packet chapter 15 - aqueous equilibrium 2010

Values of Ka for Some Common Monoprotic Acids

PAGE 1AP CHEMISTRY

CHAPTER 15 - AQUEOUS EQUILIBRIUM

DateAgendaHomework

Mon3/1Test - Chapter 14

Tues3/21/2 daypacket - common ion effect and buffersLab - Job's Method Set Upbuffer problem page 7 and 8calculate mole fraction Cu and OH in each tt

Wed3/31/2 dayLab Job's Method Part 2, 3 and 4work on lab report

Thurs3/41/2 daygraphs and discussionfinish lab reportDue Friday

Fri3/5packet - strong base to a bufferpacket - buffering capacityproblem #7 (page 18)hw problems (page 20-21)

Mon3/8Quiz - buffersSelecting a buffer, pH and pKa packet - Henderson-Hasselbalch Equation Packet -hw problems (page 26)

Tues3/9Review ProceduresLab-making a buffer (page 27)lab report

Wed3/10Titration with strong acid and strong baseTitration with a weak acid and strong baseproblem # 2 (page 33-34)

Thurs3/11Titration with weak base and strong acidin class problems (page 36-38)

Fri3/12go over homeworkQuiz - titrationsPre-lab Titrationpacket problem - weak acid-strong base p 33Study for quiz

Mon3/15Go over quizLab - Part 1: Standardization

Tues3/16Part 2: Molar Mass of an unknown acidlab calculations

Wed3/17Discuss labslab calculationsLab due Thursday

Thurs3/18Quiz - part 2 (weak base strong acid)practice problems study for test

Fri3/19Test - Aqueous Equilibrium

Values of Ka for Some Common Monoprotic Acids

NameFormulaValue of Ka

Hydrogen sulfate ion Chlorous acid Monochloracetic acid Hydrofluoric acid Nitrous acid Formic acid Lactic acid Benzoic acid Acetic acid Hydrated aluminum(III) ion Propanoic acid Hypochlorous acid Hypobromous acid Hvdrocyanic aid Boric acid Ammonium ion Phenol Hypoiodous acid HSO4- HClO2 HC2H3ClO2 HF HNO2 HCO2H HC3H5O3 HC7H502 HC2H3O2 [Al (H20)6]3+HC3H5O2 HOCl HOBr HCN H3BO3 NH4+HOC6H5HOI 1.2 x 10-2 1.2 x 10-2 1.35 x 10 -37.2 x 10-44.0 x 10-4 1.8 x 10-4 1.38 x 10-4 6.4 x 10-5 1.8 x 10-5 1.4 x 10-5 1.3 x 10-5 3.5 x 10-8 2 x 10-96.2 x 10-10 5.8 x 10-10 5.6 x 10-10 1.6 x 10-10 2 x 10-11

Stepwise dissociation constants for several common polyprotic acids

NameFormulaKa1Ka2Ka3

Phosphoric acid Arsenic acid Carbonic acid Sulfuric acid Sulfurous acid Hydrosulfuric acid Oxalic acid Ascorbic acid (Vitamin C) Citric acid H3P04 H3AsO4 H2CO3 H2S04 H2S03 H2S H2C204 H2C6H6O6 H3C6H5O7 7.5 x 10-3 5 x 10-3 4.3 x 10-7 Large 1.5 x 10-2 1.0 x 10-7 6.5 x 10-2 7.9 x 10-58.4 x 10-4 6.2 x 10-8 8 x 10-8 5.6 x 10-11 1.2 x 10-21.0 x 10-7 ~10-196.1 x 10-5 1.6 x 10-12 1.8 x l0-5 4.8 x 10-136 x 10-10

4.0 x 10-6

Values of Kb for some common weak bases

NameFormulaConjugate acidKb

Ammonia Methylamine Ethylamine Diethylamine Triethylamine Hydroxylamine Hydrazine Aniline Pyridine NH3 CH3NH2 C2H5NH2 (C2H5)2NH (C2H5)N HONH2 H2NNH2 C6H5NH2 C5H5N NH4+ CH3NH3+ C2H5NH3+ (C2H5)2NH2+ (C2H5)NH+ HONH3+ H2NNH3+ C6H5NH3+ C5H5NH+1.8 x 10-5 4.38 x 10-4 5.6 x 10-4 1.3 x 10-3 4.0 x 10-41.1 x 10-83.0 x 10-6 3.8 x 10-101.7 x 10-9

Common Ion Effect Consider the following reaction: HF(aq) ( H+(aq) + F-(aq) Initially in solution, what will take place when we add the HF?

What would happen if we were to add F- or H+ to the reaction? Which way would the equilibrium shift?

