CHEM 101L UNC Final

If a flammable substance in a beaker should catch fire while you are working but the flames are relatively contained, what should you do to extinguish the flame? What should you not do?

DO: Use the chemical fire extinguisher from the lab. Pull the pin, aim the nozzle to the beaker, and pull the lever.

DO (Best, simplest): Turn off the heat source immediately and use a watch glass to cover the beaker.

DON'T: Use water from the nearest sink to douse the flame.

How many times should you read the label on a reagent container and compare it with the lab manual?

At least twice each time you remove material from the container.

While weighing out a reagent for use in an experiment, a student finds he has leftover reagent. What should the student do?

Place the excess reagent in the appropriate solid or liquid waste container to prevent contamination of the stock reagent.

List the correct procedure for lighting and working with a Bunsen burner.

1. Make sure all loose clothing and hair have been properly restrained and will not contact the flame accidentally.

2. Prepare the workspace to be sure no flammable materials are in the vicinity of the open flame

3. Inspect the tubing to the gas valve to be sure it does not have any cracks or holes that would allow a gas to leak.

4. Inspect the burner to be sure that the air intake and fuel valves work properly and can be adjusted.

5. With the sparker in hand, turn on the gas and light the flame as soon as possible to minimize unburned gas in the air.

6. Adjust the fuel valve to set the height of the flame and adjust the air intake valve to adjust the temperature (color) of the flame according to the instructions in the lab manual.

7. Monitor the flame and all nearby materials closely until the experiment is complete. Never leave an open flame unattended.

8. Shut off the burner and gas valve as soon as experimental work is completed, or anytime you must leave the flame unattended.

List the basic order of steps for beginning any experimental procedures.

1. Take any personal safety precautions that are warranted.

2. Prepare your workspace.

3. Inspect your equipment to make sure it is functioning properly.

When should you inspect the glassware in your drawer for chips, cracks, or chemical residues?

Glassware should be inspected during an experiment (before using any glassware), before beginning an experiment, and during check-in (where any glassware can be replaced for free) in order to prevent injuries, accidents, or unwanted complications.

Suppose a beaker of solid reagent drops onto the bench and cracks. What is the best way to dispose of this?

To the extent possible, solid reagent should be added to the solid waste container and broken glass should be added to the glass waste. Hazardous materials should never be added to glass waste, and broken glass should only be allowed in solid waste with the TA's permission.

Why does the lab safety policy address loose clothing, dangling jewelry, and contact lenses?

Loose clothing/dangling jewelry: can impede movement, knock materials over, can present a fire hazard.

Contact lenses: can hold hazardous chemicals close to the eyes.

All of these items increase the risk of injury.

Your lab partner accidentally gets a mist of your solution in her eye. You immediately help her to the eye wash to rinse it. After about a minute, she feels better and is ready to get back to work. What should you do?

Keep her in the eye wash for at least 15 minutes, and then make sure she goes to student health, just to be sure.

When reading the label on a reagent container, what are the three most important pieces of information?

Name, concentration, and hazard warning(s)

What are important precautions before using a hot plate?

Remember that the ceramic top of the hot plate, and any glassware heated by the hot plate will look the same when hot as cold, and will not cool down until well after the hot plate has been shut off.

Never leave the hot plate turned on an unattended.

Set up your work space with the hot plate in a secure location away from the edge of the bench so that you won't accidentally bump it and spill hot or flammable liquids.

Set up your work space so that wires and cables cannot accidentally make contact with the ceramic surface and melt.

Set up the work space so that flammable materials - notebooks, paper towels, other reagents, etc. - are far away from the hot plate to prevent them coming in contact with the hot surface.

If you need to work with a flammable or volatile solvent, which piece of lab equipment should you be sure to use?

Fume hood with good ventilation.

Your lab partner accidentally spills some acid on his wrist and watchband. What should you do?

Let the TA inspect his wrist to see if it is okay.

