foundations of applied chemistry chem 0012 lab...

Foundations of Applied Chemistry

CHEM 0012 Lab Manual

Revised: December 2010

1

Table of Contents

PagePeriodic Table 3

Lab Write-up Instructions 4

Lab 1: Safety Lab; Introduction to Volumetric and Weighing Techniques ........ 5

Lab 2: Factors Affecting Reaction Rates ........ 17

Lab 3: Equilibrium and Le Châtelier’s Principle ........ 44

Lab 4: Introduction to pH Measurements / Acid-Base Titration ........ 56

Lab 5: pH Measurements and Acid-Base pH Titration Curves ........ 69

Lab 6: Acid-Base Equilibria: Determination of Ka and Investigation of Buffers ........ 77

Lab 7: Determination of an Equilibrium Constant using Spectroscopy ........ 85

Lab 8: Electrochemistry: The Study of Corrosion in Metals ........ 94

Lab 9: Determination of Solubility Product Constants, Ksp, using Potentiometric Method

....... 103

2

Lab Write-up Instructions

1. COVER PAGE • Experiment number and title of experiment • Student's name and partner's name(s) • Date of the Experiment • Sign your lab report

2. DATA • original data must be included in the lab report • Enter data into data sheet directly IN INK, not on a piece of scrap

paper, nor on paper towels • neatly tabulated (watch significant figures, units ...) • Data should be checked and SIGNED by instructor before leaving

laboratory.

3. GRAPH (when required)

• Proper graph paper must be used. (ie - Do not use papers with squares, engineering papers ...)

• The graph should nearly fill a 8-1/2" x 11" page. • Computer generated graphs are acceptable. • Each graph must contain the following information:

a. Title of experiment, Experiment #, Graph # (if more than one

graph) b. Descriptive Title of Graph (by convention "Y versus X") c. Chemical Equation(s) with physical states d. Label for the axes with appropriate units

- If the axis is ln [Cu2+], then the label should be "ln [Cu2+] ([Cu2+] in moles/L)"

e. Show: (i) the points used for slope calculations (ii) slope calculation on graph (iii) units, if any.

4. CALCULATIONS

• show sample calculations

• Calculations can be done in the space provided on the data sheets.

• Report all numbers with proper number of significant figures.

4

Last updated 12/10/2009 3:15 PM

Lab 1: Safety Lab; Introduction to Volumetric and Weighing Techniques

Objectives:

1. Be aware of safety practices, procedures outlined in the safety video. 2. Introduction to WHMIS and MSDS. 3. Locate the laboratory emergency safety equipment and understand why and when to use

them. 4. Understand the tolerances of lab glassware, bottle-top dispenser and the analytical

balance. 5. Learn and practice volumetric and weighing lab techniques.

Introduction: Part A: Lab Safety

In this lab you will be watching a lab safety video. Safe working practices are essential and mandatory part of all work activities. You will be given a tour of the safety features in the lab. It is important that you know where to locate the emergency safety equipment and have an understanding of their use. Students are expected to act professionally in the lab environment. WorkSafeBC (the new name for Workers' Compensation Board of BC) is an independent agency governed by a Board of Directors appointed by government. Their core mandates are:

• prevent workplace injury illness and disease, • rehabilitate, and • provide fair compensation

In order for WorkSafeBC to be effective, their focus to promote healthy and safe workplaces through enforcement, consultation and education. In 1988, an emphasis on worker safety is launched with a public awareness campaign on Alcohol and Drug Abuse in the Workplace and the Workplace Hazardous Materials Information System (WHMIS). Read more about WorkSafeBC historical milestones.

5


WorkSafeBC website

WHMIS The Workplace Hazardous Materials Information System (WHMIS) provides information about many hazardous materials used in the workplace. WHMIS refers to these hazardous materials as controlled products. Under WHMIS, workers have the right to receive information about each controlled product they use---its identity, hazards, and safety precautions.

Classification

Each controlled product is classified into one or more of the six hazard classes, Class A to F. Once classified, they are assigned one or more of the appropriate hazard symbols. There are eight WHMIS hazard symbols. Workers need to recognize these symbols and recognize what they mean.

After a controlled product has been classified, the means to communicate health and safety information about the controlled products are via:

1. WHMIS labels

2. Material Safety Data Sheets (MSDSs) are used to communicate health and safety information.

Material Safety Data Sheets (MSDSs) A Material Safety Data Sheet (MSDS) provides both workers and emergency personnel with the proper procedures for handling or working with a particular substance. MSDS's include information such as physical data (melting point, boiling point, flash point), toxicity, health effects, first aid, reactivity, storage, disposal, protective equipment, and spill/leak procedures. These MSDS are particular use when a spill or an accident occurs.

3. WHMIS education and training programs. In a teaching environment where students are expected to handle controlled products, students need to be educated to ensure the understanding of WHMIS and the hazards of the controlled products that they work with. In a work place, education programs about WHMIS are far more extensive. Workers must be trained in safe work procedures for handling, storing, disposing of the controlled products, as well as emergency procedures in the event of an accident or spill.

6


Introduction: Part B: Introduction to Volumetric Glassware When performing a chemistry lab, the procedure may include the use of various glassware for measuring volumes. In order to accurately perform the lab procedures such that the accuracy of measurements is not compromised, an understanding of the different types of glassware is required. Each piece of glassware is made to certain specifications. That is, there is a maximum measurement error associated with the glassware known as the tolerance. For example, a flask which holds 5.00 mL of liquid has a tolerance of +/- 0.02 mL (or +/- 0.4 %). This means that the actual volume that the flask hold is in the range of 4.98 to 5.02 mL. In order to avoid introducing a significant error to the analytical result, the tolerance specifications of each piece of volumetric glassware must match the required accuracy of the procedure. The following is a summary of the various types of glassware and their tolerance.

Type of Glassware Target Volume

Tolerance Accuracy Range

Graduated Erlenmeyer flask

125 mL +/- 5% 125 mL +/- 6.25 mL Graduation interval=25 mL roughly accurate

Graduated beaker 100 mL +/- 5% 100 mL +/- 5 mL Graduation interval=10 mL roughly accurate

Graduated cylinder 10 mL +/- 0.5% 10 mL +/- 0.05 mL Graduation interval=0.1 mL accurate

Graduated cylinder 100 mL +/- 0.4% 100 mL +/- 0.40 mL Graduation interval=1 mL accurate

Volumetric flask 100 mL +/- 0.08% 100 mL +/- 0.08 mL No graduation interval very accurate

Volumetric pipette 25 mL +/- 0.12% 25.0 mL +/- 0.03 mL No graduation interval very accurate

Burette 50 mL +/- 0.1% 50 mL +/- 0.05 mL Graduation interval=0.10 mL very accurate

7


8


Mechanical Dispenser Target Volume

Tolerance Accuracy Range

Bottle-Top Dispenser

50 mL +/- 1-2% 50 mL +/- 51 mL

The student must be able to use the different glassware and bottle-top dispenser with proper lab techniques such that correct measurements can be made. In this lab, the proper lab techniques will be demonstrated. It is expected that students will practice these techniques until they become proficient with handling all the lab equipment.

9


Introduction: Part C: Introduction to the Analytical Balance

Weighing a sample is often the first step in many quantitative analytical methods. An analytical balance measures masses to within 0.0001 g. Balances are sensitive to drafts, changes in temperature, or the vibrations caused by moving people. The balances are stored in a separate room to minimize these variables and are placed on concrete tables.

Balances are very expensive and are sensitive to attack by corrosive chemicals. Do not take liquid into the balance room. When possible, chemicals should be added to the weighing container outside of the balance chamber. It is important that you clean up all chemical spills.

In this experiment, you will learn to use the balance properly and be aware of the common errors encountered in weighing.

10


Apparatus:

1. 1- 125 mL Erlenmeyer flask 2. 1- 100 mL graduated beaker 3. 1- 10 mL graduated cylinder 4. 1- 25 mL pipette 5. 1- 50 mL burette 6. analytical balance 7. a plastic vial 8. a pair of tongs 9. drying oven set at 70oC. 10. a small dessicooler 11. 100.0 mL. volumetric flasks 12. burette funnels 13. Pasteur pipettes & bulbs 14. pipette racks & bulbs 15. bottle-top dispensers

Solution:

a coloured solution

11


Procedure:

Part A - Lab Safety

Visit the links and look up the WHMIS symbols and classifications. Complete the datasheet.

Part B - Introduction to Volumetric Glassware

1. Your instructor will demonstrate the following volumetric measurement techniques: - Use of a Burette - Use of a Pipette - Use of a Volumetric Flask - Use of a Bottle-top Dispenser

Download the volumetric measurement techniques as reference.

2. Using the techniques demonstrated, carry out the following tasks: a. Acclimatize the burette with the coloured solution provided. Fill the burette with

the coloured solution.

b. Obtain approximately 50 mL of distilled water in a graduated beaker. Transfer 25.00 mL of distilled water with a volumetric pipette into an Erlenmeyer flask.

c. Make a 1:4 dilution of the coloured solution provided. Transfer 25.00 mL of the coloured solution into the volumetric flask. Use distilled water to fill the volumetric flask to make it up to the mark. Mix well.

d. Measure 7.10 mL of distilled water using your graduated cylinder. Use a Pasteur pipette to carefully add the last drops to bring the bottom of the meniscus to the 7.10 mark.

e. Read the preset volume of the bottle-top dispenser. Use a graduated cylinder and measure the volume dispensed by the bottle-top dispenser. The volume collected should be within 2% of the preset volume. If the volume being dispensed is outside the preset volume, make sure that the bottle-top dispenser is not pumping air and try again.

Ask your instructor to check your work before you clean the glassware.

Part C - Introduction to the Analytical Balance

1. Zero the balance. 2. Use a pair of tongs and transfer a plastic vial into the balance. Determine the mass of the

vial. 3. Use a pair of tongs and place the vial in a drying oven for 5 minutes. Remove the vial and

place it in a desicooler for transporting the warm vial to the balance room. Reweigh immediately while it is still warm. Record the change in weight every 30 seconds for the next 5 minutes or until the mass stabilizes.

4. Touch the vial with your hand. Roll the vial in your palm for 10 seconds. Reweigh the vial. 5. Record all the mass measurements in the datasheet.

12


Datasheet:

Part A: Lab Safety Go to the Worksafe BC website and identify the following hazard symbols.

WHMIS SYMBOLS Classification

13


Datasheet:

Part B: Introduction to Volumetric Glassware Proper techniques in using the following glassware:

burette

Instructor initial:

volumetric flask

Instructor initial:

pipette

Instructor initial:

graduated cylinder

Instructor initial:

bottle-top dispenser Instructor initial:

14


Datasheet:

Part C: Introduction to the Analytical Balance

Mass of clean vial: ______________________________________ g

Mass of warm vial:

___________________________ g (immediately from the oven)

___________________________ g (30 sec)

___________________________ g (60 sec)

___________________________ g (90 sec)

___________________________ g (120 sec)

___________________________ g (150 sec)

___________________________ g (180 sec)

___________________________ g (210 sec)

___________________________ g (240 sec)

___________________________ g (270 sec)

___________________________ g (300 sec)

Mass of the vial touched by your hands:

___________________________ g

15


Postlab Questions: Part A:

1. Look up the MSDS for nitric acid. (a) List 3 physical properties. (b) Which section of the MSDS can you find information on treatment when nitric acid is causing skin irritation? (c) What is the treatment for nitric acid causing skin irritation? (d) How should nitric acid be stored?

2. Cite the reference of your nitric acid MSDS source.

Part B:

1. What is the function/use for each of the following glassware? (a) graduated beaker (b) Erlenmeyer flask (c) graduated cylinder (d) volumetric flask (e) volumetric pipette (f) burette (g) bottle-top dispenser

Part C:

1. What can you conclude about weighing an object that is not at room temperature? 2. Compare the mass of the vial that is handled by tongs and the mass of the vial that is

handled by your hands. What is the mass difference and state some reasons to account for the difference in mass?

16


Lab 2: Factors Affecting Reaction Rates

Objectives:

To determine how concentration, temperature, and catalyst affect the speed of a chemical reaction.

Introduction: The rate law is a mathematical expression showing how the rate of the reacting species changes with concentrations. In this experiment, you will determine the rate law of the following reaction by observing how concentration and temperature affect the speed of the reaction.

S2O82- (aq) + 2 I- (aq) 2 SO4

2- (aq) + I2 (aq) (2-1)

persulfate ion

iodide ion

sulfate ion iodine

The factors, which affect the rate of this reaction are:

1. Concentrations of the reactants, S2O82- and I- ions (Part A)

2. Temperature (Part B) 3. Catalyst (Part C)

17


Introduction: Part A: Effect of Concentration on the Rate of a Reaction

Our method for measuring the rate of the reaction involves what is called a "clock" reaction. The rate law for this reaction is of the form:

rate = k [ S2O82- ]x [ I- ]y eqn (2-2)

where [ S2O82- ] and [ I- ] are the concentration in moles/L of persulfate ion and iodide ion

respectively, and k is the rate constant. The rate constant will have a unique value for any particular process at a given temperature. The magnitude of the rate constant will tell whether a reaction will proceed quickly. A small rate constant indicates a slow reaction. A large rate constant indicates a rapid reaction. The rate constant of a chemical reaction will vary with temperature.

