Experimental - Review the Knowledge You Need to Score High - 5 Steps to a 5 AP Chemistry (2015)

5 Steps to a 5 AP Chemistry (2015)

STEP 4. Review the Knowledge You Need to Score High

CHAPTER 19. Experimental

SPECIAL NOTE

At the time this book was revised, the new AP Chemistry experiments had not been released. However, the following experiments should be very similar to those included on the exam. In studying the experiments, focus on the design of each experiment and the general concepts each experiment covers.

IN THIS CHAPTER

Summary: The free-response portion of the AP exam will contain a question concerning an experiment, and there may also be a few multiple-choice questions on one or more of these experiments. This chapter reviews the basic experiments that the AP Exam Committee believes to be important. You should look over all of the experiments in this chapter and pay particular attention to any experiments you did not perform. In some cases you may find, after reading the description, that you did a similar experiment. Not every AP class does every experiment, but any of these experiments may appear on the AP exam.

The free-response questions on recent exams have been concerned with the equipment, measurements, and calculations required. In some cases, sources of error are considered. To answer the question completely, you will need an understanding of the chemical concepts involved.

To discuss an experiment, you must be familiar with the equipment needed. In the keywords section at the beginning of this chapter is a complete list of equipment for the experiments (see also Figure 19.1, on the next spread). Make sure you are familiar with each item. You may know an item by a different name, or you may need to talk to your teacher to get additional information concerning an item.

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Figure 19.1 Common laboratory equipment.

In some cases, the exam question will request a list of the equipment needed, while in other cases you will get a list from which to choose the items you need. Certain items appear in many experiments. These include the analytical balance, beakers, support stands, pipets, test tubes, and Erlenmeyer flasks. Burets, graduated cylinders, clamps, desiccators, drying ovens, pH meters, volumetric flasks, and thermometers are also commonly used. If you are not sure what equipment to choose, these serve as good guesses. Most of the remaining equipment appears in three or fewer experiments.

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You will need to know the basic measurements required for the experiment. For example, you may need to measure the initial and final temperatures. Do not make the mistake of saying you measure the change in temperature. You calculate the change in temperature from your measured initial and final temperatures. You do not need to give a lot of detail when listing the required measurements, but you need to be very specific in what you measure. Many students have lost exam points for not clearly distinguishing between measured and calculated values.

The basic calculations fall into two categories. Simple calculations, such as the change in temperature or the change in volume, are the easiest to forget. Simple calculations may also include mass-to-mole conversions. The other calculations normally involve entering values into one of the equations given at the beginning of the previous chapters of this book.

Beginning with the 2007 AP Chemistry exam, experimental questions began to be incorporated into the free-response questions. This means that you will need to have a good understanding of the experiments in order to discuss not only the experiment itself, but the underlying chemical concepts. Therefore, when studying each experiment, refer back to the appropriate material to review the concepts involved.

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Keywords and Equations

Pay particular attention to the specific keywords and equation in the chapters associated with the individual experiments.

A = abc (A = absorbance; a = molar absorbtivity; b = path length; c = concentration)

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Experiment 1: Finding the Formula of a Compound

Synopsis

The formula of a compound is determined by using the mass of the original substance, usually a metal, and the mass of a compound of that substance, usually an oxide. (See Chapter 7 Stoichiometry.)

Equipment

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Measurements

1. the mass of the crucible and cover

2. the mass of the original sample and the crucible and cover

3. the mass of the reacted sample and the crucible and cover

The last measurement must be done after the sample has cooled to room temperature.

Calculations

The mass of the sample is calculated by taking the difference between masses 1 and 2. The mass of the substance that combined with the original substance is calculated from the difference between masses 2 and 3. The moles of the substances must be calculated by dividing each mass by the molar mass of the substance. The empirical formula is calculated from the simplest ratio of the moles of the elements present.

Comments

This procedure will allow you to calculate the empirical formula of the substance. The experiment is often performed by reacting magnesium metal with atmospheric oxygen to form magnesium oxide.

Experiment 2: The Percentage of Water in a Hydrate

Synopsis

The amount of water present in a sample is determined by weighing a compound before and after heating. The difference in mass is due to the loss of water. (See Chapter 7 Stoichiometry.)