This shift is called the common ion effect. We are adding more of one of the ions already in solution. This makes this solution less acidic by inhibiting the HF from dissociating. This will alter our pH, however, our favorite steps for solving these types of problems are still the same! YEA!! Heres a problemtry it! Problem #1Calculate the pH, and the percent dissociation of the acid, in each of the following solution. 0.200 M CH3COOH (Ka= 1.8 x 10-5)

STEP 1: List the major species

STEP 2: Choose the species that can dominate the reaction by comparing K values

STEP 3: Find the initial concentrations of all species involved.

STEP 4: List the initial, change and final concentrations in terms of x. ICE

STEP 5: Check the 5% rule

STEP 6: Find the pH. [H+] and % dissociation of HF.

Problem #2Find the pH and the percent dissociation of the acid, in the following solution. 0.200 M CH3COOH in the presence of 0.500 M NaCH3COO. STEP 1: List the major species

STEP 2: Choose the species that can dominate the reaction by comparing values




STEP 6: Find the pH, [H] and % dissociation of CH3COOH.

Buffered Solutions

These are one of the most important acid-base solutions. The most common and maybe most important example is blood. A buffered solution, such as blood, has the ability to resist a change pH. All a buffered solution is a weak acid and its salt (HF and NaF) or a weak base and its salt (NH3 and NH4Cl). Try it!

Problem # 3Calculate the pH of a solution that contains 0.250 M formic acid, HCOOH(Ka 1.8x10-4) and 0.100 M sodium formate NaCOOH.

STEP 1: List the major species - include whatever acid/base properties the species have




STEP 5: Check the 5 % rule

STEP 6: Find the pH, [H+] and % dissociation

Buffer Review Problems

1. Calculate the pH of each of the following solutions.

a. 0.100 M propanoic acid (HC3H5O2 Ka = 1.3 x 10-5)

b. 0.100 M sodium propanoate (NaC3H5O2)

c. Pure H2O

d. A mixture containing 0.100 M HC3H5O2 and 0.100 M NaC3H5O2e. Compare the percent dissociation of the acid in a with the acid in d. Explain the large difference in the percent dissociation of the acid.

2. Calculate the pH of a solution which is 1.00 M HF and 1.00 M KF.

3. Calculate the pH of a solution which is 0.50 M CH3NH2 and 0.70 M CH3NH3Cl.

Basic Buffers

To buffer a solution at a basic pH requires a weak base and its salt. For example, one might use NH3 (Kb 1.8 x 10-5) acid NH4Cl.

Problem #4Consider a solution containing 0.30 M NH3 and 0.20 M NH4Cl.







Strong Base Quick refresher. Calculate the pH of the solution formed by added 10.0 mL of 6.0 M NaOH to 500 mL of water. You know the deal...

Addition of a Strong Base to a Buffer

Problem #5Calculate the pH of the 0.250 M HCOOH/0.100 M NaCOOH buffer that we used before after the addition of 10.0 mL of 6.0 M NaOH to the original buffered solution volume of 500 mL.

Strategy: Think of this problem in TWO parts- a stoichiometry part (determine the number of moles of weak acid AFTER the addition of the strong base) and the equilibrium part (which weve been doing- but now there are new concentrations)

STOICHIOMETRY PART STEP 1: List the major species

STEP 2: decide what reaction will dominate- REMEMBER- the goal of acid meeting base is to NEUTRALIZE!

STEP 3: Find the number of MOLES of the major species.

Moles of HCOOH initial,

Moles of COOH- initial,

Moles of OH- added =

STEP 4: Make a MOLE ICE chart. This chart will involve numbers and NOT xs

STEP 5: Using the new mole values for whatever species are left, calculate the new initial concentrations.

EQUILIBRIUM PART- back to same old stuff...

...) STEP 1: List the major species - (ok, we know what they are, but just for old time sake)

STEP 2: Choose the species that can dominate the reaction




STEP 6: Find the pH

1 Practice with Buffers

Problem #6A solution is prepared by adding 31.56 g of NaCN and 22.30 g of HCN to 600.0 mL of water (Kb for HCN 6.2 x 10-10) A. What is the pH of this solution?







B. What is the pH after adding 50.0 mL of 3.00 M HCl?STOICHIOMETRY PART STEP 1: List the major species



Moles of HCN initial:

Moles of CN-initial: Moles of H+ added:


STEP 5: Using the new mole values for whatever species are left, calculate the new initial concentrations.

EQUILIBRIUM PART- back to same old stuff... STEP 1: List the major species (ok. we know what they are, but just for old times sake...)

STEP 2: Write the dominant reaction.

STEP 3: List the initial concentrations of the major species.