Table (2-1) gives of some reactions and their rate equations. The order of a reaction is very useful because it allows us to predict the influence of concentration of the speed of the reaction. For a 1st order reaction (example (c)) doubling the concentration, doubles the reaction rate. But if a reaction is 2nd order (example (a)) doubling the concentration increases the reaction rate by 4 times.

REACTION RATE LAW OVERALL ORDER (a) 2 HI H2 + I2 Rate = k[HI]2 2

(b) 2 NO + H2 N2O + H2O Rate = k[NO]2 [H2]1 3

(c) 2 N2O5 4 NO2 + O2 Rate = k [ N2O5 ] 1

Table (2-1) - Some chemical reactions and their rate equations.

The exponent 'x' in eqn (2-2) is the order of the reaction with respect to the S2O82- ion. The

exponent 'y' is the order of reaction with respect to the I- ion. The overall order of the reaction is the sum of x+y. The powers to which the concentrations are raised, ‘x’ and ‘y’, MAY or MAY NOT be the same as the stoichiometric coefficients in the balanced equation (2-1). In general, the order of a reaction CANNOT be determined by inspecting the balanced chemical equation but must be determined experimentally.

In this experiment, a known amount of sodium thiosulfate, Na2S2O3, is added to the reaction mixture. As reaction (2-1) proceeds, it will start to form I2 (aq). The iodine formed will be consumed according to the following reaction.

I2 (aq) + 2 S2O32- (aq) → 2 I- (aq) + S4O62- (aq) (2-1-2)

After the S2O32- ions are exhausted, the formation of any more I2 will react with the starch to turn the solution blue.

18


Part B: Effect of Temperature on the Rate of a Reaction

For almost all reactions, an increase in temperature will lead to an increase in the reaction rate. An increase in temperature will increase the average kinetic energy of the molecules. This will lead to an increase in the number of collisions per unit time. But collision alone is an insufficient criterion for the production of the products. The kinetic energy of the colliding molecules must be greater than the required minimum energy level before the reactants can be converted to products. This energy level is called the activation energy, Eact. Activation energy is unique for a particular chemical reaction.

If a reaction is investigated at a number of different temperatures, the rate constant, k, usually shows quite a dramatic increase, often several orders of magnitude, over a relatively small temperature range. For many reactions there is approximately a two- to three-fold increase in rate for every 10oC rise in temperature. For a given temperature, reactions that have large activation energies would be slower than the ones that have smaller activation energies.

The relationship between the temperature and the rate constant ‘k’ is given by the Arrhenius equation,

eqn (2-3)

or

eqn (2-4)

where

A is the pre-exponential or frequency factor, a constant related to the collision frequency

R is the gas constant (8.314 J / K mole) T is the absolute temperature (K) k is the rate constant at temperature T

Eact is the activation energy, the energy required by the reacting species for their collisions to be effective (ie - those that lead to the formation of products)

Equation (2-4) shows that 'ln k' is a linear function of the reciprocal absolute temperature. The activation energy can be determined experimentally by measuring the rate constant, k, at several different temperatures, T. A garph of ln k vs. 1/T yields a straight line with a slope of (– Eact / R) and an intercept of ln A.

Part C: Effect of a Catalyst on the Rate of a Reaction

A catalyst is a substance which increases the rate of a reaction, but remains unchanged in the process. A catalyst may function by providing an alternate path for which the reactants come together. In order for the rate of of the catalyst-assisted reaction to increase, the alternate pathway will have a lower activation energy, Eact, and therefore proceeds more rapidly.

In this experiment, a small quantity of Cu2+ is introduced to the persulfate-iodide clock reaction as a catalyst.

Cu2+ S2O8

2- (aq) + 2 I- (aq) 2 SO42- (aq) + I2 (aq) (2-5)

19


Apparatus:

1. Water baths set at: 20oC, 30oC, 40oC 2. lead donuts 3. 22 - 125 mL Erlenmeyer flasks (14 for Part A, 6 for Part B, 2 for Part C) 4. 2 plastic buckets to make ice slurry 5. ice 6. Timers that display seconds 7. Alcohol thermometers

Solutions:

1. 0.00500 M Na2S2O3 – in dispenser 2. 0.200 M KI – in dispenser 3. 0.100 M (NH4)2S2O8 – in dispenser 4. 0.1 M CuSO4 5. starch indicator

20


Procedure:

Part A - Effect of Concentration on Reaction Rate

Note: 1. All Erlenmeyer flasks provided are CLEAN and DRY. Use them as is.

2. Part A is carried out at 20oC 1. The following solutions are available in bottle-top dispensers. Pay attention to the preset volumes and dispense the proper amounts into the Erlenmeyer flasks.

(i) 0.00500 M Na2S2O3

(ii) 0.200 M KI

(iii) 0.100 M (NH4)2S2O8

(iv) distilled H2O

2. Obtain fourteen 125 mL Erlenmeyer flasks. Label them as follows:

A1, A2, A3, A4, A5, A6 and A7 B1, B2, B3, B4, B5, B6 and B7

3. Prepare the seven ‘A’ solutions according to Table (2-2).

Note: The total volume of each Erlenmeyer flask is 30.00 mL and 3 drops of starch indicator

Table (2-2) - Contents of the seven 'A' solutions.

SOLUTION 0.00500 M Na2S2O3

(mL)

0.200 M KI

(mL)

distilled H2O (mL)

3 % starch indicator

A1 10.0 20.0 0.0 3 drops

A2 10.0 20.0 0.0 3 drops

A3 10.0 20.0 0.0 3 drops

A4 10.0 20.0 0.0 3 drops

A5 10.0 15.0 5.0 3 drops

A6 10.0 10.0 10.0 3 drops

A7 10.0 5.0 15.0 3 drops

21


4. Prepare the seven ‘B’ solutions according to Table (2-3).

Note: The total volume of each Erlenmeyer flask is 20.00 mL Table (2-3) – Contents of the seven ‘B’ solutions.

SOLUTION 0.100 M (NH4)2S2O8

(mL)

distilled H2O (mL)

B1 5.0 15.0

B2 10.0 10.0

B3 15.0 5.0

B4 20.0 0.0

B5 20.0 0.0

B6 20.0 0.0

B7 20.0 0.0 5. Use lead donuts to stabilize the fourteen Erlenmeyer flasks in the 20oC water bath. Allow the flasks to come to thermal equilibrium by leaving them in the bath for at least 5 minutes. Measure the temperature of the water bath with a thermometer. Record the actual temperatures on the data sheet. 6. Pour the content of solution A1 rapidly into B1 while swirling. Start the timer immediately and record the time (in seconds) for the appearance of the blue colour. For the duration of the reaction, keep swirling the flask containing the combined solutions and keep the flask immersed in the 20oC water bath. 7. Repeat step 6 for the remaining 6 pairs of solutions, pouring A2 into B2, A3 into B3, A4 into B4, A5 into B5 and A6 into B6. Record the time (in seconds) for the appearance of the blue colour.

22

Last updated 12/13/2010 4:47:00 PM 23

Procedure: Part B - Effect of Temperature on Reaction Rate


2. Part B is carried out at 0oC, 20oC, 30oC, 40oC 3. Four pairs of A4/B4 solutions will be used at these temperatures.

1. Obtain eight 125 mL Erlenmeyer flasks. 2. Prepare four 'A4' solutions and four 'B4' solutions by following instructions given in Table (2-2) and Table (2-3). 3. Obtain two plastic buckets and prepare an ice slurry in each bucket. Use lead donuts to stabilize the Erlenmeyer flasks. Put one pair of 'A4/B4' solutions in the 0oC ice bath. Put the second pair of solution in the 20oC water bath, the third pair of solution in the 30oC water bath and the fourth pair of solution in the 40oC water bath. Allow the flasks to immerse in the bath for 5 minutes to come to thermal equilibrium. 4. Measure the temperatures of the ice slurry and water baths with a thermometer. Record the actual temperatures on the data sheet. 5. Pour the content of solution A4 rapidly into B4 while swirling. Start the timer immediately and record the time (in seconds) for the appearance of the blue colour. For the duration of the reaction, keep swirling the flask containing the combined solutions and keep the flask immersed in the water bath. 6. Repeat step 5 for the remaining three pairs of solutions.


Procedure:

Part C - Effect of a Catalyst on Reaction Rate


1. Obtain two 125 mL Erlenmeyer flasks.

2. Prepare one 'A4' solutions and one 'B4' solutions by following instructions given in Table (2-2) and Table (2-3).

3. To solution A4 add 1 drop of 0.1 M CuSO4·5H2O solution.

4. Use lead donuts to stabilize the two Erlenmeyer flasks in the 20oC water bath. Ensure that the temperature of the water bath is the same as the temperature used in Part A. Allow the flasks to come to thermal equilibrium by leaving them in the bath for at least 5 minutes. Measure the temperature of the water bath with a thermometer. Record the actual temperatures on the data sheet.

5. Pour the content of solution A4 rapidly into B4 while swirling. Start the timer immediately and record the time (in seconds) for the appearance of the blue colour. For the duration of the reaction, keep swirling the flask containing the combined solutions and keep the flask immersed in the 20oC water bath.

24


Datasheet:

Part A - Effect of Concentration on Reaction Rate

Temperature: _____________________________________

Experiment Solution Time (sec) 1 A1/B1

2 A2/B2

3 A3/B3

4 A4/B4

5 A5/B5

6 A6/B6

7 A7/B7

25

Last updated 12/13/2010 4:53:00 PM 26

Datasheet:

Part B - Effect of Temperature on Reaction Rate

Experiment Solution Time (sec)

Temperature (oC)

8 (in ice)

A4/B4

9 (near 20oC)

A4/B4

10 (near 30oC)

A4/B4

11 (near 40oC)

A4/B4

Last updated 12/13/2010 4:53:00 PM 27

Datasheet:

Part C - Effect of a Catalyst on Reaction Rate

Temperature: _____________________________________


12 A4/B4


Postlab Questions:

Part A - Effect of Concentration on Reaction Rate 1. Transfer data Experiments 1, 2, 3, and 4 from the datasheet to Table (2-4) below and calculate:

• the initial concentration of [S2O82-] and [I -],

• the concentration of I2 formed to consume the S2O32- added,

• the rate of formation of I2.

(a) Show a sample calculation for Experiment 1 in the space provided. Enter the calculated results for Experiment 1 into the first row of Table (2-4) below.

1. Calculate the initial concentration of [S2O82-] for Experiment 1.

2. Calculate the initial concentration of [I -] for Experiment 1. [Note: The concentration of [I -] is the same for Experiments 1, 2, 3, and 4]

3. (a) Calculate the concentration of I2 formed to consume the S2O32- added.

(b) Calculate the rate of formation of I2 for Experiment 1.

28

Last updated 12/13/2010 4:56:00 PM 29

(b) Repeat the above calculations for Experiments 2, 3, and 4 and complete Table (2-4).

Table (2-4): Summary of calculations for Experiments 1 to 4.

Temperature of Experiments: _____________________________________


[S2O82-]

(moles/L) [I-]

(moles/L) Rate of

Formation of I2 (M/sec)

1

A1/B1

2

A2/B2

3

A3/B3

4

A4/B4

2. (a) Determine the order of the reaction with respect to [ S2O8

2- ] using the calculated results from Experiments 1 and 2 from Table (2-4). [Note: The order of the reaction with respect to [ S2O8

2- ] is the value of 'x' in the rate equation rate = k [ S2O8

2- ]x [ I- ]y .] Show a sample calculation in the space provided. Enter your calculated result in the first row of Table (2-5). Sample calculation of 'x' using Experiments 1 and 2 calculated values.


(b) Repeat the calculation of 'x', using the other pairs of experiments as indicated in Table (2-5) below. Calculate the average value of 'x' . The average value of 'x' should be rounded the nearest integer. Complete Table (2-5). Table (2-5): Calculation of 'x' based on Experiments 1, 2, 3 and 4.

Experiment pairs

'x', The order of the reaction with respect to [S2O8

2-]

Average value of

'x'

1 and 2

2 and 3

3 and 4

1 and 3

1 and 4

2 and 4

The average value of ‘x’

___________

Rounded to the nearest integer, the order of the reaction with respect to

[S2O82-] is

_____________

3. Transfer the calculated results for Experiment 4 from Table (2-4) to Table (2-6).

30


4. Transfer data for Experiments 5, 6 and 7 from the datasheet to Table (2-6) and calculate:

• the initial concentration of [S2O82-] and [I -],

• the concentration of I2 formed to consume the S2O32- added,

• the rate of formation of I2.

(a) Show a sample calculation for Experiment 7 in the space provided. Enter the calculated results for Experiment 7 into the 4th row of Table (2-6) below.