Equipment

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Measurements

1. the mass of the crucible and cover

2. the mass of the original sample and the crucible and cover

3. the mass of the heated (dried) sample and the crucible and cover

The last measurement must be done after the sample has cooled to room temperature.

Calculations

The mass of the hydrate is calculated from the difference between masses 1 and 2. The mass of the water lost is calculated from the difference between masses 2 and 3. The percentage of water is calculated by dividing the mass of the water lost by the mass of the hydrate and multiplying the result by 100%.

A variation of this experiment uses the mass of the anhydrous material (calculated from the difference between masses 1 and 3). The moles of the anhydrous material and water are then calculated from their respective masses and molar masses. The simplest ratio of the moles gives the empirical formula.

Comments

The experiment often uses hydrates of copper(II) sulfate, magnesium sulfate, calcium sulfate, or barium chloride.

Experiment 3: Molar Mass by Vapor Density

Synopsis

The molar mass (molecular mass) of a volatile substance is determined in this experiment. The mass of a sample of vapor is initially determined. This mass, along with the volume of the container, the pressure, and the temperature, is used with the ideal gas equation to calculate the molar mass. (See Chapter 8 Gases.)

Equipment

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Measurements

1. the mass of the flask

2. the mass of the flask plus condensed vapor

3. the temperature of the water bath used to heat the flask

4. the barometric pressure

5. the number of milliliters of water required to fill the flask

Calculations

There are a variety of ways to do the calculations. Most of these, however, involve the calculation of the number of moles (n) from the ideal gas equation: n = PV/RT. The mass of the vapor sample is calculated from the difference between measurements 1 and 2. The temperature (measurement 3) is converted to kelvin. The pressure (measurement 4) is converted to atmospheres. Measurement 5 is converted to liters. Inserting the various numbers into the ideal gas equation allows you to calculate the number of moles. The molar mass is calculated by dividing the mass of the sample by the moles.

Comments

Variations in this experiment usually combine the ideal gas equation with the mass of the sample.

Experiment 4: Molar Mass by Freezing-Point Depression

Synopsis

The molar mass (molecular mass) of a solute is determined by measuring its effect on the freezing point of a solvent. (Refer to Figure 19.2 for general experimental setup.) A cooling curve is constructed by plotting the temperature of a solution that is slowly cooling versus time. After the solution completely freezes, the difference between the solution’s freezing point and the pure solvent’s freezing point is calculated. The change in the freezing point is then related to the molality of the solution. (See Chapter 6 Reactions and Periodicity and Chapter 13 Solutions and Colligative Properties.)

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Figure 19.2 Freezing-point depression apparatus.

Equipment

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Measurements

1. the mass of the empty test tube

2. the mass of the test tube plus the solvent

3. the mass of the test tube plus solute plus solvent

The solid solvent or solution is heated above its melting point and then allowed to cool.

4. repeated measurements of the temperature

5. the times at which the preceding measurements were made

The preceding series of measurements is done one time for the solvent and one time for the solution, or the melting point of the solvent may be obtained (not measured) from a table.

Calculations

The mass of the solvent is calculated from the difference between measurements 1 and 2. The mass of the solute is calculated from the differences between masses 2 and 3. The mass of the solvent is converted to kilograms.

The temperature and time measurements (4 and 5) are plotted, and a smooth curve is drawn.

The temperature difference between the “level” regions of the solvent plot and the solution plot (or the difference between the solution plot and the tabulated freezing point of the solvent) is used to calculate the change in temperature (ΔT) between the freezing point of the solvent and the solution. The change in temperature divided by the freezing-point depression constant (from a table) will give the molality of the solution. The molality of the solution times the kilograms of solvent yields the moles of solute. Finally, the mass of the solute divided by the moles of solute gives its molar mass.

Comments

One variation in this experiment is to add the solute to the test tube before the solvent.

Experiment 5: Molar Volume of a Gas

Synopsis

The volume occupied by a mole of a gas is calculated in this experiment. A sample of a solid substance is heated, decomposing it into several products, including a gas. The mass of the gas is determined by the weight difference of the solid before and after heating and is then converted to moles. The volume of the gas, the pressure, and temperature are measured. (See Chapter 8 Gases.)