Remove the watch and wristband immediately, and rinse his wrist for at least 15 minutes to be sure all hazardous material has been washed away.

For which of the following situations should you be sure to notify your TA and fill out an incident report form?

You begin to feel faint and dizzy in lab because you had to skip lunch.

You accidentally pick up a beaker from the hot plate, not realizing that it was already hot, and the edge of the beaker leaves a small red mark on your thumb.

During lab check-out at the end of the semester, your lab partner accidentally drops a beaker, and a small chip of glass causes a scratch on your ankle because you weren't wearing socks.

During check-in you discover a broken funnel in your drawer, and the broken end of the stem causes a small nick through your glove.

You arrive in lab, but realize that your cold medicine is making you feel groggy.

If your ailment does not originate in lab, you must still let your TA know and he will propose an excused absence. If you cannot drive, you should not come to lab.

No matter how small the injury, you must fill out an incident report.

List the steps you would take if you find broken glassware in your lab drawer, or if a piece of glassware breaks during the lab.

1. Notify your TA.

2. Use a pair of leather gloves found in the lab to prevent nicks or cuts.

3. Clean up large pieces and dispose of the designated glass only waste container.

4. Use the dustpan and brush to sweep up small shards and dispose of them in the designated glass waste.

What is your BEST resource for understanding the nature of the chemical hazards of materials you work with in lab?

Safety Data Sheets

Consider the following scenario: A student has prepared a series of analytical solutions to be run on an instrument in another room. The student has worn gloves throughout the preparation of the solutions, and has wiped up all drips and spills that occurred during that process. The student removes her gloves before leaving the lab room and carries her solutions in scintillation vials in a large clean beaker to the room where the instrument is located. Once in the new room, the student obtains a new pair of gloves before dispensing the solutions for analysis.

What changes could be made to improve her procedures?

The student could have chosen a plastic bin, or other less breakable container to carry her solutions to the instrument so that she wasn't carrying glassware with her bare hands.

The student could have worn a single glove on the hand she used for transporting the solutions to the instrument room, and used her non-gloved hand for opening doors and pushing elevator buttons.

The student could have worked with her lab partner so that one of them could carry the materials with gloved hands, while the other opened doors and pressed elevator buttons with non-gloved hands

When should you be sure to use a sealed container and/or secondary containment in the laboratory?

When transporting materials from one lab room to another.

What is the most important consideration to preserve safety when it is necessary to feed glass tubing, thermometers, or other apparatus through a rubber stopper?

Using the correct hand position so that if something goes wrong you can avoid coming into contact with broken or sharp ends.

Which of the following is a behavior that can pose a safety risk in the laboratory environment?

Sitting on lab benches, lab floors, or hallway floors, while waiting for lab to begin or for another student to finish.

Leaving your lab drawer open while you set up your apparatus and obtain your reagents.

Staying focused on your own experiment and not being distracted by what nearby groups are doing.

Looking away from your work to answer a question from your TA or lab partner.

Moving rapidly around the lab to be sure to finish the experiment in time.

When the procedure calls for making a more dilute solution of an acid, or mixing an acid with other solutions, what is the correct order of steps?

Always Add Acid - Either add all of the water or non-acid component first, or add a significant portion, before adding the acid to the mixture. This helps to minimize the heat generated, which could otherwise create dangerous fumes or reactions.

When should you not wear gloves in the laboratory setting?

When entering data into your ELN on your laptop.

When writing notes on scratch paper using a pen and paper.

Define accuracy.

How closely obtained/measured values agree with the "true" value.

Define precision.

How closely obtained/measured values agree with each other.

Define systematic error and give examples.

An error that occurs due to a flaw in the experimental design or in the equipment. These types of errors can be detected and corrected.

Water in a heated beaker evaporating.

Air pressure not being 1 atm.

Incorrectly tared balance

A person consistently takes an incorrect measurement.

An instrument is worn out and measurements are difficult to read.