1. Calculate the initial concentration of [S2O82-] for Experiment 7.

[Note: The concentration of [S2O82-] is the same for experiments 4, 5, 6 and 7]

2. Calculate the initial concentration of [I -] for Experiment 7.

3. (a) Calculate the concentration of I2 formed to consume the S2O32- added.

(b) Calculate the rate of formation of I2 for Experiment 7.

(b) Repeat the above calculations for Experiments 5, and 6 and complete Table (2-6).

31

Last updated 12/13/2010 4:57:00 PM 32

Table (2-6): Summary of calculations for Experiments 5 to 7.

Expt Solution Time (sec)

[S2O82-]

(moles/L) [I-]

(moles/L) Rate of


4

A4/B4

5

A5/B5

6

A6/B6

7

A7/B7

5. (a) Determine the the order of the reaction with respect to [ I- ] using the calculated results from Experiments 4 and 5 from Table (2-6). [Note: The order of the reaction with respect to[ I- ] is the value of 'y' in the rate equation rate = k [ S2O8

2- ]x [ I- ]y .]


Show a sample calculation in the space provided. Enter your calculated result in the first row of Table (2-7) below. Sample calculation of 'y' using Experiments 4 and 5 calculated values.

33

Last updated 12/13/2010 4:58:00 PM 34

(b) Repeat the calculation of 'y', using the other pairs of experiments as indicated in Table (2-7) below. Calculate the average value of 'y'. The average value of 'y' should be rounded to the nearest integer. Complete Table (2-7) below.

Table (2-7): Calculation of 'y' based on Experiments 4, 5, 6, and 7.

Experiment pairs

'y', The order of the reaction

with respect to [I -]

Average value of

'y'

4 and 5

The average value of 'y'

____________

Rounded to the

nearest integer, the order of the reaction with respect to [I-] is

___________

5 and 6

6 and 7

4 and 6

4 and 7

5 and 7

6. Based on your experimental results, summarize your results for the reaction

S2O82- (aq) + 2 I- (aq) 2 SO4

2- (aq) + I2 (aq) (2-1)

in Table (2-8).


Table (2-8): The experimentally determined rate law.

Experimentally Determined

Rate Law =

Order of reaction with respect to [ S2O8

2- ], ‘x’ =

Order of reaction with respect to [ I - ], ‘y’ =

Overall reaction order =

7. Using the experimentally determined rate law (Table 2-8), calculate the rate constant, k, for each experiment and the average rate constant, k, for the reaction. [Note: Report k with the correct unit.] Show a sample calculation in the space provided. Enter your calculated result in the first row of Table (2-9).

Sample calculation of 'k' for Experiment 1.

35


Table (2-9): Calculation of the rate constant 'k' based on Experiments 1 to 7.

Experiment Rate constant, k.

(Report k with the appropriate unit.)

Average value of

k

1

2

3

4

5

6

7

The average rate constant is

_______________________

(Report k with the appropriate unit.)

for the temperature

________oC

36


Postlab Questions:

Part B - Effect of Temperature on Reaction Rate 1. Transfer from Part A (Table 2-4) the calculated concentrations of S2O8

2- and I- for the A4/B4 solutions.

[ S2O82- ] (moles/L)

[ I- ] (moles/L)

2. (a) Using the rate law determined from Part A, calculate:

• the rate of formation of I2, • the concentration of I2 formed to consume the S2O3

2- added, • the rate constant, k

for Experiment 8. Show a sample calculation for Experiment 8 in the space provided. Enter the calculated results for Experiment 8 into the first row of Table (2-10) below. 1. Calculate the rate of formation of I2 for Experiment 8 (in ice).

2. Calculate the rate constant, k, for Experiment 8 (in ice).

38

Last updated 12/13/2010 5:05:00 PM 22

(b) Repeat the above calculations for Experiments 9, 10, and 11 and completeTable (2-10) below. \ Table (2-10): Summary of calculations for Experiments 8 to 10.

Experiment [ S2O82-] and [ I- ] in the

A4/B4 reaction mixture (moles/L)

Temperature(oC)

Time (sec)

Rate of Formation of I2

(mmole/mL sec)

Rate constant, k

(Report k with the appropriate

unit.) 8

[ S2O82- ] = _____________

[ I- ] = _________________

9

10

11

Last updated 12/13/2010 5:10:00 PM 39

3. Using the results entered in Table (2-10), complete Table (2-11). Table (2-11: Summary of ln k and and temperature results for Experiments 8, 9, 10, and 11. Experiment ln k Temperature

(oC) Transfer from Table (2-10)

Temperature(K)

1/Temperature (1/K)

8

9

10

11

4. Graph the values of ln k (on the y-axis) and the values of 1/T (on the x-axis). Draw the best straight line through the points (i.e. - trendline). Your graph must display the following information:

1. The title of the graph. The general format of the title is "{Y-axis} versus{ X-axis} for the {chemical equation}". Substitute the curly bracket information with appropriate information for your graph.

2. Labelled y-axis. Include appropriate units. For graphs involving log or ln functions, since log or ln functions are unitless, the unit of the log or ln argument should be written in bracket. For example, if the y-axis is ln [A], where [A] is the concentration of species A, the label of the y-axis would be ln [A] ([A] moles/L)

3. Labelled x-axis. See above for log or ln functions. 4. Trend line, is the equation of the least-squares regression line. From the trend line, write:

o the slope of the line reported to proper significant figures and with the appropriate unit.

o the intercept of the line reported to proper significant figures and with the appropriate unit.

5. R2, is the coefficient of determination. This measures the percentage of variation in the dependent variable given by the trend line. It has a value between zero and one. A value close to one indicating a good fit. A value of one indicating a best fit.

6. Conditions of the experiment. The temperature and pressure at which the experiment was carried out.


5. Recall from equation (2-4),

eqn (2-4)

Determine the activation energy, Eact, and the pre-exponential or frequency factor, A. Show your calculation in the space provided. Determine Eact.

Determine the pre-exponential or frequency factor, A.

40

Last updated 12/13/2010 5:14:00 PM 41

25

Postlab Questions:

Part C - Effect of a Catalyst on Reaction Rate 1. Transfer:

data from Experiment 12 the calculated initial concentration of [S2O8

2-] and [I -] fromTable (2-6)

to Table (2-12). 2. Show your calculation of the rate of formation of I2 in the space provided. Enter the calculated results into the first row of Table (2-12). (a) Calculate the concentration of I2 formed to consume the S2O2

2- added. (b) Calculate the rate of formation of I2 for Experiment 12.

Last updated 12/13/2010 5:16:00 PM 42

Table (2-12): Summary of calculations for Experiment 12.

Temperature of Experiment: _____________________________________

Expt Solution Time (sec)

[S2O82-]

(moles/L) [I-]

(moles/L) Rate of


12

A4/B4

9

A4/B4


3. Compare the rate of formation of I2 for Experiment 11 with that which was obtained for Experiment 4, where no catalyst was added. In the space provided calculate how much faster the reaction is with the catalyst. By how much does the catalyst cause the rate of the reaction to increase?

43


Lab 3: Equilibrium and Le Châtelier’s Principle

Objectives:

To explore the effect of changing the concentrations of reactants and products on the equilibrium composition of four equilibrium systems.

Introduction: Most chemical reactions do not result in a 100% yield of products based on the stoichiometry of the reaction. This is usually due to the equilibrium state that is reached when the forward rate of reaction equals the rate of the reverse reaction. In this lab, the effect of qualitative changes on a number of reactions at equilibrium will be studied.

Le Châtelier’s Principle states that

“If a change in conditions is imposed on a system at equilibrium, the equilibrium position will shift in a direction that

tends to reduce that change in conditions.”

For example, the change in conditions could be either the temperature or concentration and the effects observed. It should be noted that for a system, there exists many equilibrium positions but only have one equilibrium constant at a specific temperature.

Click on this link to view an animation of Le Châtelier’s Principle

(Flash plugin required)

44


In this experiment, we will study the equilibrium of four systems and observe the reaction of the equilibrium systems as predicted by Le Châtelier’s Principle. The four systems are:

Part A - The Equilibrium of Co(II) Complex Ions The element cobalt can form compounds in two different oxidation states, +2 and +3. The +2 state is more common. The chloro complex of cobalt (II), CoCl42-, is tetrahedral while the aquo complex of cobalt (II), Co(H2O)6

2+, is octahedral. Both of these complexes exhibit different colours. Cobalt complexes are used as drying agents with the colour change indicating when the drying agent should be changed. The equilibrium reaction is: Co(H2O)6

2+ (aq) + 4 Cl- (aq) CoCl42- (aq) + 6 H2O (l) ΔH = +50 kJ (3-1) pink blue Part B - The Equilibrium of the thiocyano-iron(III) complex ion When colourless aqueous solutions of iron (III) ion, Fe3+, and thiocyanate ion, SCN-, are combined, the reaction that occurs produces the thiocyanoiron (III) complex ion, FeSCN2+, which is responsible for the equilibrium mixture's deep red colour.

Fe3+ (aq) + SCN- (aq) FeSCN2+ (aq) (3-2) colourless colourless red-brown

The colour of the thiocyanoiron (III) complex ion, FeSCN2+, solution will indicate how the equilibrium system is being affected. Part C - The Equilibrium of a Mg+2 precipitate Reactions which form precipitates are written as an equilibrium reaction using the solubility product. If there is a precipitate MX, then the Ksp expression is:

MX (s) M+(aq) + X-(aq) (3-3)

Part D: The Equilibrium of an Acid-Base Indicator An acid-base indicator can be used to observe an equilibrium reaction. Indicators are weak acids which show one colour in the acid form, HInd, and another colour in the basic form, Ind-. At the pKa of the indicator there is equimolar amounts of the conjugate forms and the observed colour is a mixture of the two. Bromothymol blue is a yellow-green-blue indicator which has a pKa of 7.0. The reaction of the indicator bromothymol blue can be illustrated as follows:

HInd (aq) H+ (aq ) + Ind- (aq) (3-4) yellow blue

45


Apparatus:

1. Large test tubes 2. Hot plate (1 per student)

Solutions:

1. 0.1 M CoCl2·6H2O (15 mL per student) 2. concentrated HCl (place in fumehood) 3. 0.1 M AgNO3 4. 0.1 M Mg(NO3)2 5. 6M NH4OH 6. NH4Cl solid 7. pH 7 buffer solution (3 mL per student) 8. Bromothymol blue indicator 9. 0.1 M Fe(NO3)3 in 0.1M HNO3 (2 mL per student) 10. 1 M Fe(NO3)3 in 0.1M HNO3 (0.1 mL per student) 11. 0.1 M KSCN in 0.1M HNO3 (2 mL per student) 12. 1 M KSCN in 0.1M HNO3 (0.1 mL per student) 13. 0.1 M NaCl 14. 6M NaOH 15. ice 16. 1 M HCl 17. 1 M NaOH

46


Procedure:

Part A - The Equilibrium of Co(II) Complex Ions

1. Obtain four clean and dry test tubes and label them Test tube #1, #2, #3 and Control.

2. Record the initial colour of the stock CoCl2· 6H2O solution.

3. Pour 12 mL of 0.1 M CoCl2· 6H2O into a clean, dry 50 mL beaker. In the fume hood, add concentrated HCl drop wise and mix with a glass rod until a permanent colour change is observed. Record the observed colour.

4. Divide the solution equally into four test tubes. Use the Control test tube for colour comparison.

5. In Test tube #1 add water with mixing until a colour change is produced. Record the observed colour.

6. Heat Test tube #1 in a hot water bath (add boiling chips) and you should see a colour change. (If you don’t then you have likely added too much water. Try again with another sample.) Record the observed colour.

7. Cool Test tube #2 in an ice water bath and record the observations. Record the observed colour. Keep your ice bath for Part B.

8. Heat Test tube #3 in a hot water bath and record your observations. Record the observed colour. Keep your water bath for Part B.

9. Dispose of the cobalt solutions in the waste bottle.

Part B - The Equilibrium of the thiocyano-iron(III) complex ion You should compare the colour of each tube with the reference Test tube #1. Note the colour of the 0.1 M Fe(NO3)3.

1. In a 100 mL beaker, combine: o 1.5 mL of 0.1 M Fe(NO3)3, ** USE THE CORRECT CONCENTRATION** o 1.5 mL of 0.1 M KSCN, ** USE THE CORRECT CONCENTRATION** o 50 mL H2O.

Pour 5 mL of the solution into nine numbered test tubes.

2. Add two drops of H2O to Test tube #1, which will serve as reference for colour. Record your observations.

3. Add two drops of 1M Fe(NO3)3 to Test tube #2. Record your observations.

4. Add two drops of 1M KSCN to Test tube #3. Record your observations.

5. Add 8 drops 6M NaOH to Test tube #4. The precipitate Fe(OH)3 will take a few minutes to form. Record your observations.