Equipment

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Measurements

1. the barometric pressure

2. the mass of the test tube plus solid sample (before the reaction)

3. the mass of the test tube plus sample after the reaction

4. the temperature of the water

5. volume of water displaced into the beaker, if the variation is used

The mass of the test tube after the reaction must be determined after the test tube has completely cooled to room temperature.

Calculations

The temperature must be converted to kelvin (T1), and the volume of water is normally expressed in liters.

The mass of gas generated is calculated by taking the difference between measurements 2 and 3. Using the molar mass of the gas, the mass of gas is converted to moles of gas (n).

For the variation, the vapor pressure of water at the recorded temperature is found in a table. The pressure of the gas (P1) is the difference between the value in the table and measurement 1.

The volume of water in the beaker is the volume of the gas (V1).

Calculate the volume (V2) of the gas at STP (T2 and P2) using the combined gas law. The molar volume of the gas is the volume at STP (V2) divided by the moles of gas (n).

Comments

The most common procedure is to produce oxygen gas by decomposing KClO3. A common variation is to measure the volume of gas produced by displacing water from a flask. The volume of water displaced is the volume of gas generated at that temperature and pressure. From a measurement of the atmospheric pressure and the temperature of the gas, the volume of gas at STP can be calculated.

Experiment 6: Standardization of a Solution

Synopsis

The concentration of a solution is determined by titration with a sample of known composition. (See Chapter 7 Stoichiometry.)

Equipment

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Refer to Figure 19.3 for general titration setup.

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Figure 19.3 General acid–base titration setup

Measurements

1. the mass of an empty flask

2. the mass of the flask plus the sample

3. the initial reading of the buret

4. intermediate readings from the buret

5. the final reading of the buret

6. the pH of the solution at various times during the reaction

Calculations

The mass of the sample is calculated from the differences between masses 1 and 2. The volume added is calculated by taking the difference between measurement 3 and either measurement 4 or 5.

A plot of pH versus the volume added is made. This graph or the difference between measurements 3 and 5 gives the volume of titrant.

The volume of titrant is converted to liters.

The mass of the sample is converted to moles by using the molar mass. The moles of titrant may be calculated from a consideration of the moles of sample and the balanced chemical equation. The moles of titrant divided by the liters of solution gives the molarity of the solution.

Comments

A solution could be prepared by dissolving a known amount of solute in a volumetric flask and diluting to volume.

The course of the titration could be followed with an acid–base indicator instead of a pH meter.

Experiment 7: Acid–Base Titration

Synopsis

The concentration of an acid or a base may be determined by titrating a solution of an unknown concentration with a solution of a known concentration. (See Chapter 6 Reactions and Periodicity and Chapter 7 Stoichiometry.)

Equipment

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Measurements

1. the volume of the solution of acid or base using a pipet

2. the initial reading of the buret

3. intermediate readings from the buret

4. the final reading of the buret

5. the pH of the solution at various times during the reaction

Calculations

The volume added is calculated by taking the difference between measurement 2 and either measurement 3 or 4.

A plot of pH versus the volume added is made. This graph or the difference between measurements 2 and 3 gives the volume of titrant.

The volume of titrant is converted to liters.

The pipeted volume is converted to moles by multiplying the liters of solution by its molarity. The moles of titrant are determined using the mole ratio in the balanced chemical equation for the acid–base reaction. The molarity of the solution is calculated by dividing the moles of titrant by the liters of titrant used.

Comments

It does not matter whether an acid is titrated by a base or vice versa.

Experiment 8: Oxidation–Reduction Titration

Synopsis

The concentration of either an oxidizing or a reducing agent may be determined by titrating a solution of an unknown concentration versus a solution of a known concentration or containing a known mass of solute. (See Chapter 7Stoichiometry.)

Equipment

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Measurements

1. the volume of a solution of an oxidant or reductant using a pipet

2. the initial reading of the buret

3. the final reading of the buret at the endpoint of the titration

Calculations

The volume of titrant added is calculated by the difference between measurements 2 and 3.

The volume of titrant is then converted to liters.