Not reading the meniscus of a liquid at eye level.

Affects the accuracy of results*

Define random error and give examples.

An error that arises from uncontrolled, and often times uncontrollable, variables in an experiment which has equal chance of being positive and negative. This type of error cannot be detected or corrected for.

Reading volume on a flask from a different angle each time.

Mass measurements can be affected as water enters or leaves the specimen being measured.

Estimations between the lines may be incorrect.

Affects the precision of results*

Why does ice float in water?

Ice is less dense than water.

How does temperature affect the density of a solution?

Temperature and density are inversely proportional. As temperature increases, density decreases.

Does calibration using standards help to improve the precision or accuracy of the experiment?

Accuracy.

Calibration defines the accuracy and quality of measurements recorded using a piece of equipment.

If a reaction is exothermic, what happens to the heat of each species?

Surroundings absorb heat and system releases heat.

For endothermic reactions, will the final temperature be higher or lower?

For exothermic reactions, will the final temperature be higher or lower?

ENDOTHERMIC: higher

EXOTHERMIC: lower

Should you add acid to water, or water to acid?

Acid to water.

If you add water to acid, you form an extremely concentrated solution of acid initially and the solution may boil very violently, splashing concentrated acid. If you add acid to water, the solution that forms is very dilute and the small amount of heat released is not enough to vaporize and spatter it.

What is the larger quantity: solvent or solute?

solvent

What is a precipitate?

Solid that forms in a solution in a chemical reaction.

Why do we use two styrofoam cups in a calorimeter set up?

The cups are great insulators and release little heat to the surroundings.

Why do you need a stir bar in the calorimeter?

The stir bar distributes the heat throughout the calorimeter.

Does making a bond release or require energy?

Release; bond is a more stable form.

Why do bond-breaking reactions require energy?

The atoms are more stable when bonded because they are attracted to each other.

Would you expect a bond forming reaction to be endo- or exothermic?

exothermic

Explain the difference between a stoichiometric reaction and a catalytic reaction. Give one example for each that was not a part of the experiment, and be specific with a type (redox, combustion, or acid-base reaction).

Catalytic reaction: lower activation energy than the corresponding uncatalyzed reaction, resulting in a higher reaction rate at the same temperature for the same reactant concentrations.

Example: Potassium permanganate as a catalyst for the decomposition of hydrogen peroxide. Redox reaction.

Stoichiometric reaction: quantities of the reactants and products show that all of the reactants are consumed and none remain after the end of a chemical reaction.

Example: Ethane combusts in the presence of oxygen to form water and carbon dioxide.

What is a decomposition reaction? Name on that you have learned in 101L.

One reactant yields two or more products.

Decomposition of NaOH (s)

What happens to a species when it is reduced?

It gains electrons.

Why is enthalpy change considered a state function? How does this relate to Hess's Law?

Enthalpy change is considered a state function because it is only dependent on two thermodynamic state properties (like temperature and pressure) of the state of the substance at the moment of change in enthalpy. The path required to reach these states has no effect on enthalpy change.

Purpose of Experiment 1: Mass, Volume and Density

The purpose of this experiment is to determine the density of water at room temperature to understand how temperature affects density. This lab is important for understanding how solutes can affect density, and for the researcher to have a better understanding of precision and accuracy.

Procedures for Experiment 1: Mass, Volume, and Density

Fill buret with assigned water sample and clamp it to a beaker on a hot plate.

Heat the water after noting initial temperature of the water.

Turn off hotplate upon reaching 80 degrees C and remove beaker.

Record temperature values upon seeing a change in volume.