6. Add 4 drops of AgNO3 to Test tube #5. The precipitate is AgSCN. Record your observations.

7. Add 4 drops of 0.1 M NaCl to Test tube #6. Record your observations. 47


8. Place Test tube #7 in an ice water bath. Record observations.

9. Place Test tube #8 in a boiling water bath. Record observations.

10. Add 1 mL of distilled water to Test tube #9. Record your observations.

11. Now add an additional 4 mL of water and record your observations.

12. Dispose of the reagents as instructed.

Part C - The Equilibrium of a Mg2+ precipitate

1. Into a test tube add: o 1 mL water, o 2 drops of 0.1 M Mg(NO3)2, and o 3 drops of 6M NH4OH.

Record your observations.

2. Add a small amount (1/4 spatula) of solid NH4Cl to the test tube and mix to dissolve. Record your observations. What is the product? Consult a solubility table.

Part D: The Equilibrium of an Acid-Base Indicator

1. Obtain a pH 7 buffer solution and pour 3 mL into a 50 mL beaker. Add 5 drops of bromothymol blue indicator. Record your observations.

2. Add 1 M HCl drop wise with mixing until the solution is acidic and the indicator shows a colour change. Record your observations.

3. Add 1M NaOH drop wise to return to the original colour and continue until the solution is basic and a new colour is reached. Record your observations.

48


Datasheet:

Part A - The Equilibrium of Co(II) Complex Ions

Solutions Colour

CoCl2· 6H2O (step 2)

CoCl2· 6H2O + HCl (step 3)

Test tube #1 + H2O (step 5)

Test tube #1 + H2O + heat (step 6)

Test tube #2 + ice water bath (step 7)

Test tube #3 + hot water bath (step 8)

49


Datasheet:

Part B - The Equilibrium of the thiocyano-iron(III) complex ion

Test tubes Observations

Test tube #1 - Fe(NO3)3 + KSCN + H2O(step 2)

Test tube #2 + 1 M Fe(NO3)3 (step 3)

Test tube #3 + 1M KSCN (step 4)

Test tube #4 + 6 M NaOH (step 5)

Test tube #5 + 4 drops AgNO3 (step 6)

Test tube #6 + 0.1 M NaCl (step 7)

Test tube #7 + ice (step 8)

Test tube #8 + boiling water bath (step 9)

Test tube #9 + 1 mL distilled water (step 10)

Test tube #9 + 5 mL distilled water (step 11)

50


Datasheet:

Part C - The Equilibrium of a Mg2+ precipitate

Test tubes Observations

Mg(NO3)2 + NH4OH (step 1)

Test tube + NH4Cl (step 2)

Part D: The Equilibrium of an Acid-Base Indicator

Solutions Observ ations

pH 7 Buffer + bromothymol blue (HInd) (step 1)

Solution + HCl (step 2)

Solution + NaOH (step 3)

51


Postlab Questions:

Part A - The Equilibrium of Co(II) Complex Ions 1. Write the equilibrium equation for step 3 when concentrated HCl was added to CoCl2· 6H2O. Equilibrium reaction: 2 .In Step 5, circle the shift in equilibrium observed (if any) when water is added to Test tube #1.

no shift in equilibrium

shifts to the right

shifts to the left

3. In Step 5, the colour of the solution is due to the presence of which ion? 4. In Step 6, circle the shift in equilibrium observed (if any) when Test tube #1 is heated.


shifts to the right

shifts to the left

5. Circle the correct response. Which statement is true?

heating favours an exothermic process

heating favours an endothermic process

6. In Step 6, the colour of the solution is due to the presence of which ion? 7. In Step 7, circle the shift in equilibrium observed (if any) when Test tube #2 is in the ice bath.


shifts to the right

shifts to the left


cooling favours an exothermic process

cooling favours an endothermic process

52


9. In Step 7, the colour of the solution is due to the presence of which ion? 10. In Step 8, circle the shift in equilibrium observed (if any) when Test tube #3 is in the hot water bath.


shifts to the right

shifts to the left


heating favours the equilibrium products

heating favours the equilibrium reactants

12. In Step 8, the colour of the solution is due to the presence of which ion? Part B - The Equilibrium of the thiocyano-iron(III) complex ion 1. Write the equilibrium equation for step 1 when Fe(NO3)3, KSCN and water was combined. Equilibrium reaction:

2. In Step 2, circle the shift in equilibrium observed (if any) in Test tube #1 when 2 drops of water is added.


shifts to the right

shifts to the left

3. In Step 3, circle the shift in equilibrium observed (if any) when Test tube #2 when 2 drops of 1 M Fe(NO3)3 is added.


shifts to the right

shifts to the left

4. In Step 4, circle the shift in equilibrium observed (if any) when Test tube #3 when 2 drops of 1 M KSCN is added.


shifts to the right

shifts to the left

53


5. In Step 5, circle the shift in equilibrium observed (if any) when Test tube #4 when 8 drops of 6 M NaOH is added.


shifts to the right

shifts to the left

6. In Step 6, circle the shift in equilibrium observed (if any) when Test tube #5 when 4 drops of AgNO3 is added.


shifts to the right

shifts to the left

7. In Step 7, circle the shift in equilibrium observed (if any) when Test tube #6 when 4 drops of 0.1 M NaCl is added.


shifts to the right

shifts to the left

8. In Step 8, circle the shift in equilibrium observed (if any) when Test tube #7 when it is placed in ice water bath.


shifts to the right

shifts to the left

9. In Step 9, circle the shift in equilibrium observed (if any) when Test tube #8 when it is placed in hot water bath.


shifts to the right

shifts to the left

10. In Step 10, circle the shift in equilibrium observed (if any) when Test tube #9 when 1 mL of water is added.


shifts to the right

shifts to the left

11. In Step 11, circle the shift in equilibrium observed (if any) when Test tube #9 when 5 mL of water is added.


shifts to the right

shifts to the left

12. Write the equilibrium constant expression and explain the observation when water is added.

54


Part C -The Equilibrium of a Mg2+ precipitate 1. Write the equilibrium equation (written as Ksp) for step 1 when Mg(NO3)2 and NH4OH was combined. Equilibrium reaction: 2 .In Step 2, circle the shift in equilibrium observed (if any) in the test tube when NH4Cl is is added.


shifts to the right

shifts to the left

Part D -The Equilibrium of an Acid-Base Indicator 1. Write the equilibrium equation (written as Ka) for step 1 when the pH 7 buffer and 5 drops of bromothymol blue (HInd) is combined. Equilibrium reaction: 2 .In Step 2, circle the shift in equilibrium observed (if any) in the test tube when 1 M HCl is added to the solution.


shifts to the right

shifts to the left

3 .In Step 3, circle the shift in equilibrium observed (if any) in the test tube when 1 M NaOH is added to the solution.


shifts to the right

shifts to the left

55

Lab 4: Introduction to pH Measurements / Acid-Base Titration

Objectives:

1. obtain a quick measure of whether a sample is acidic or basic with the use of broad-range pH paper.

2. carry out volumetric titration using proper volumetric and titration techniques. 3. determine the concentration of a solution using volumetric titration.

Introduction: One of the most important analysis in industry is the measurement of acidity, neutrality, or basicity in chemical processes. Typical industries employing such tests are in resin manufacturing, pulp and paper, mining, food production, pharmaceuticals, and environmental monitoring. Chemical feed stocks, intermediate and finished products, and effluents (waste water, etc.) are all monitored.

Acid-base titration is a quantitative analysis used to determine the concentration of an unknown acid or base solution. In this experiment, we will know when to end the titration by adding a pH indicator to one of the solutions and observing a permanent colour change.

Last updated 06/01/2008 12:45 AM 56

Introduction: Part A: Introduction to pH Measurements Types of Acids and Bases Acids, by definition, are proton donors, H+. Acids such as HCl, HNO3 and HClO4 are classified as strong acids. They are strong acids because in solution, they dissociate 100% into its ions. The molecular form of the acid does not exist. The dissociation process can be represented as follows:

HCl (aq) + H2O (l) H3O+ (aq) + Cl- (aq)

HNO3 (aq) + H2O (l) H3O+ (aq) + NO3- (aq)

HClO4 (aq) + H2O (l) H3O+ (aq) + ClO4- (aq)

Weak acids will also release H+ ions in solution, but they do not release all of the available H+ ions. Weak acids ionize partially in water. Most of H+ is still in the original acid molecule. The less amount of H+ released, the weaker the acid. Examples of weak acids are acetic acid, CH3COOH, or vinegar, citric acid and oxalic acid.

CH3COOH (aq) + H2O (l) H3O+ (aq) + CH3COO- (aq)

The double-headed arrow is used to indicate that only a few H3O+ ions are formed. Bases, by definition, are proton acceptors. Bases such as NaOH and Ba(OH)2 are classified as strong bases. They are strong bases because in solution, they dissociate 100% into its ions. The molecular form of the acid does not exist. The dissociation process can be represented as follows:

NaOH (aq) + H2O (l) Na+ (aq) + OH- (aq)

Ba(OH)2 (aq) Ba2+ (aq) + 2 OH- (aq)

Weak bases will also release OH- ions in solution. Weak bases ionize partially in water. The less amount of OH- released, the weaker the base. Examples of weak bases are ammonia, NH3 (aq), and ethylamine, CH3CH2NH2 (aq).

NH3 (aq) NH4+ (aq) + OH- (aq)

CH3CH2NH2 (aq) CH3CH2NH3+ (aq) + OH- (aq)

Last updated 06/01/2008 12:45 AM 57

Click on this link to view animation

of strong acid and weak acid ionization


The pH Concept Most common pH measurements are carried out in aqueous solution (i.e. in water). Water is attracted to the proton (H+) to form the hydronium ion, H3O+. The pH scale is frequently used instead of acid concentration. pH is defined as the negative log10 of hydrogen ion concentration, where [H3O+] is moles/L.

pH = -log[H3O+] (4-1)

For example, to calculate the pH of a 0.10 M HCl solution, we substitute [H3O+] = 0.10 M because HCl is a strong acid, all the HCl will completely dissociate to form H3O+ and Cl- ions.

pH = -log 0.10 = 1.0

The pH scale ranges from 0 to 14. When pH is 7, the solution is neutral. When pH is less than 7, the solution is acidic. When pH is greater than 7, the solution is basic. Acidic solutions contain more H3O+ ions than OH- ions. Basic solutions contain more OH- ions than H3O+ ions. Neutral solutions contain equal amounts of H3O+ ions and OH- ions. pH Indicators Acids and bases are colourless ions in solutions, but pH indicators are either weak organic acids/bases that change colour over a narrow pH range (see Table 4.1), Some colour change on the acidic side of the pH scale (i.e. pH < 7) and some change colour on the basic side (i.e. pH >7). Most indicators have a predominant colour change. Table 4.1 - pH range for the colour change of some common pH indicators.

Indicator pH Colour change pH Colour change thymol blue 1.2 - 2.8 red yellow 8.0 - 9.6 yellow blue methyl orange 3.2 - 4.2 red orange-yellow methyl red 4.8 - 6.0 red yellow bromothymol blue 6.0 - 7.6 yellow blue phenol red 6.6 - 8.0 yellow red cresol red 7.0 - 8.8 yellow red 0.4 - 1.8 red yellow phenolphthalein 8.2 - 9.8 colourless pink thymolphthalein 9.4 - 10.5 colourlesss blue alizarin yellow 10.1 - 12 yellow red If the solution pH is 6.4, the first three indicators will all be yellow in colour.

Last updated 06/01/2008 12:45 AM 58

http://highered.mcgraw-hill.com/olcweb/cgi/pluginpop.cgi?it=swf::525::530::/sites/dl/free/0072857684/322633/acid_ionization.swf::Acid%20Ionization




http://www.adobe.com/shockwave/download/download.cgi?P1_Prod_Version=ShockwaveFlash

Introduction: Part B: Acid-Base Titration 1. Stoichiometric Ratio: Moles versus Equvalents The general acid-base neutralization reaction is

Acid + Base Salt + Water

An example of acid-base neutralization reaction is:

Example 1: HCl (aq) + NaOH (aq) NaCl (aq) + H2O (l)

where 1 mole HCl reacts with 1 mole NaOH to give products

In acid-base reactions, instead of using moles to determine the stoichiometric ratio between the reactants, we can use the term equivalents. An acid equivalent is equal to one mole of H+ or H3O+ ions. Similarly, a base equivalent is equal to one mole of OH- ions. Applying the term equivalents to the above reaction,

HCl (aq) + NaOH (aq) NaCl (aq) + H2O (l)

1 equivalent acid

reacts with

1 equivalent base to give products

In the above example, whether we are using moles or equivalents, at the point of complete neutralization when all the acid is consumed by the base, the ratio of acid and base is 1 mole of acid :1 mole of base or 1 equivalent of acid:1 equivalent of base. Let's take a look at another example where the acid-base ratio is not 1:1.

Example 2: H2SO4 (aq) + 2 NaOH (aq) Na2SO4 (aq) + H2O (l)

where 1 mole H2SO4

reacts with 2 moles NaOH to give products

The stoichiometric mole ratio between the acid, H2SO4, and the base, NaOH, is 1 mole of acid :2 moles of base. When we use equivalents to describe the ratio between the acid and base in the balanced reaction,

• the acid equivalent of H2SO4 is equal to 2 (i.e. each mole of H2SO4 supplies 2 moles of H+ or H3O+ ions),

• the base equivalent of NaOH is also equal to 2 (i.e. each mole of NaOH supplies 1 mole of OH- and there are 2 moles of NaOH in the reaction).