The pipeted volume is converted to moles by multiplying the liters of solution by its molarity. The moles of titrant are determined from the mole ratio in the balanced chemical equation for the reaction. The molarity of the solution is calculated by dividing the moles of titrant by the liters of titrant used.

Comments

Common oxidants are potassium permanganate and potassium dichromate.

Iron(II) and oxalates are commonly chosen as reductants.

Refer to Figure 19.3 for general titration set-up.

Experiment 9: Mass/Mole Relationships in a Chemical Reaction

Synopsis

The initial masses of various reactants may be determined and then converted to moles. A similar calculation may be done for the products. (See Chapter 7 Stoichiometry.)

Equipment

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Measurements

1. mass of samples of different reactants

2. mass of samples of different products

Calculations

Convert the masses of the reactants and products to moles using their molar masses. Using the mole ratios from the balanced chemical equation, it is possible to determine how much material should react or be produced. These calculated values can be compared to the observed values.

Comments

Nearly any reaction may be used.

Experiment 10: Finding the Equilibrium Constant

Synopsis

The value of an equilibrium constant is calculated by measuring (or calculating) the equilibrium concentrations of the reactants and products. A calibration curve is constructed by measuring the absorbance of a colored solution versus its concentration. Known quantities of the reactants are mixed, and the calibration curve is used to determine the concentration of the colored substance in the resultant solution. (See Chapter 15 Equilibrium.)

Equipment

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Spectrophotometers

Some of the AP recommended experiments require the use of a spectrophotometer. A spectrophotometer is an instrument that is used to measure the amount of light absorbed (or percentage transmitted) by a particular solute in a solution. In order to determine the absorbance (A) of a sample, the instrument is set to a particular wavelength; a solution, contained in a holder called a cuvette, is placed in a sample chamber and an absorbance reading is taken. This procedure may be repeated for other solutions or wavelengths. The cuvette is a standard size to ensure a given path length (b).

A plot of absorbance versus wavelength may be used to identify a component of a solution or to determine the wavelength of maximum absorbance (maximum molar absorptivity = a). A more common plot is one of absorbance versus concentration. For this type of plot the instrument is set at the wavelength of maximum molar absorptivity and the absorbances of solutions of various known concentrations (c) are measured. This plot should be a straight line. This linear relationship is called Beer’s law and has the form of A = abc. The concentration of an unknown solution may be determined by measuring its absorbance and using the plot to find its concentration.

Measurements

Calibration Curve

Quantities of one or more reactants are pipeted into a volumetric flask.

The solutions are diluted to a known volume in the volumetric flask. Measurements of the absorbance are made with a spectrophotometer.

Equilibrium Concentrations

Different quantities of various reactants are pipeted into a volumetric flask and diluted to a known volume.

Measurements of the absorbance of these solutions are made with a spectrophotometer.

Calculations

Calibration Curve

The concentration of the absorbing species is calculated using the initial pipeted volumes and the final volumetric flask volume. These concentrations are plotted versus the absorbance of the solution.

Equilibrium Concentrations

The final concentration of each reactant is calculated from the final volume and the volume and concentration of the solution pipeted into the volumetric flask. The calibration curve is used to find the equilibrium concentration. Using the balanced chemical equation, the equilibrium concentrations of the other substances may be calculated.

The equilibrium concentrations are inserted into the reaction quotient expression, and the equilibrium constant is calculated.

Comments

Any equilibrium may be used, as long as one substance is colored.

Experiment 11: pH Measurements and Indicators for Acid–Base Titrations

Synopsis

The acidity of various substances is determined with a pH meter or acid–base indicators. This may also be done by mixing or diluting solutions. (See Chapter 15 Equilibrium.)

Equipment

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Measurements

1. Weigh each acid or base into a volumetric flask, and dilute to volume.

2. Pipet a sample of an acid or base of known concentration into a volumetric flask, and dilute to volume.

3. Pipet different solutions into a flask.

4. Measure the pH of the solutions using a pH meter.

Calculations

All weighed samples are converted to moles by using the molar mass, and the moles are divided by the volume of the volumetric flask in liters to yield molarity.

The concentrations of the diluted solutions (measurement 2) are calculated by using the dilution equation.

The concentrations of the other solutions (measurement 3) are calculated from the balanced chemical equation and the dilution equation.

