kr 2.32 0.0398 Xe 4.19 0.0511

H2 0.244 0.0266

N2 1.39 0.0391 O2 1.36 0.0318

Cl2 6.49 0.0562 CO2 3.59 0.0427

CH4 2.25 0.0428

NH3 4.17 0.0371 H2O 5.46 0.0305

Partial Pressure of Water Vapor for Select Temperatures Temperature

(°C) Pressure (mmHg)

Temperature (°C)

Pressure (mmHg)

15 12.8 26 25.2

16 13.6 27 26.7

17 14.5 28 28.3 18 15.5 29 30.0

19 16.5 30 31.8 20 17.5 31 33.7

21 18.7 32 35.7 22 19.8 33 37.7

23 21.1 34 39.9

24 22.4 35 42.2 25 23.8 36 44.6

Common Gas Law Constant Values

𝑅 = 8.31 𝐿 ∙ 𝑘𝑃𝑎


𝑅 = 0.0821 𝐿 ∙ 𝑎𝑡𝑚


62.364 𝐿 ∙ 𝑚𝑚𝐻𝑔


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