When complete neutralization takes place when the ratio of acid to base is 2 equivalents of acid : 2 equivalents of base.

H2SO4 (aq) + 2 NaOH (aq) Na2SO4 (aq) + H2O (l)

2 equivalents acid

reacts with

2 equivalents base to give products

Last updated 06/01/2008 12:45 AM

The point is, if we think in terms of equivalents instead of moles, at the point of complete

59

neutralization as predicted from the balanced chemical reaction, the number of equivalents of acid is ALWAYS EQUAL to the number of equivalents of base. This is true regardless of whether the stoichiometric mole ratio of acid-base neutralization reaction is 1:1 or 1:2 as seen in the above examples. In summary, at the point of complete neutralization as predicted from the balanced chemical reaction,

the number of equivalents of acid = the number of equivalents of base

This concept will be further explored when we do titration calculations. 2. Concentration: Molarity (M) versus Normality (N) Normality (N) is another way a chemist describes a solution's concentration. Normality is defined as the number of equivalents of solute per liter of solution:

normality = number of equivalents / 1 L of solution

There is a very simple relationship between normality and molarity:

N = n × M (4-2)

where n is an integer, the number of equivalents/mole of solute.

For an acid solution, n is the number of H+ provided by a formula unit of acid.

• For example: 3 M H2SO4 solution is the same as a 6 N H2SO4 solution because H2SO4 provides 2 equivalents of H+ per formula unit of acid. Therefore, the concentration of a 3 M H2SO4 solution is a 6 N H2SO4 solution (2 equivalents x 3 M).

• For example: 3 M HCl solution is a 3 N HCl solution because HCl provides 1 equivalents of H+ per formula unit of acid. Therefore, the concentration of a HCl solution expressed in molarity or normality has the same numerical value, 3 M HCl or 3 N HCl (1 equivalent x 3 M)..

For a basic solution, n is the number of OH- provided by a formula unit of base.

• For example: 1 M Ca(OH)2 solution is the same as a 2N Ca(OH)2 solution because Ca(OH)2 provides 2 equivalents of OH- per formula unit of base. Therefore, the concentration of a 1 M Ca(OH)2 solution is a 2 N Ca(OH)2 solution (2 equivalents x 1 M)..

• For example: A 3 M LiOH solution is the same as a 3N LiOH solution because LiOH provides 1 equivalents of OH- per formula unit of base. Therefore, the concentration of a LiOH solution expressed in molarity or normality has the same numerical value, 3 M LiOH or 3 N LiOH (1 equivalent x 3 M)..

Bottom line! The normality of a solution is NEVER less than the molarity. It can be the same, but never less. (HINT: to remember this, think of the position of the letter 'N' in the alphabet, it's further down. Therefore, N can never be smaller than M.)

Last updated 06/01/2008 12:45 AM 60

http://nobel.scas.bcit.ca/chem0012/labs/4-intro_pH/4-introduction_partB.htm#calc

3. Use of a pH Indicator to Detect the End of a Titration When an acid solution of known concentration, the standard solution, and a base solution of unknown concentration are reacted to the point where the number of acid equivalents equals the number of base equivalents (or vice versa), the equivalence point is reached. This means that at the equivalence point, the two solutions have been mixed in the exact proportions according to the balanced chemical equation such that all of the acid and all of the base are consumed. This is the theoretical point as predicted by the balanced chemical equation. In practice, during a titration, the titrant (the solution in the burette) is slowly added to the solution in the Erlenmeyer flask such that enough of it has been added to completely neutralize the other solution in the Erlenmeyer flask. How do we know when “enough” titrant from the burette has been added to the solution in the Erlenmeyer flask? We look for a colour change in the solution in the Erlenmeyer flask. Prior to the start of a titration, a few drops of a suitable pH indicator (Table 4.1) is added to the solution in the Erlenmeyer flask. A common indicator that is used in acid-base titration is called phenolphthalein. Phenolphthalein is an acid-base indicator which is colourless in acid solution and turns pink when the solution becomes slightly basic. When a colour change in the solution in the Erlenmeyer flask is observed, this is known as the end point of the titration is reached. The end point of the titration indicates that we have added slightly more titrant than necessary to neutralize the acid (or base) in the Erlenmeyer flask. This slightly more titrant that is added is necessary to bring about an observable colour change in the solution in the Erlenmeyer flask. This is a built-in error in using titration as a technique to determine solution concentration. The important thing is to minimize the titration error by being sensitive to observe the colour change and stopping the addition of the titrant once the colour change is observed.

Click on this link to test your understanding of acid-base neutralization


4. Titration Calculations Expressing the concentrations of the acid and the base in normality has its advantage in titration calculations. The fundamental titration equation is

Nacid · Vacid = Nbase · Vbase (4-3)

where Nacid= the normality of the acid (equivalence/L),

Vacid= the volume of the acid (L),

Nbase= the normality of the base (equivalence/L),

Vbase= the volume of the base (L),

This equation is derived from the idea that at the equivalence point, the number of equivalents of the acid is the same as the number of equivalents of the base. In a titration, usually, three of the four variables are known.

Last updated 06/01/2008 12:45 AM 61

http://nobel.scas.bcit.ca/chem0012/labs/4-intro_pH/4-introduction_partA.htm#table

http://highered.mcgraw-hill.com/olcweb/cgi/pluginpop.cgi?it=swf::525::530::/sites/dl/free/0072857684/322633/09_neut_reactions.swf::Neutralization%20Reactions



http://www.adobe.com/shockwave/download/download.cgi?P1_Prod_Version=ShockwaveFlash

http://nobel.scas.bcit.ca/chem0012/labs/4-intro_pH/4-introduction_partB.htm#equiv

http://nobel.scas.bcit.ca/chem0012/labs/4-intro_pH/4-introduction_partB.htm#equiv


A typical set up would be using a known concentration of an acid solution to determine the concentration of a base solution. In the set up of the titration, a known volume of the acid (example: Vacid = 25.00 mL) of a known concentration (example: Nacid = 0.1023 N) is pipetted into an Erlenmeyer flask. The base solution is prepared in a burette and is used as the titrant.To complete the set up, a few drops of colour indicator is added to the Erlenmeyer flask. During the titration, the base solution is slowly added to the acid. When the endpoint of the titration is reached, a colour change in the acid solution is observed. Vbase is the volumet of base that is used to bring the reaction to the endpoint. Vbase is the volume read on the burette (example: Vbase = 24.35 mL). Substitute the numbers in the above example into equation (4-3), the normality of the base solution can be calculated.

(0.02500 L) (0.1023 N) = Nbase (0.02435 L)

Nbase= 0.1050 N

The same titration set up can also be used for determining the normality of the acid if the normality of the base is known.

62


Apparatus:

1. 6 - large test tubes – 1 per pair of students 2. test tube rack – 1 per pair of students 3. 6 - 20 mL beakers – 1 per pair of students 4. pH paper

Solutions:

1. Part A: 5 mL of the following test solutions: o 0.1 M HCl o 0.1 M CH3COOH o 0.1 M NaOH o 0.1 M NH4OH

2. Part A: Indicators o alizarin yellow o bromothymol blue o cresol red o methyl orange o methyl red o phenolphthalein o phenol red o thymol blue o thymolphthalein

3. Part B: o 0.1 M NaOH - 150 mL per student o 0.15 M HCl - 100 mL per student o Phenolpthalein indicator o 10 mL volumetric pipettes o pipet bulb and rack o 50 mL buret o buret reader

63

Procedure:

Part A - pH approximations with pH indicators

1. Label and fill a series of clean test tubes with the solutions indicated on the datasheet. use a clean glass rod and apply a drop of the sample to the broad-range pH paper on a watch glass. Record your results on the datasheet.

2. Working on a sample at a time, divide the sample (approximately 0.5 mL) into a series of small test tubes. Knowing the approximate pH of the sample, apply a few drops of the appropriate indicator to narrow down the pH range of the sample. record results on the datasheet.

Part B - Acid-Base Titration

Download the volumetric measurement techniques and review the techniques before proceeding.

1. Obtain approximately 100 mL of a standard base solution (NaOH) in a clean and dry 250 mL beaker and cover it with a watch glass. Label the beaker "NaOH - base".

2. Empty the burette which is usually stored in distilled water and acclimatize the buret with small amounts of NaOH. Fill the buret. Check the stopcock for leaks. Remove any air bubbles near the stopcock. Lower the level of the solution below the zero mark.

3. Obtain approximately 70 mL of acid solution (HCl) in a clean and dry 100 mL beaker and cover it with a watch glass. Label the beaker "HCl - acid". Pour a small amount of acid into a clean and dry 50 mL beaker and use this to acclimatize a 10.00 mL pipette.

4. Rinse three 125 mL Erlenmeyer flasks with distilled water and discard the washings.

5. Pipette 10.00 mL portions of acid solution into each flask. Rinse the inner walls with 15-20 mL of distilled water with a wash bottle.

6. Add 3 drops of phenolpthalein indicator to each flask. 7. Record the initial burette reading. Put a piece of white

paper under the Erlenmeyer flask to observe the colour change easily. Your titration set up should resemble Figure 1.

8. Titrate with the NaOH solution while swirling the Erlenmeyer flask. Rinse the inner walls of the Erlenmeyer flask occasionally.

9. Titrate until the acid solution in the Erlenmeyer flask changes to a pink colour permanently (see Figure 2). The colour should last at least 40 seconds.

10. Record the final burette reading. 11. Repeat with the other two acid samples. Repeat your

titration until you have at least two volumes that are within +/- 0.05 mL.

Figure 1: Acid-base titration

Last updated 06/01/2008 12:45 AM 64

http://nobel.scas.bcit.ca/chem0012/labs/1-safety/0012-techniques.pdf

Figure 2: Flask #1 shows a good endpoint where the colour of the solution changes to a pale pink colour. Flask #2 shows a bright pink endpoint. The brighter colour indicates that the endpoint of this titration for flask #2 is farther away from the equivalence point than flask #1.

Last updated 06/01/2008 12:45 AM 65


Datasheet:

Part A - pH approximations with pH indicators

Sample Distilled water 0.10 M HCl 0.1 M CH3COOH 0.10 M NaOH 0.1 M NH4OH

pH paper test - approximate pH range and colour

1st indicator: Colour and approximate pH range

2nd indicator: Colour and approximate pH range

3rd indicator: Colour and approximate pH range

Deduced pH range

66


Datasheet:

Part B - Acid-Base Titration Solutions:

Solution Concentration

Acid: ___________________ _________________________ (copy from the bottle)

Base: ____________________ to be determined

Trial 1 2 3 4 5

Final buret reading (mL)

Initial buret reading (mL)

Volume of base added (mL)

67

Postlab Questions:

Part A - pH approximations with pH indicators List the solutions tested in order of lowest pH to highest pH. Explain the order observed.

Part B - Acid-Base Titration

1. Write the balanced chemical reaction for the titration:

Using data from trials : __________________________________ Average volume of base: ________________________________

2. Calculation of the normality of NaOH:

Last updated 06/01/2008 12:45 AM 68


Lab 5: pH Measurements and Acid-Base pH Titration Curves

Objectives:

1. to standardize a pH meter and then use it to determine the pH of various solutions. 2. to determine pH titration curves for a strong acid/strong base and a strong acid/weak

base titration.

Introduction:

pH Measurements

In the last lab, we learned that we can approximate the pH of solutions with the use of broad-range pH paper. pH paper can be used for rough estimation of pH. They are made of paper impregnated with pH indicators (see Lab 4 - Table 4.1). pH papers are useful when 0.5 to 1 pH unit accuracy is needed. When high precision is required, pH measurements determined using potentiometric method is used. The most widely used method for measuring pH is the combination pH electrode and a pH meter. The combination pH electrode, as the name suggests, contains two different electrodes in one probe. One is a reference electrode that does not change in voltage. The second is a glass membrane sensitive to pH. It is the voltage difference between these two electrodes that the meter reads and converts into pH.

Standardizing and using a pH Meter

Each pH meter will have a slightly different way of standardizing. But there are a number of general rules that are useful to note.

1. The standardization of a pH meter must be calibrated with standard buffer solutions that span the range expected. Usually the pH meter is standardized using two buffers of known pH. The electrode is first immersed in the pH 7 buffer. If the pH meter is used to measure in the acidic range, then the second buffer used in the standardization is the pH 4 buffer. If the pH meter is used to measure in the basic range, then the second buffer used in the standardization is the pH 10 buffer.

2. Ensure that the buffer solutions used in the standardization of the pH meter are accurate pH buffers. Buffer solutions that have been poured into beakers can usually be kept for the entire lab period. It should be discarded at the end of the lab period.