The pH may be estimated by adding an acid–base indicator to any of the prepared solutions.

The hydrogen ion concentration, hydroxide ion concentration, or pOH may be calculated from the pH. One or more of these may be used to determine the concentration of all other species in the solution.

Comments

The original acids and bases may all be solids or solutions, or a mixture of both.

Experiment 12: The Rate and Order of a Reaction

Synopsis

The rate equation for a reaction is determined in this experiment. Known quantities of solutions are mixed, and the time required for a change is recorded. (See Chapter 14 Kinetics.)

Equipment

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Measurements

1. Weigh solid samples.

2. Use a pipet or a buret to measure the volume of any solutions.

3. Use the buret to measure the volume of any gas formed (multiple measurements at different times may be required).

4. Measure the temperature of the solutions.

5. Measure time intervals or record the time after mixing when an observable change occurs.

Note: When using a buret, an initial and a final measurement are always needed.

Calculations

Using the molar mass, calculate the moles of all weighed samples. The moles of substances are converted to molarities by dividing by the volume (in liters) of the solution. Molarities may also be determined from pipet or buret readings using the dilution equation. (If a buret is used, one of the volumes is calculated from the difference between the initial and final readings.) The dilution equation may be needed to calculate the concentration of each reactant immediately after all the solutions are mixed.

If a gas is being generated, plot the volume of gas formed versus time. The volume of the gas formed is the difference between the initial buret reading and the buret reading at a particular time. The slope of this graph is the rate.

The rate may also be determined by taking the amount of any reactant divided by the time required.

Tabulate the concentrations of the reactant solutions and the rates for various trials. The rate law may be determined by comparing values in this table. (See the Kinetics chapter.)

Comments

Many types of reaction may be used. The simple recording of the time required for a noticeable change is particularly applicable to “clock” reactions.

Experiment 13: Enthalpy Changes

Synopsis

In this experiment the heat change associated with a process is calculated. Various substances are added to a calorimeter (usually a polystyrene cup), and the initial and the final temperatures are measured. (See Chapter 7Stoichiometry.)

Equipment

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Refer to Figure 9.1 showing two types of calorimeters in Chapter 9.

Measurements

1. The masses of various substances are determined.

2. The volume of a solution or solvent is determined with a graduated cylinder.

3. Measure the initial temperature.

4. Measure the final temperature.

Calculations

The masses are converted to moles using the molar mass. The volumes of solutions may be converted to moles using the molarity.

The mass of a solvent or solution may be calculated from the volume and density.

A total mass may need to be calculated by adding the individual masses together.

The change in temperature (ΔT) is calculated from the difference between the final and initial temperatures.

The energy (joules or calories) is calculated by multiplying the mass times the change in temperature times the specific heat (from a table).

The enthalpy change is calculated by dividing the calculated energy by the moles, mass, or some other designated quantity.

Comments

Any reaction or phase change may be used.

Experiment 14: Qualitative Analysis of Cations and Anions

Synopsis

Solutions containing various ions are tested for the presence of certain specified ions. The formation of colors, gases, or precipitates indicates the presence or absence of certain ions. (See Chapter 6 Reactions and Periodicity.)

Equipment

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Measurements

This is primarily an observation experiment. No precise measurements are needed.

Calculations

No calculations are required.

Comments

A wide variety of cations or anions may be used. Many of the results are a simple application of the solubility rules.

Experiment 15: Synthesis and Analysis of a Coordination Compound

Synopsis

A coordination compound is usually synthesized from a transition metal ion that is in solution. The compound is filtered from the solution and tested in various ways to determine the composition of the substance. (See Chapter 6Reactions and Periodicity.)

Equipment

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Measurements

Synthesis

Quantities of the reactants are weighed and added to a beaker. A solvent may or may not be present initially.

The mass of the product is determined by weighing.

Analysis

Weigh a sample of the compound into a volumetric flask, and dilute to volume.

Use a pipet to measure samples to be diluted.

Measure absorbency of the diluted solutions with a spectrophotometer. This may or may not require the separate construction of a calibration curve.