Goal: Compare densities of different types of water (purified, distilled, contaminated)

Calculations for Experiment 1: Mass, Volume, and Density

(Mass of buret with sample) - (Mass of buret) = (Mass of water)

(Largest volume increment) - (Displaced Volume) - (Volume reading with thermometer) = (Corrected Volume)

(Mass of water) / (Corrected volume) = Density

[ (True Value) - (Mean) ] / (True Value) = (% error)

Standard deviation equation

square root of [ (the sum of x - mean)^squared ] / (n-1)

CV equation

100 * (standard deviation) / mean

Conclusion of Experiment 1: Mass, Volume, and Density

In this experiment, it was determined that the density of water decreases as temperature increases. It was also determined that water samples containing more solutes, like contaminated water and purified water, have higher densities on average.

Redox reaction

Involves a transfer of electrons between two species

Acid Base Reaction

Also called a neutralization reaction. When an acid and a base combine to form water and a salt.

Combustion reaction

A compound and an oxidant are combined to form heat and another product.

Combustion reactions between hydrocarbons and oxygen form water and carbon dioxide.

Purpose of Experiment 2: Aqueous reactions

The purpose of this experiment is to use calorimetry to understand the temperature changes that occur in a chemical reaction in aqueous solution. This lab is important because it allows the researchers to link observable changes and molecular changes in a chemical reaction.

Measuring the change in temperature of several chemical reactions in aqueous solution.

Possible chemical reactions in aqueous media

1. precipitation (forms a precipitate)

2. acid-base (water + salt, observed through litmus test)

3. complexation (usually a color change)

4. gas-forming

Procedures from Experiment 2: Aqueous Reactions

Place two Styrofoam cups inside of each other with a stir bar and distilled water inside, a lid on top, and a temperature probe inserted around the edge.

Place the apparatus on a stir plate, then lift the lid, add ammonium chloride, replace the lid, and turn on the stir plate.

Record the temperature at specified intervals and make observations by lifting the lid occasionally.

Dispose of solution according to procedures (liquid or solid waste) and repeat this process with sodium carbonate and distilled water, sodium bicarbonate with citric acid, and hydrochloric acid with magnesium.

Calculations for Experiment 2: Aqueous Reactions

(Grams of solid reactant) * (inverse of molar mass of solid reactant) = (Moles of solid reactant)

(Final Temperature) - (Initial Temperature) = (Change in temperature)

(Change in temperature) / (moles of reactant) = (Change in temperature / moles of reactant)

Conclusion for Experiment 2: Aqueous Reactions

Decrease in temperature of the surrounding water meant the reaction was endothermic, and vice versa. Calorimetry can be used to determine if reactions are endothermic or exothermic.

From this experiment, I learned that in each reaction, even if the same processes are occurring (dissolving, disassociating, gas-forming), the concept of exothermic and endothermic reactions are separate concepts.

Purpose of Experiment 4: Conductometric Titrations

The purpose of this experiment is to monitor the dissolution of barium hydroxide (a conductive solution) and then to perform a titration with sulfuric acid. After the titration, the researchers are meant to observe the change in conductivity of the solution as the ion concentration decreases. Initially, the dissolution will contain barium and hydroxide ions, and then, upon adding the sulfuric acid, molecular water and barium sulfate will form, decreasing the concentration of the barium and hydroxide ions. The conductivity will be measured using a conductivity probe, which detects ions in solution. Because ions in solution generate electricity, I hypothesize that a decrease in ion concentration would decrease conductivity of the solution.

Procedures for Experiment 4: Conductometric Titrations

Combine barium hydroxide and water and record conductivity with probe.

Calculate predicted volume of titrant and add titrant to approach predicted volume.

Record conductivity while adding more titrant to determine a minimum conductivity.

Repeat process for a total of three trials.

What experimental uncertainty is reduced by rinsing the buret with the sulfuric acid solution before filling it?

Rinsing the buret with the sulfuric acid solution before filling it ensures that there is no excess moisture in the buret that will affect the concentration of the sulfuric acid.

What is the end point of titration?

The volume of titrant necessary to reach the lowest conductivity in an acid base titration.