3. The pH electrode should be rinsed in distilled water and blotted dried before it is immersed in a new solution.

4. It takes time for a pH meter to obtain a stable reading. Let the meter equilibrate to each buffer. Some meters will beep or display on screen to indicate that the meter is reading a stabilized pH reading.

5. Stirring the solution will help the pH electrode to equilibrate to the solution. 6. Ensure that the room temperature stays constant and the solution temperature is near

room temperature. pH is a function of temperature. The pH of the buffers will change at different temperatures.

69


pH Titration Curves

In the previous lab, a pH indicator was added to one of the solutions and the titration was carried out until a permanent colour change was observed. The progress of an acid-base titration is often monitored by plotting the pH of the solution against the volume of added titrant to give a pH curve. A pH curve will reflect the changes in pH that occur as the titrant from a burette is added to the solution in the Erlenmeyer flask. The shape of the pH titration curve will depend on the types of acid and base used: strong acid, weak acid, strong base or weak base. In this experiment, we will study two pH titration curves for:

1. strong acid-strong base titration 2. strong acid-weak base titration

For the neutralization reactions where the reactions are between a strong acid and a strong base, the equivalence point will occur at pH 7. For reactions involving weak acids and bases, the equivalence point will not occur at pH 7. We will study each pH titration curve, identify the equivalence point and explain the behaviour of the observed pH shift of the equivalence point.

Click on this link to view the pH titration curve for

an acid-base neutralization reaction


70


Apparatus:

1. large test tubes – 6 per pair of students 2. test tube rack – 1 per pair of students 3. 20 mL beakers – 6 per pair of students 4. 200 mL tall form beaker 5. pH combination electrode – 1 per pair of students 6. pH meter – 1 per pair of students

Solutions:

1. pH 4 and pH 7 buffers 2. Part A: 0.1 M unknown solutions - 50 mL per pair of students 3. Part B:

o 0.1 M NaOH - 150 mL per student o 0.15 M HCl - 100 mL per student o 0.15 M NH4OH - 100 mL per student o 10 mL volumetric pipettes - 2 per student o pipette bulb and rack o 50 mL burette o burette reader

71

Last updated 12/13/2010 5:33:00 PM 72

Procedure:

Part A - pH Measurements with a pH meter

1. Calibrate the pH meter as demonstrated by your lab instructor. The buffer solutions are reusable for the entire lab period.

2. Obtain about 10 mL of 0.1 M "unknown" solutions in the large test tubes. Measure the pH of the solution using the pH meter. Rinse the electrode in between measurements. Record the results.

3. Recalibrate the pH meter. Obtain about 10 mL of the natural samples provided in the large test tubes. Measure the pH of some natural samples.

Part B - Acid-Base pH Titration

NOTE: In this lab, fill the buret with HCl solution

Record the pH titration of a Strong Acid (HCl) /Strong Base (NaOH) Titration:

1. Obtain approximately 100 mL of a standard acid solution (HCl) in a clean and dry 250 mL beaker and cover it with a watch glass. Label the beaker "HCl - acid".

2. Empty the burette, which is usually stored in distilled water and acclimatize the buret with small amounts of HCl solution. Fill the burette. Check the stopcock for leaks. Remove any air bubbles near the stopcock. Lower the level of the solution below the zero mark.

3. Obtain approximately 70 mL of base solution (NaOH) in a clean and dry 100 mL beaker and cover it with a watch glass. Label the beaker "NaOH - base". Pour a small amount of base solution into a clean and dry 50 mL beaker and use this to acclimatize a 10.00 mL pipette.

4. Rinse a 125 mL Erlenmeyer flask with distilled water and discard the washings. 5. Pipet 10.00 mL of the NaOH solution into the Erlenmeyer flask. Rinse the inner walls with

15-20 mL of distilled water with a wash bottle. 6. Stir the solution to ensure uniform mixing. 7. Dip the pH combination electrode into the NaOH solution and clamp in place. Your pH

meter should display a high pH reading (i.e. pH > 7). 8. Slowly add the HCl solution at 2.00 mL aliquot and record the pH of the solution after

each addition. 9. When the pH of the solution is approximately 11, slow down the addition of the HCl to

adding drop wise. Record the pH of the solution after each addition. 10. When the pH of the solution is approximately 4, increase the addition of HCl to 2.00 mL

aliquot. Continue recording the pH after each addition. When the pH remains fairly constant, add two 2.00 mL aliquots and record the pH readings.

Record the pH titration of a Strong Acid (HCl) /Weak Base (NH4OH) Titration:

1. Refill the burette with HCl solution. 2. Obtain approximately 70 mL of base solution (NH4OH) in a clean and dry 100 mL beaker

and cover it with a watch glass. Label the beaker "NH4OH - base". Pour a small amount of base solution into a clean and dry 50 mL beaker and use this to acclimatize a 10.00 mL pipette.

3. Rinse a 125 mL Erlenmeyer flask with distilled water and discard the washings. 4. Pipet 10.00 mL of the NH4OH solution into the Erlenmeyer flask. Rinse the inner walls

with 15-20 mL of distilled water with a wash bottle. 5. Stir the solution to ensure uniform mixing. 6. Dip the pH combination electrode into the NH4OH solution and clamp in place. Your pH

meter should display a high pH reading (i.e. pH > 7). 7. Slowly add the HCl solution at 2.00 mL aliquot and record the pH of the solution after

each addition. 8. When the pH of the solution is approximately 8, slow down the addition of the HCl to

adding drop wise. Record the pH of the solution after each addition. 9. When the pH of the solution is approximately 3, increase the addition of HCl to 2.00 mL

aliquot. Continue recording the pH after each addition. When the pH remains fairly constant, add two 2.00 mL aliquots and record the pH readings.


Datasheet:

Part A - pH Measurements with a pH meter

pH Measurements of 0.1 M unknown samples

Sample A B C D E F

Measured pH

Acidic, Basic or Neutral?

pH Measurements of Natural Samples

Sample Vinegar Battery Acid Baking Soda Washing Soda

Measured pH

Acidic, Basic or Neutral?

73


Datasheet:

Part B - Acid-Base pH Titrations

Strong Acid/Strong Base Titration Strong Acid/Strong Base Titration Volume of Acid (mL) pH Volume of Acid (mL) pH

74


Datasheet:

Part B - Acid-Base pH Titrations

Strong Acid/Weak Base Titration Strong Acid/Weak Base Titration Volume of Acid (mL) pH Volume of Acid (mL) pH

75


Postlab Questions:

Part A - pH Measurements with a pH meter 1. From the unknown samples in Part A, which sample is most likely HCl? Explain your choice. 2. From the unknown samples in Part A, which sample is most likely CH3COOH? Explain your choice. 3. From the unknown samples in Part A, rank the samples from strongest to weakest acidity.

Part B - Acid-Base pH Titration

4. Write the balanced chemical reaction for the strong acid/strong base titration.

5. Write the balanced chemical reaction for the strong acid/weak base titration. 6. Graph the pH titration curves on two separate graph papers. Your graph should fill the size of the piece of graph paper. The y-axis is pH. The x-axis is volume of HCl added. Each graph should contain:

• Title of graph • A balanced chemical reaction • Labelled axes, include units where applicable • A smooth curve drawn through all the points plotted • The equivalence point • The pH at the equivalence point

7. Explain why the pH of the equivalence points for the two titrations is different.

76

Last updated 12/13/2010 6:33:00 PM 77

Lab 6: Acid-Base Equilibria: Determination of Ka and Investigation of Buffers

Objectives:

To study the buffer capacity of a buffer.

Introduction: For a general weak acid of the form HA that is dissolved in water, the following equilibrium is set up.

HA (aq) H+(aq) + A- (6-1)

where HA (aq) is the weak acid, H+(aq) is the hydrogen ion and A- is the conjugate base. The equilibrium constant Ka, for this reaction is called the acid dissociation constant of the acid HA.

(6-2)

At the half equivalence point, half the acid has been converted to the salt. Therefore, [A-] = [HA] and equation (6-2) becomes

[H+] = Ka (6-3)

or

pH = pKa. (6-4)

The Henderson-Hasselbalch equation is used for the calculation of the pH or composition of a buffer solution. With mixtures consisting of weak acids and their salts the convenient approximations [HA] = total acid concentration, and [A- ] = salt concentration, can often be made. Hence an acetic acid/acetate buffer solution containing 0.1 M acetic acid and 0.05 M sodium acetate would have a pH of 4.4. pH = pKa + log [conjugate base] = 4.7 + log (0.05) = 4.7 + (-0.3) = 4.4 (6-5) [conjugate acid] (0.1)

Buffer solutions behave as follows: when a strong acid is added, the H+ from that acid combine with a portion of the anion to form undissociated acid, thereby removing most of the added H+ from the solution

Last updated 12/13/2010 6:33:00 PM 78

H+ (aq) + A- (aq) HA (aq)

When strong base is added part of the undissociated acid reacts to form anions

OH- (aq) + HA (aq) A- (aq) + HOH (l)

In Part A, we saw that at the half equivalence point, where [A-] = [HA] and equation (6-2) becomes

[H+] = Ka (6-3)

Using the Henderson-Hasselbalch equation (6-5)

(6-6)

When [A-] = [HA], equation (6-6) reduces to

pH = pKa + log 1 (6-7) and

pH = pKa. (6-8)

This means that an acid is half dissociated when the pH of the solution is numerically equal to the pKa of the acid.

Click on this link to view a buffer solution


Last updated 12/13/2010 6:33:00 PM 79

Apparatus:

1. Burettes – 2x per group of students 2. 200 mL tall-form beakers – 2x per group of students 3. Stir bar 4. Stir plate 5. pH meter – 1 per pair of students

Solutions:

1. pH 4 and pH 7 buffers 2. 0.10 M NaOH - 80 mL per pair of students 3. 0.10 M HCl - 80 mL per pair of students 4. 0.100 M acetic acid - 60 mL per pair of students 5. 0.100 M sodium acetate - 60 mL per pair of students 6. Phenolpth alein indicator 7. Metacresol purple or bromocresol green indicator

Last updated 12/13/2010 6:33:00 PM 80

Procedure:

1. Calibrate the pH meter as demonstrated. The following is the calibration procedure for some pH meters.

a. Immerse pH electrode into the pH 7 buffer solution. b. Adjust the Standardized knob to display pH=7.00. c. Rinse the electrode with distilled water and blot it dry. d. Immerse the pH electrode into the pH 4 buffer solution. e. Adjust the Temperature knob to display pH=4.00.

[Note: For the newer digital pH meter and electrode units, remember to open the vent in the electrode before using the unit.]

Recipe for making a buffer solution:

Buffer Solution 0.100 M acetic acid (Dispenser)

0.100 M sodium acetate (Dispenser)

A 10.00 mL 50.00 mL B 20.00 mL 40.00 mL C 30.00 mL 30.00 mL D 40.00 mL 20.00 mL E 50.00 mL 10.00 mL

2. Your instructor will assign a buffer solution ID to your group. Using the above recipe, prepare 60.00 mL of the assigned buffer. The solutions are in bottle-top dispensers that are set to 10.0 mL. Record the pH of the prepared buffer. Cover the beaker with a watch glass.

3. Record the exact concentrations of the NaOH and HCl solutions. (They are approximately 0.10 M.) Acclimatize and fill a burette with NaOH and another burette with HCl. Label the burettes. Do not take more than 60 mL of either reagent to start.

4. Use a 25.0 mL graduated cylinder to measure 20.0 mL of your buffer into two clean and dry 200 mL tall-form beakers.

5. Put a magnetic stirrer in one and immerse the pH electrode into the buffer. Check that the probe is immersed and not touching the spinning stir bar. Add three drops of phenolphthalein indicator. Record the initial pH of the solution.

6. Titrate with 1 mL of NaOH and record the pH and volumes after waiting about 20 seconds after each addition. Continue until pH 6.

7. Continue by adding smaller volumes of 0.1 mL NaOH until the basic equivalence point is reached (it will be close to the colour change of the indicator). Record the pH and volumes after each addition after mixing for 20 seconds. Aim to record a pH and volume reader for every 0.2 pH change.

8. Continue to add a few millilitres of NaOH beyond the equivalence point and record the volume and pH. You can stop when pH > 12.

9. Using the other buffer sample prepared in step 4, repeat the experiment by using 0.10 M HCl and either metacresol purple or bromocresol green as an indicator. The colour change of the indicator may not coincide with the actual equivalence point but it can be used as a signal as the equivalence point is reached. You can stop when pH < 2.

10. Your instructor will instruct you on submission of your data and how to download the class data for all the other buffer solutions to carry out data analysis.