Calculations

Calculate the moles of each reactant from the masses and molar masses. Then calculate the yield based on the limiting reagent. The mass of the product, determined at the end of the synthesis, divided by the mass calculated from the limiting reagent times 100%, gives the percent yield.

There are numerous possible analysis calculations.

Comments

Commonly synthesized coordination compounds include K3[Fe(C2O4)3] and [Co(NH3)6]Cl3.

Experiment 16: Gravimetric Analysis

Synopsis

The amount of a substance present in a sample is determined by taking a solution containing that substance and precipitating a compound containing that substance. The precipitate is then dried and weighed. (See Chapter 7Stoichiometry.)

Equipment

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Measurements

1. A sample is weighed and then dissolved.

2. Excess (unmeasured) reactant is added to the solution to form a precipitate.

3. The empty crucible and cover are weighed.

4. The crucible and cover containing the dried precipitate are weighed.

Calculations

The mass of the precipitate is found as the difference between measurements 3 and 4.

The mass of the dried precipitate is converted into moles by using the molar mass. Through use of a stoichiometric ratio, the moles of precipitate are converted to the moles of the substance of interest. The moles of this substance are converted to its mass using the molar mass.

The mass of the substance of interest divided by the mass of the sample and then multiplied by 100% gives the percent of a substance in the sample.

Comments

Common precipitates used include AgCl and BaSO4.

Experiment 17: Colorimetric Analysis

Synopsis

This experiment involves determining the amount of a substance in a solution. A calibration curve is constructed plotting the measured absorbance versus the concentration of a colored substance. The concentration of an unknown solution may be determined by reversing this process. (See Chapter 7 Stoichiometry.)

Equipment

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Measurements

Using either a pipet or a buret, quantities of standard solutions are measured. (If a buret is used, separate measurements of the initial and final volumes are needed.) Solvent may be added to dilute the samples if needed. These are the known solutions from which a calibration curve will be constructed.

The absorbance of each solution is determined with a spectrophotometer.

The absorbance of one or more unknown solutions is determined with a spectrophotometer.

Calculations

Concentrations of the known solutions are calculated using the dilution equation.

A plot of absorbance versus concentration for the known solutions is made.

The plot allows the absorbance of the unknown solution(s) to be converted to concentration.

Comments

Any colored substance may be used.

Experiment 18: Chromatographic Separation

Synopsis

Two or more substances are separated by the differences in their affinity to paper or some other material. (Refer to Figure 19.4.) This affinity is related to the intermolecular forces between the substances. (See Chapter 12 Solids, Liquids, and Intermolecular Forces.)

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Figure 19.4

Equipment

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Measurements

In some cases, the distance a spot travels on the chromatographic media is measured. This requires the use of a ruler.

Calculations

There are usually no calculations.

Comments

Instead of an ion exchange resin or silica gel, it is possible to use filter paper as a chromatographic medium.

Experiment 19: Properties of Buffer Solutions

Synopsis

Buffer and non-buffer solutions are prepared. The pHs of these solutions are determined before and after other substances—usually acids or bases—are added. (See Chapter 15 Equilibrium.)

Equipment

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Measurements

1. volumes of the pipeted solutions

2. pH of various solutions

Calculations

The concentrations of the solutions may be calculated by using the dilution equation. Concentrations may then be converted to moles by multiplying the concentration by the liters of solution. This procedure applies to buffer components or any reactant species.

The moles of substances may be determined from the initial moles and stoichiometry. Combined with the liters of solution, these may be used to determine the final concentrations.

Comments

The pH values of the solutions may be used in several ways, depending upon the goal of the experiment.

Experiment 20: An Electrochemical Series

Synopsis

The reactivity of several metals with solutions containing ions of other metals is observed. This experiment involves the development of an activity series. (See Chapter 6 Reactions and Periodicity.)

Equipment

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Measurements

This experiment is based on observations, not measurements.

Calculations

Observations, not calculations, are needed.

Comments

The more active metal will displace a less active metal from a solution. Hydrogen is usually included in the series by using acids.

Experiment 21: Electrochemical Cells and Electroplating

Synopsis

Electrochemical cells are constructed, and their cell potentials are determined with a voltmeter. Electroplating is accomplished by using an external power supply, usually a battery, to plate a metal onto an electrode. (See Chapter 16Electrochemistry.)