Calculations for Experiment 4: Conductometric Titrations

(Volume of barium hydroxide) (Molarity of Barium Hydroxide) (Molar ratio) * (Inverse of molarity of sulfuric acid) = (Volume of sulfuric acid)

(b2 - b1) / (m1 - m2) = (Volume of titrant extrapolated at minimum conductivity)

(Volume of sulfuric acid at minimum conductivity) * (Molarity of sulfuric acid) = (Moles of sulfuric acid)

Suppose you and your partner accidentally added sulfuric acid to the distilled water in the beaker instead of barium hydroxide; how would the shape of the conductivity plot change and why?

Since sulfuric acid is also aqueous, the acid would dissociate into its ionic components in solution. The shape of the conductivity plot would remain the same because the conductivity would be high before any base has been added because the solution is separated into ions, and the conductivity would continually decrease as barium hydroxide was added as the sulfate ions in solution bond with barium ions. The conductivity would then rise again after reaching the minimum conductivity and the end point of titration because barium hydroxide would continue to dissociate in solution after all sulfate ions have been bonded to barium ions.

Conclusion of Experiment 4: Conductometric Titrations

Through the utilization of titration to perform a neutralization reaction, it was determined that conductivity of a solution and ion concentration of a solution are directly proportional.

Purpose of Experiment 5: Gas Forming Reactions and Rockets

The purpose of this lab is to launch a pipette-bulb rocket using the products of a combustion reaction. To obtain the fuel for the rocket, two gas-producing reactions will be conducted in which the products will be used to perform the combustion reaction necessary to propel the rocket. The order of addition of O2 and H2 will be compared in the combustion reaction to see which gas being added first will produce the farthest rocket launch.

Procedures for Experiment 5: Gas Forming Reactions and Rockets

Build the gas-generating apparatus.

Mix the hydrogen peroxide with yeast in the apparatus test tube and collect the oxygen gas produced in the water-filled pipette bulb on the apparatus.

Mix the zinc and hydrochloric acid in the apparatus test tubes and collect the hydrogen gas produced in the water-filled pipette bulb on the apparatus.

Record how much oxygen gas and hydrogen gas were produced and the ratios used in the pipette bulb that will be launched.

Take the pipette bulb to the TA to ignite the gases and record all observations about the launch.

What experimental components hindered the launch of the rocket in Experiment 5?

Capped the bulb with a gloved finger only, excess water on the bulb prevented full combustion

Why were you asked to full the pipet bulb up with water before adding the gases?

The displacement of water allowed us to determine how much gas was being collected.

Reactions in Experiment 5: Rockets

Redox: Zinc and HCl

Decomposition: Hydrogen peroxide

Conclusion for Experiment 5: Rockets

Based on the experimental data, the order of gas addition did not have a strong effect on the distance traveled by the pipette bulb.

Purpose for experiment 6: Thermochemistry

The purpose of this lab is to determine the enthalpy change of the dissolution of sodium hydroxide, the neutralization of hydrochloric acid by sodium hydroxide, and the dissolution and neutralization of sodium hydroxide by hydrochloric acid. A calorimeter will be used to monitor the temperature changes of each of the reactions. In order to accurately and precisely measure the temperature of the reactions, a temperature probe with be used that connects to the computer and consistently records the temperature and the calorimeter provides precise results.

Procedures for Experiment 6: Thermochemistry

Set up the calorimeter.

Perform Reaction 1 in the calorimeter and record temperature values.

Perform Reaction 2 in the calorimeter and record temperature values.

Perform Reaction 3 in the calorimeter and record temperature values.