Last updated 12/13/2010 6:33:00 PM 81

Datasheet:

Assigned Buffer Solution: __________________________________________

Initial pH of the buffer solution ___________________________________________

Last updated 12/13/2010 6:33:00 PM 82

82Titration with NaOH

Buffer/Base Titration Buffer/ Base Titration Volume of NaOH (mL) pH Volume of NaOH (mL) pH

Last updated 12/13/2010 6:33:00 PM 83

Titration with HCl

Buffer/Acid Titration Buffer/Acid Titration Volume of HCl (mL) pH Volume of HCl (mL) pH

Last updated 12/13/2010 6:33:00 PM 84

Postlab Questions:

1. Graph the pH titration data for your buffer solution. Combine the data from the NaOH and HCl titrations on the same graph. The y-axis is pH. The x-axis is volume of titrant (NaOH and HCl) added. Use 1 mL as the large scale on the x-axis. Place the y-axis with an increasing pH scale near the middle of the graph paper. Plot the titration volume of HCl on the left side of the y-axis with increasing volume going right to left. Plot the titration volume of NaOH on the right side of the y-axis with increasing volume going from left to right. Draw a smooth curve through the data points and label the following:

initial pH of the buffer solution the basic equivalence point the acidic equivalence point the pKa

Below is an example of the pH titration curve of a acetic acid/sodium acetate buffer solution:

2. Your instructor will provide you with instructions on how to name your graph. (Example: Buffer X from John Doe_Jane Doe.xls) 3. Your instructor will provide you with instructions to download the class data from the course website. 4. If you can, on the same graph, graph the titration data for the different buffer solutions (A to E). Use different colours to distinguish the different buffer solutions. Create a legend to identify coloured curves to the different buffer solutions. When done properly, you will be able to view the shift of the titration curves from Buffer A to E. Otherwise, you will need to inspect each graph separately in order to answer the questions below. 5. The buffer capacity of a buffer solution is the amount of acid or base that can be added to a volume of a buffer solution before its pH changes significantly. Estimate the buffer capacity of the different buffer solutions for: (a) NaOH, and (b) HCl. Explain. 6. From the class data, which buffer solution has the greatest buffer capacity for (a) NaOH, and (b) HCl. Explain.

Last updated 12/13/2010 6:57:00 PM 85

Lab 7: Determination of an Equilibrium Constant using Spectroscopy

Objectives: To determine the concentration of iron (III) thiocyanate ions, FeSCN2+, in various iron (III) nitrate, Fe(NO3)3, and potassium thiocyanate, KSCN. The results of these measurements will determine the equilibrium constant for the formation of FeSCN2+.

Introduction: Chemical reactions occur to reach a state of chemical equilibrium. The equilibrium state can be characterized by specifying its equilibrium constant. In this experiment you will determine the value of the equilibrium constant for the reaction between the iron (III) ion and thiocynate ion (SCN-).

Fe3+ (aq) + SCN- (aq) FeSCN2+ (aq) (7-1)

where

(7-2)

To find the value of Keq, it is necessary to determine the concentration of each of the three species in solution at equilibrium. This can be done using UV-visible spectroscopy.

Part A - Determination the Calibration Curve of FeSCN2+ by Spectrometry

When a chemical reaction reaches chemical equilibrium, the rates of the forward and the reverse reactions are equal. The concentrations of all the species become constant. In this experiment, we take advantage of the fact that FeSCN2+ is a coloured compound and that the concentration of this compound can be determined by measuring its absorbance using spectrophotometric methods. This method requires a calibration curve using samples of known concentration.

Beer's law relates absorbance, A, the light absorbed on passing through a length of solution, l, of a concentration, c, of the absorbing solute, c. According to Beer's law, the amount of light absorbed by a coloured species in solution at a specific wavelength is directly proportional to the concentration of the coloured species.

A = ε · l · c (7-3)

The absorbance is measured for a series of standard solutions with known concentrations. If a solute obeys Beer's law, according to equation (7-3), a graph of the "Absorbance versus Concentration" will yield a straight line of slope equal ε · l. This is known as a calibration curve for

Last updated 12/13/2010 6:57:00 PM 86

the solute. From the calibration curve, the absorbance of the solute with an unknown concentration can be determined. In usual practice of this technique, a blank is used to cancel out absorption by any other solutes, so that the measured absorbance is directly proportional to the concentration of the solute of interest. This is routinely done by people working in the fields of medicine, forensic science, and chemistry.

In this lab, the solute of interest is FeSCN2+. A series of standard FeSCN2+ solutions will be prepared from solutions of varying concentrations of SCN- and constant, high concentrations of Fe3+ and H+. The high concentration of the Fe3+ ion, relative to that of the SCN- ion, drives reaction (7-1) to the right; essentially to completion. Solutions of Fe3+ are weakly coloured and the SCN- ion is colourless. The primary absorber in the mixture will be the FeSCN2+. FeSCN2+ is known to obey Beer's law over a wide range of concentrations with maximum absorption at a single wavelength of 447 nm. Thus, we will use this setting to make measurements of its concentration in equilibrium mixtures. The high [H+] will ensure that the iron (III) ion will not form the insoluble compound, iron (III) hydroxide. The high [Fe3+] will ensure that all SCN- reacts to form FeSCN2+. The FeSCN2+ complex forms slowly in about 1 minute and then it decomposes slowly due to its reaction with light. For best results, the absorbance value should be read between 2 and 4 minutes after preparation and all samples read after the same time interval.

Part B - Determination of an Equilibrium Constant

A second series of solutions will be prepared for the determination of the equilibrium constant. These solutions will have a fixed high concentration of H+, a constant low concentration of Fe3+

and varying low concentration of SCN-. The absorbance readings will enable the equilibrium concentration of FeSCN2+ to be determined directly. Using the equilibrium concentration of FeSCN2+, the equilibrium concentrations of the reactants can be calculated using their initial concentrations.

Last updated 12/13/2010 6:57:00 PM 87

Apparatus:

1. spectrophotometer - 1 per pair students 2. cuvettes - 4 matching cuvettes per pair of students 3. 25.00 mL volumetric flask - 8 per pair of students 4. 10.00 mL burette - 1 per pair of students

Solutions:

1. 0.10 M HNO3 2. 0.150 M Fe(NO3)3 in 0.10 M HNO3 – in dispenser 3. 1.50 x 10-3 M Fe(NO3)3 in 0.10 M HNO3 – in dispenser 4. 3.00 x 10-3 M SCN- in 0.10 M HNO3 5. 5.00 x 10-4 M SCN- in 0.10 M HNO3

Last updated 12/13/2010 6:57:00 PM 88

Procedure:

Part A - Determinination the Calibration Curve of FeSCN2+ by Spectrometry Note: For Part A of the experiment, we will be using the 0.150 M Fe(NO3)3 solution. We can assume that all of the SCN- is present as the complex ion, FeSCN2+ due to the high concentration of Fe3+.

Work in pairs. Deliver all volumes of Fe3+ solutions with a 10.00 mL volumetric pipette and SCN-

solutions using a burette. Add the solutions into a 25.00 mL volumetric flask and dilute to the mark with 0.10 M

nitric acid. All cuvettes should be wiped clean and dry on the outside with a tissue. All solutions should be free of bubbles. Always position the cuvette with its reference mark facing the reference mark on the

spectrophotometer.

1. Turn on the spectrophotometer, set Mode to Absorbance, and set the wavelength to 447 nm.

2. Obtain a cuvette and rinse (2x with small portions) of distilled water and fill the cuvette ¾ full. Wipe the outside of the cuvette and avoid touching the bottom half of the cuvette.

3. Insert the cuvette into the sample compartment with its orientation mark aligned with the mark in the spectrophotometer.

4. Prepare Solutions #1 according to Table 7-1 using a 25.00 mL volumetric flasks. Dilute to the mark with 0.10 M HNO3 and mix well.

Solution # 1 (blank)

2 3 4 5 6

Volume of 0.150 M Fe(NO3)3 in 0.10 M HNO3 – from the bottle-topdispenser

10.00 10.00 10.00 10.00 10.00 10.00

Volume of 5.00 x 10-4 M SCN- in 0.10 M HNO3 – use a burette

0.00 2.00 3.00 5.00 7.00 9.00

Table 7-1: Volume of reagents used to make solutions.

5. Obtain a clean cuvette and rinse (2 x with small portions) with Solution #1, the blank solution, and fill the cuvette ¾ full. Insert the cuvette into the sample compartment with its orientation mark aligned with the mark in the spectrophotometer and measure the absorbance. This is the absorbance of unreacted Fe3+ in the solution. All the standards have a large excess of Fe3+, which absorbs light in this wavelength region. This value must be subtracted from the absorbance of the standard solutions.

6. Repeat Step 4 and 5 to prepare and measure the absorbance for Solutions #2 to #6 in order of increasing SCN- concentration and read the absorbance of each sample once the solution is prepared. Remember to rinse the cuvette with each sample and use water as the reference.

7. Insert the cuvette filled Solution #2 and read the absorbance of the solution. 8. Check that the absorbance values are in direct proportion with one another so the graph

will be linear. Prepare alternate solutions to any which fall outside the linear region and measure the absorbance.

Last updated 12/13/2010 6:57:00 PM 89

Procedure:

Part B - Determination of an Equilibrium Constant Note: For Part B of the experiment, we will be using the 1.50 x 10-3 M Fe(NO3)3 solution.

Work in pairs. Deliver all volumes of Fe3+ solutions with a 10.00 mL volumetric pipette and SCN-

solutions using a burette. Add the solutions into a 25.00 mL volumetric flask and dilute to the mark with 0.10 M

nitric acid. All cuvettes should be wiped clean and dry on the outside with a tissue. All solutions should be free of bubbles. Always position the cuvette with its reference mark facing the reference mark on the

spectrophotometer.

1. Prepare the solutions shown in Table 7-2 and measure the absorbance of each using water as a reference to zero the instrument and Solution #7 as the new blank. As in Part A, measure the absorbance of the solution once it has been prepared.

2. Prepare alternate solutions to any which have absorbance outside the linear region of Beer’s Law.

Solution # 7 (blank)

8 9 10 11 12

Volume of 1.50 x 10-3 M Fe(NO3)3 in 0.10 M HNO3 – from the bottle-top dispenser

10.00 10.00 10.00 10.00 10.00 10.00

Volume of 3.00 x 10-3 M SCN- in 0.10 M HNO3 - use a burette

0.00 2.00 3.00 5.00 7.00 8.00

Table 7-2: Volume of reagents used to make solutions.

Last updated 12/13/2010 6:57:00 PM 90

Datasheet:


Wavelength used for measurements: ______________________________________

Solution # [FeSCN2+] Absorbance

1 zero (blank)

2

3

4

5

6

Last updated 12/13/2010 6:57:00 PM 91

Datasheet:

Part B - Determination of an Equilibrium Constant

Wavelength used for measurements: ___________________________________________

Solution # [FeSCN2+] Absorbance

7 zero (blank)

8

9

10

11

12

Solution #

initial [Fe3+]

initial [SCN-]

equilibrium [FeSCN2+]

equilibrium [Fe3+]

equilibrium [SCN-]

Keq

8

9

10

11

12

Last updated 12/13/2010 6:57:00 PM 92

Postlab Questions:


1. Calculate [FeSCN2+] for each tube Part A and enter the concentrations on the datasheet. Show a sample calculation for Solution #3 below.

2. Plot a calibration graph of net absorbance versus [FeSCN2+]. Include the (0,0) point. Clearly label the graph. Obtain the equation of the straight line for the calibration graph written in terms of the variables used in this experiment. Attach the graph to the report sheet. Part B - Determination of an Equilibrium Constant 3. Determine the [FeSCN2+] at equilibrium from the calibration graph and measured absorbance. Enter the concentration values on the datasheet. 4. Calculate the equilibrium concentrations of Fe3+ and SCN- for Solutions #8 to #12. Enter the concentration values on the datasheet.

Last updated 12/13/2010 6:57:00 PM 93

35. Calculate the equilibrium constant, Keq, for Solutions #8 to #12. Show sample calculation of Keq for Solution #10 below.

6. Calculate the average Keq.

Average Keq = ______________________


Lab 8: Electrochemistry: The Study of Corrosion in Metals

Objectives:

To study the process of corrosion in metals

Introduction:

Facts:

1. For corrosion to occur, normally it is necessary for oxygen and water to be present. 2. Corrosion usually increases with an increase in strong electrolyte concentration. 3. If dissimilar metals are in contact, the one that will corrode will be the metal that's higher

in the activity series. 4. The use of a sacrificial anode is a deliberate use of the dissimilar metal theory. 5. Corrosion can occur on a single metal - usually at 'stress' points - or at surface

imperfections. 6. There are two main types of corrosion: Galvanic (due to difference in metals); Local (due

to differences in electrolyte concentration). 7. Corrosion processes occur in cells usually with the metal (or metals), electrolyte, and

oxygen present. 8. Corrosion processes are redox processes.

Definitions:

1. Corrosion is oxidation and occurs at the anode. 2. Reduction occurs at the cathode. 3. Corrosion is a combination of an oxidation ½ reaction and a reduction ½ reaction. One

cannot occur without the other.

Predictions:

There are various and sometimes sophisticated methods used to determine whether a specific reaction will occur spontaneously or not. Some reactions that do not occur spontaneously can be forced to using an external source of energy (i.e. electrolysis).