Equipment

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Measurements

The cell potentials are measured with a voltmeter. Refer to Figure 16.1 of the electrochemical cell in Chapter 16.

Calculations

Normally, no calculations are required.

Comments

Cells may vary in composition or concentrations.

Experiment 22: Synthesis and Properties of an Organic Compound

Synopsis

Any of a number of chemical reactions can be used to synthesize an organic compound. After synthesis, the compound is purified and tested. (See Chapter 18 Organic Chemistry.)

Equipment

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Measurements

Quantities of reactants are measured by using either mass or volume measurements. In some cases, the mass of the product is measured.

Calculations

When the mass of the product (actual yield) is measured, normally a percent yield is required. The mass of the limiting reagent is converted, through moles, to the theoretical yield of product. The percent yield is calculated by dividing the actual yield by the theoretical yield, then multiplying the resulting value by 100%.

Comments

Many different compounds could be synthesized.

Image Common Mistakes to Avoid

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1. You measure initial and final values, but calculate the change.

2. You use an analytical balance to weigh the mass (grams), but not the moles.

Image Review Questions

Below you will find a multipart free-response question like the ones in Section II of the exam. Use this question to review the content of this chapter and practice for the AP Chemistry exam. To make this an even more authentic practice for the actual exam, time yourself following the instructions provided.

Free-Response Question

You have 15 minutes to answer the following question. You may use a calculator and the tables in the back of the book.

Question 1

A sample of a solid, weak monoprotic acid, HA, is supplied, along with solid sodium hydroxide, NaOH, a phenolphthalein solution, and primary standard potassium hydrogen phthalate (KHP).

(a) Describe how a standardized sodium hydroxide solution may be prepared for the titration.

(b) Sketch a graph of pH versus volume of base added for the titration.

(c) Sketch the titration curve if the unknown acid was really a diprotic acid.

(d) Describe the steps to determine Ka for HA.

(e) What factor determines which indicator should be chosen for this titration?

Answer and Explanation for the Free-Response Question

(a) A sample of sodium hydroxide is weighed and dissolved in deionized water to give a solution of the approximate concentration desired. (Alternatively, a concentrated NaOH solution could be diluted.)

1. Samples of dried KHP are weighed into flasks and dissolved in deionized water.

2. A few drops of the appropriate acid–base indicator (phenolphthalein) are added to each sample.

3. A buret is rinsed with a little of the NaOH solution; then the buret is filled with NaOH solution.

4. Take an initial buret reading.

5. NaOH solution is titrated into the KHP samples until the first permanent pink color.

6. Take the final buret reading.

7. Using the molar mass of KHP, determine the moles of KHP present. This is equal to the moles of NaOH.

8. The difference in the buret readings is the volume of NaOH solution added (convert this to liters).

9. The molarity of the NaOH solution is the moles of NaOH divided by the liters of NaOH solution added.

10. (Repeat the procedure for each sample.)

Give yourself 2 points for this entire list, if the items are in order. If three or more items are in the wrong order or missing, you get only 1 point. You get 0 points for three or fewer items.

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You get 1 point for this graph. You get an additional point for noting that the equivalence point is greater than 7.

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You get 1 point for this graph. You must show two steps.

(d) There are several related ways to do this problem. One method is to split the sample into two portions. Titrate one portion to the equivalence point. Add the titrated sample to the untitrated sample, and add a volume of deionized water equal to the volume of NaOH solution added. The pH of this mixture is equal to the pKa of the acid (this corresponds to a half-titrated sample).

You get 1 point for anything concerning a half-titrated sample and an additional point for pH = pKa.

(e) The pH at the equivalence point must be close to the pKa of the indicator.

You get 1 point for this answer.

There are a total of 8 points possible.

Image Rapid Review

Reviewing the experiments should include looking at the synopsis, apparatus, calculations and comments as well as the appropriate concept chapters, if needed.

• Pay particular attention to any experiment you did not perform.

• Be familar with the equipment used in each experiment.

• Know the basic measurements required in each experiment.

• Know what values are measured and which are calculated.

• Pay attention to significant figures.

• Balances are used to measure the mass of a substance, not the moles.