Reactions for Experiment 6: Thermochemistry

Reaction 1: Dissolution of solid sodium hydroxide

Reaction 2: Neutralization of liquid sodium hydroxide and hydrochloric acid

Reaction 3: Dissolution and neutralization of solid sodium hydroxide and hydrochloric acid

Hess's Law

Delta H of reaction = (Sum of delta h for products) - (sum of delta h for reactants)

Equations for Experiment 6: Thermochemistry

qcal = | qcold + qhot |

q = mcDT

ccal = (Dqcal) / (DTcold)

Tmax = y intercept for equation when temperature begins to decrease

qsoln = (msoln) (DT) (cH2O)

qrxn = - | qsoln + qcal |

DH = (qrxn) / (moles of reagent)

Calculations for Experiment 6: Thermochemistry

(mliq) + (msolid) = (msoln)

qcal = (ccal) * DT

Conclusion from Experiment 6: Thermochemistry

The enthalpy changes of the three reactions with sodium hydroxide were calculated by using change in temperature through calorimetry. This calculation of enthalpy change through change in temperature in exothermic reactions demonstrates Hess's Law by showing, with minimal variance from theoretical values, that the sum of the change in enthalpy of two reactions is equal to the change in enthalpy for that combined reaction.

Purpose of Experiment 7: Beer's Law

The purpose of this experiment is to design an experiment to determine the concentration of Allura Red AC dye in Gatorade and Powerade by utilizing Beer's Law. Data will be used with Beer's Law in order to compare the amount of dye in Gatorade and Powerade by measuring the absorbance values of dilutions of red Gatorade and Powerade and plugging those absorbance values into the calibration curve regression line. This experiment also allows the researchers to compare diluted concentrations and neat concentrations in order to determine the amount of Allura Red AC in both red Gatorade and Powerade.

Procedures for Experiment 7: Beer's Law

Prepare the stock standard solution.

Calibrate the spectrometer with different concentrations of Allura Red AC dye.

Use the spectrometer to determine the absorbance of the varying concentrations of the dye.

Prepare standard concentrations of Gatorade and Powerade.

Use the spectrometer to determine the absorbance of the Gatorade and Powerade.

Equations for Experiment 7: Beer's Law

Absorbance = (Molar absorptivity e) (the path of light l) (Concentration of analyte C)

For the most accurate determination, the absorbance of the samples should fall in the middle of your calibrated range (~0.5-0.7 absorbance units). What will you do if the absorbance for your sample is greater than 1.2? Explain the process in detail, including the instruments used. Remember that you must preserve the precision to 4 significant figures.

If the absorbance for your sample is greater than 1.2, we must consult our TA to determine whether calibration standards should be prepared again or if any standard concentrations should be adjusted. Because of the relationships seen in the equation for absorbance (A=ειC), a higher absorbance would result from a higher concentration (considering ε and ι are constant). To lower the absorbance, a decrease in concentration would be required and the solution providing the higher absorbance value should be diluted. For this experiment, 501 nm wavelength gave the highest absorbance value. If it was necessary to decrease this maximum wavelength, lower concentrations should be used.

Conclusion for Experiment 7: Beer's Law

In this experiment, we put Beer's Law into practice while understanding measurements regarding concentrations and how they relate to wavelength and absorbance ranges. While comparing different concentrations of solutions, the higher concentrations had a higher absorbance value, thus proving that absorbance and concentration are directly proportional.

Purpose of Experiment 8: Spectroscopy

The purpose of this experiment is utilize spectroscopy through observing the continuous and line spectra of different mixtures to determine the wavelengths, frequencies, and energies of the different mixtures and to identify unknown mixtures. Comparative observations and measurements will allow the researchers to make connections between observed line and emission spectra and energy levels.

Procedures for Experiment 8: Spectroscopy

Calibrate the spectroscope.

Use the spectroscope to record the spectral characteristics of three sources of light.

Use the spectroscope and spectrometer to determine the line and emission spectra of hydrogen gas.

Use the spectroscope to observe the spectral characteristics of different metals.

Identify an unknown metal from the previously observed spectral characteristics of different metals.