In electrochemistry, the Table of Standard Reduction Potentials is used to predict whether the reaction is spontaneous or not. All the ions are one molar concentration and gases ard one atmosphere pressure, at 25oC. The Table of Standard Reduction Potentials shows the reduction ½ reaction and its potential. If the reaction is reversed, then it becomes an oxidation ½ reaction and the sign of the potential is reversed (see below and find it in the Standard Table).

94


Example:

Oxidation ½ reaction Zn (s) Zn2+ (aq) + 2e- Eo = + 0.76 V

Reduction ½ reaction Cl2 (aq) + 2e- 2 Cl- (aq) Eo = + 1.36 V

Net reaction Zn (s) + Cl2 (aq) Zn2+ (aq) + 2 Cl- (aq) Eo = +2.12 V

The cell potential of the above reaction is +2.12 V and is spontaneous. Any negative cell potential calculated for a cell will not give a spontaneous reaction. If the cell potential is negative then the reaction will not occur to any great extent.

Example: Will a reaction occur if a solution of Cu2+ (aq) is mixed with an aqueous solution of I- (aq) to form I2 (s) and Cu (s)?

Oxidation ½ reaction 2 I- (aq) I2 (s) + 2e- Eo = - 0.53 V

Reduction ½ reaction Cu2+ (aq) + 2e- Cu (s) Eo = + 0.34 V

Net reaction Cu2+ (aq) + 2 I- (aq) I2 (s) + Cu (s) Eo = - 0.19 V

Concentration Cells

A concentration cell can be constructed using the same ½ cells but with different concentrations of the same electrolyte. When a cell that has the following notation is constructed.

Cell Notation

We can convey an electrochemical cell by using the cell notation rather than drawing a picture or an actual diagram of the cell. Cell notation lists the metals and ions involved in the reaction. A vertical line, |, denotes a phase boundary. A double line, ||, represents the salt bridge. The anode is always written on the left and the cathode on the right. The general form is:

anode | electrolyte of anode || electrolyte of cathode | cathode

For example, the cell notation for the concentration cell made with Fe metal dipped into different concentration of Fe+2 solutions can be represented as follows:

Fe | Fe+2 (0.10 M) || Fe+2 (1.0 M) | Fe

The half-cell written to the left of the double line, ||, is the anode.

Fe (s) Fe+2 (0.10 M) + 2 e- (8-1)

The half-cell written to the right of the double line, ||, is the cathode.

Fe+2 (1.0 M) + 2 e- Fe (s) (8-2)

95


Apparatus:

Part B:

1. pH meter - 1 per pair of students 2. 3% NaCl solution 3. steel wool 4. iron nails - 1 per pair of students 5. zinc strips - 1 per pair of students 6. aluminum strip - 1 per pair of students 7. copper strip - 1 per pair of students

Solutions:

1. 0.10 M HNO3 2. 0.150 M Fe(NO3)3 in 0.10 M HNO3 3. 1.50 x 10-3 M Fe(NO3)3 in 0.10 M HNO3 4. 3.00 x 10-3 M SCN- in 0.10 M HNO3 5. 5.00 x 10-4 M SCN- in 0.10 M HNO3

96


Procedure:

Part A - Corrosion and Electrodes Four petri dishes containing an agar-salt solution. Once the agar cools, this is an electrolyte supported in gel matrix. It contains 7.5 g NaCl, 5 g powdered agar, 5 mL of 5% potassium ferricyanide and 1 mL of 1% phenolphthalein solution. The petri dishes will have the following metal or metals inserted in the agar gel.

Dish 1 Dish 2 Dish 3 Dish 4 Fe Fe / Zn Fe / Cu Cu / Fe / Zn

Record your observations of the petri dishes provided. You may observe one or more of the following results:

• Oxidation of Fe to Fe2+. In the agar-salt solution any iron that is oxidized will react with the 5& potassium ferricyanide and form a blue colour described as "Prussian Blue". If iron is not oxidized, there is no blue colour in the agar gel.

• Oxidation of Zn to Zn2+. In the agar-salt solution any zinc metal that is oxidized will give a white-grey colour.

• Oxidation of O2 to OH-. The agar-salt solution contains phenophthalein, an acid-base indicator that is pink in basic conditions. If oxygen is reduced then the hydroxide that is formed will cause the agar to turn pink.

Part B - Activities of Metals

1. On a piece of paper towel, polish the metal electrode strips with some steel wool. Wipe the electrodes clean with a damp paper towel.

2. Place two iron nails in a beaker of electrolyte made of 3% NaCl. Keep the nails separated but connect one to the reference lead and the other to the input lead of the potentiometer. Zero the instrument.

3. Replace the iron on the input lead with an aluminum electrode. Read the millivolt reading (take the highest reading).

4. Replace the aluminum on the input lead with a zinc electrode. Read the millivolt reading (take the highest reading).

5. Replace the zinc on the input lead with a copper electrode. Read teh millivolt reading (take the highest reading).

Part C - Concentration Cells - Demonstration

1. Construct a galvanic cell that has two different concentrations of Cu2+. Use a salt bridge to connect the two ½ cells and a copper electrode for each ½ cell.

2. Cell 1: The first galvanic cell has a beaker of 1.0 M Cu2+ versus a beaker of 0.10 M Cu2+. Measure the cell potential.

3. Cell 2: The second galvanic cell has a beaker of 1.0 M Cu2+ versus a beaker of 0.001 M Cu2+. Measure the cell potential.

Part D - Corrosion - Demonstration

1. Record your observations for a nail in water saturated with oxygen gas. 2. Record your observations for a nail in water saturated with nitrogen gas.

97


Datasheet:

Part A - Corrosion and Electrodes

Sketch to record your observations, colours, and any important results of the nails in the petri dishes. Dish 1:

Fe

Dish 2:

Fe / Zn

Dish 3:

Fe / Cu

Dish 4:

Cu / Fe / Zn

Part B - Activities of Metals

Reference Input Reading (mV)

Fe Fe 0.0

Fe Al

Fe Zn

Fe Cu

98


Part C - Concentration Cells – Demonstration

Cell 1: 1 M Cu2+ versus 0.1 M Cu2+ ____________________________________ mV

Cell 2: 1 M Cu2+ versus 0.001 M Cu2+ ____________________________________ mV

Part D - Corrosion - Demonstration Record your observations, colours and any important results below.

Fe nail in O2 saturated water Fe nail in N2 saturated water

99


Postlab Questions:

Part A - Corrosion and Electrodes 1. Answer the following questions by writing the answers in the Table .8-1.

a. What metal serves as the cathode?

b. What metal serves as the anode?

c. What is the oxidation ½ reaction?

d. What is the reduction ½ reaction?

Dish Cathode Anode Oxidation ½ reaction Reduction ½ reaction

1

2

3

4

2. Why is zinc referred to as the sacrificial anode when it is paired with Fe?

100


Part B - Activities of Metals 3. For the potential measurements between Fe & Al, Fe & Zn, and Fe & Cu, identify:

• Which metal is the anode?

• Which metal is the cathode?

Pair Cathode Anode

Fe / Al

Fe / Zn

Fe / Cu

4. Based on your experimental results, list the metals Fe, Al, Zn and Cu in order of strongest reducing agent to the weakest reducing agent. Part C: Concentration Cell 5. a. Write the cell notation for Cell 1. Include the concentrations.

b.

Which solution of the cell undergoes oxidation?

c.

Which solution of the cell undergoes reduction?

101


6. a. Which cell (1 or 2) has the higher potential?

b. What is the trend in cell potential as the concentration of one of the half-cell is decreased?

Part D: Corrosion Demonstration 7. Which system shows more corrosion? Explain your answer.

102


Lab 9: Determination of Solubility Product Constants, Ksp, using Potentiometric Method

Objectives: The concentration of Ag+ from an insoluble silver halide AgX is determined using a silver value.

Introduction:

A concentration cell is an electrochemical cell that has two equivalent half-cells of the same material but differing only in concentration. As a result, the concentration cell produces a voltage. For example, consider a cell where the anode and the cathode are a piece of copper metal, Cu(s). At the anode, the copper metal is dipped into a 0.10 M Cu +2 solution. The following oxidation half-reaction occurs at the anode.

Cu (s) Cu +2 (aq, 0.10 M) + 2 e- (9-1)

At the cathode, the copper metal is dipped into a 1.0 M Cu +2 solution. The following reduction half-reaction occurs at the cathode.

Cu+2 (aq, 1.0 M) + 2 e- Cu (s) (9-2)

The anode and cathode compartments are separated by a salt bridge which allows the ions in the solution to flow. The cations will migrate to the cathode and the anions will migrate to the anode. The overall cell reaction is

Cu+2 (aq, 1.0 M) + 2 e- Cu+2 (aq, 0.10 M) + 2 e- (9-3)

103


The potential generated by the above cell is governed by the Nernst Equation

eqn (9-4)

where

Ecell is the potential of the concentration cell

Eocell is the standard cell potential, in this case, Eo

cell = 0 *

R is the gas constant, 8.314 J/moleK T is the room temperature in Kelvin

n is the number of electrons transferred, n=2 in this case

C2 is 0.10 M

C1 is 1.0 M

F is Faraday's constant, 96500 coulombs/mole of electrons

* Eocell is zero because the standard half-cell potentials cancel (Eo

copper (oxid) + Eocopper (red) = 0).

In the concentration cell, the half-reactions are the same but the concentrations are different. This results in a nonstandard cell potential and Ecell does not equal zero. There is a 10-fold higher concentration in the cathode than the anode. As a result, at 298K, the overall potential of the above concentration cell is 0.0296 V or 29.6 mV.

Cell Notation

We can convey an electrochemical cell by using the cell notation rather than drawing a picture or an actual diagram of the cell. Cell notation lists the metals and ions involved in the reaction. A vertical line, |, denotes a phase boundary. A double line, ||, represents the salt bridge. The anode is always written on the left and the cathode on the right. The general form is:

anode | electrolyte of anode || electrolyte of cathode | cathode

The cell notation for the above concentration cell is,

Cu | Cu+2 (0.10 M) || Cu+2 (1.0 M) | Cu

In this experiment, we will set up a concentration cell as shown in Figure 9-1.

104


Figure 9-1 - A concentration cell.

105


Apparatus:

1. Digital pH meter 2. Lab jack 3. magnetic stirrer 4. salt bridge

Solutions:

1. 0.010 M AgNO3 2. 0.010 M NaCl 3. 0.010 M NaBr 4. 0.010 M NaI 5. saturated NH4NO3

106


Procedure:

1. Polish two silver electrodes and arrange apparatus as shown in Figure 9-2.

Figure 9-2 A concentration cell to determine the Ksp value of AgCl

2. Connect the electrode in the 0.010 M Ag+ to the positive terminal of the pH meter and the second electrode to the negative terminal of the pH meter.

3. Add about 10 to 15 drops of AgNO3 to the beaker containing 0.010 M sodium chloride to form AgCl precipitate (ie - until milky). Start the stirrer and record the voltage of the cell together with the temperature. Enter the data on the datasheet for Cell 1. Describe the cell using cell notation.

4. Repeat this procedure using 0.010 M sodium bromide solution (Cell 2) and 0.010 M sodium iodide solution (Cell 3) in place of the sodium chloride solution.

107


Datasheet:

Temperature: _____________________________________

CELLS CONCENTRATION CELL NOTATION | ΔE| (millivolts)

1 (cell with

NaCl solution)

2 (cell with

NaBr solution)

3 (cell with

NaI solution)

108


Postlab Questions:

1. Write the Ksp equilibrium reactions for the three insoluble solids, AgCl, AgBr, and AgI.

Ksp equilibrium reactions

AgCl

AgBr

AgI

2. Calculate the concentration of Ag+, [Ag+]2, in each of the silver halide solutions (ie - AgCl, AgBr, and AgI). For silver concentration cells at 25oC, eqn (9-4) becomes

eqn (9-5)

where E = measured potential in volts, n = 1 mole / equivalent, F = Faraday's constant, 96500 coulombs/mole e- R = gas constant, 8.314 jouiles/mole K T = temperature in Kelvin [Ag+]1 = 0.0100 M.

Provide one sample calculation using AgCl

109


Concentration for each cell is:

CELLS [Ag+] (moles/L)

(I)

(II)

(III)

3. Calculate Ksp for each of the three salts. You may assume that the halide concentration in each case is 0.0100 M. For example: Ksp = [Ag+]2 [X-]

where [Ag+]22 is determined from eqn (9-5), [X-] is assumed to be 0.0100 M.

Provide the calcuation of Ksp for AgCl as a sample calculation.

110


Tabulate calculated values of Ksp together with the accepted values of the solubility product constants. Cite the reference source and comment on your results.

Ksp (experimentally determined)

Ksp (accepted values)

AgCl

AgBr

AgI

4. Calculate the % error in the Ksp values for each of the salt. % error = | Ksp (expt) - Ksp (accepted value) | / Ksp (accepted value) * 100

Provide one sample calculation using % error of Ksp(AgCl)

111


% error

Ksp (AgCl)

Ksp (AgBr)

Ksp (AgI)

5. What are the possible sources of errors?

112

foundations of applied chemistry chem 0012 lab...

Documents