Calculations for Experiment 8: Spectroscopy

Wavelength = [ (Lowest position on spectroscope) - (Y-intercept from calibration equation) ] / (slope of calibration equation)

(Speed of light) / (wavelength) = frequency of photon

Conclusion for Experiment 8: Spectroscopy

The position of an emission spectra on a spectroscope can provide sufficient data to determine the wavelength of a light source, which can be used to determine frequency, energy, and energy levels of electrons; the spectrometer provides a more accurate wavelength through the optical fiber, which provides sufficient data to determine the color of emission lines of a light source. By using both instruments, the researchers could more accurately predict the identity of the unknown mixture.

The volume marking at the sealed bottom of a buret segment is 48.60 mL, and has markings for every 0.1 mL. When water is added, the bottom of the meniscus sits about halfway between the line for 28.3 mL and the next mark up. What volume of water is contained in the buret segment?

20.35 mL

When calculating the displacement volume of the thermometer, the volume reading on the buret is 29.37 mL without the thermometer, and 28.72 mL with the thermometer in place. How many significant figures should be recorded for the displacement volume of the thermometer?

2

The molecular equation of a reaction is NaHCO3 (aq) + HCl (aq) →NaCl (aq) + H2O (l) + CO2 (g). Both the initial solutions and the final solution were transparent and colorless, but bubbles were seen during the reaction. Based on this equation and the observations of the reaction, the spectator ions in the ionic equation would be:

Na and Cl

The temperature at the start of a reaction was 26.5°C and at the end of the reaction was 18.5°C. Based on these observations the reaction of sodium bicarbonate with citric acid would be:

endothermic, because energy is transferred to the reaction from the surroundings.

Based on your observations in Experiment 5, an exothermic reaction may be identified via observation in the lab when it:

Causes a loud popping noise and/or produces bubbles

T/F: The following three actions could produce a significant source of error in the results for Experiment 5:

Using an excess of yeast in the reaction with H2O2.

Using an excess of HCl in the reaction with zinc.

Collection both gases in the same pipet bulb.

False

To carry out a calorimetry experiment quantitatively, what is necessary?

Calculations must be made in moles rather than grams.

The thermometer must be calibrated to give accurate temperature readings in Celsius.

The calorimeter must be calibrated to account for heat loss of the system.

All solutions must be equilibrated to room temperature before beginning the experiment.

In calorimetry, the enthalpy change of the reaction, qrxn, is defined to be:

-(qsoln + qcal) and the heat gained or lost by the reaction in units of joules.

In a calorimetry experiment, 50.0 mL each of two dilute aqueous solutions are combined and produce an extrapolated ΔT = -22.4˚C. If the density of each solution is assumed to be equal to that of water (at 20˚C, d = 0.9982 g/mL), and the specific heat of the solution is also equal to that of water (4.186 J/g˚C), the value of qrxn would be:

9.36 kJ

Which of the following is a true statement about the bright yellow flame seen during a flame test of a 1 M solution of NaCl in a Meker burner flame?

The yellow color is characteristic of the changes in energy of an electron in the Na+ metal ion.

Given the Rydberg equation: ΔE = RH (1/(nf2)- 1/(ni2))

In the Balmer series for hydrogen, what is the initial energy level (ni) for a photon with a wavelength of 410.2 nm? REMEMBER: h = 6.626 x 10-34 J⋅s; c = 2.998 x 108 m/s; RH = 2.18 x 10-18 J

6

What is the purpose of a calibration curve?

general method for determining the concentration of a substance in an unknown sample by comparing the unknown to a set of standard samples of known concentration.

What causes absorption

Intensity of light passing through sample compared to the blank in the spectrometer

Define the variables in Beer's Law

A: (no units), absorption, the intensity of light passing through the sample relative to the intensity of light passing through the blank (l0 = -log(l/l0)

ε: (1/M*cm), molar extinction coefficient/ molar absorptivity, describes how strongly the analyte absorbs light at a particular wavelength

ι: (cm), the path length of light through the solution contained in the cuvet

C: (M), the concentration of the analyte in solution