STEP 4: Make on ICE chart showing the change in equilibrium

STEP 5: Substitute back into the equilibrium expression and solve the easy way.


STEP 7: Calculate the pH

C. What is the pH after a further addition of 80.0 mL of 4.00 M NaOH? STOICHIOMETRY PART STEP 1: List the major species



Moles of HCNinitial

Moles of CN-initial

Moles of OH- added


STEP 5: Using the new mole values for whatever species ore left, calculate the new initial concentrations.

EQUILIBRIUM PART- back to some old stuff... STEP 1: List the major species (ok, we know what they are, but just for old time sake...)

STEP 2: Write the dominant reaction.

STEP 3: List the initial concentrations of the major species.

STEP 4: Make an ICE chart showing the change in equilibrium

STEP 5: Substitute back into the equilibrium expression and solve the easy way.


STEP 7: Calculate the pH

Buffers

Problem #7 Consider a 1.0 L solution containing 0.30 M NH3 and 0.20 M NH4Cl. A. Calculate the pH of this solution B. Calculate the pH after 0.050 mols HCl is added to this solution. C. Calculate the pH after 0.030 mol of NaOH(s) is added to the original solution.

Buffering Capacity A buffer with a large capacity contains large concentrations of buffering components and so can absorb a relatively large concentration of buffering components and so can absorb a relatively large amount of protons or hydroxide ions ond show little pH change. The pH of a buffered solution is determined by the ratio [A-]/[HA]. The capacity of a buffered solution is determined by the magnitudes of [HA] and [A-] Problem #8Calculate the pH of a 0.500 L solution that contains a 0.15 M HCOOH (Ka= 1.8 x 10-4 and 0.20 M NaCOOH. Then calculate the pH of the solution after the addition of 10.0 mL of 12.0 M NaOH.

HW Strong base/Strong acid added to buffers

1a. Calculate the pH of a solution which is 1.00 M HNO2 and 1.00 M NaNO2 .

1b. Calculate the pH after 0.10 mole of NaOH is added to 1.00 L of the above solution.

1c. Calculate the pH after 0.20 mol of HCl is added to 1.00 L of the original solution.

2a. Calculate the pH of a solution which is 0.60 M HF and 1.00 M KF.

2b. Calculate the pH after 0.10 mole of NaOH is added to 1.00 L of the above solution.

2c. Calculate the pH after 0.20 mol of HCl is added to 1.00 L of the solution above.

Problem #9Selecting a Buffer We wish to buffer a solution at pH = 10.07. Which one of the following bases (and conjugate acid salts) would be most useful? A. NH3 (Kb = 1.8 x 10-5) B. C6H5NH2 (Kb = 4.2 x 10-10) C. N2H4(Kb = 9.6 x 10-7)

pH and pKa

The Ka of propionic acid HC3H5O2 is 1.34 x 10-5. What is the pH when [HC3H502]= [C3H5O-]?

We dont know the values, but we know they are equal. Write the equation for the dissociation of HA and then write the K expression and solve

Or, maybe try writing the equation for the K reaction with water. Write the Kb expression and solve the problem that way.

How do the values compare?

Try solving the problem using the Henderson-Hasselbaich equation. Henderson-Hasselbalch EquationFor the dissociation of a weak acid,

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

Ka = [H+][A-]= [H+] [A-]

[HA]

[HA]

Now take the log of both sides

log Ka = log [H+] + log [A-]

[HA](remember that the log of a product is the sum of the logs, ie. Log xy = log x + log y) Multiply both sides of the equation by -1-log Ka = -log [H+] - log [A-]

[HA]

pKa = pH log [A-]

[HA]

pH = kPa + log [A-]

[HA]

Example:The acid HOCl has a pKa value of 7.5. Calculate the pH of a solution

containing 0.25 M HOCl and 0.75 M NaOCl.

pH = kPa + log [A-]

[HA]

= 7.50 + log 0.75

0.25

= 7.50 + log (3.0)

=7.50 + 0.48

= 7.98Review of ProceduresRemember: The most important part of doing an acid-base problem is the analysis at the beginning of the problem.

Does a reaction occur?

What is it?

What equilibrium dominates?

The best way to answer these questions successfully is to write the major species in solution. Then ask the question: Does a reaction occur that goes to completion? The situations to look for are:

Has OH- been added to a solution containing an acid?

Has H+ been added to a solution containing a base?

In both of these situations, the reaction that occurs can be assumed to go to completion. After the reaction has been allowed to go to completion, again write the major species. Now check each one for acid or base properties and select the dominant equilibrium by looking at the values of the various equilibrium constants. In almost every case, one equilibrium will dominate and can be used to solve for the [H+] or [OH-].

When faced with an acid-base problem, the best strategy is to assume that it is not like exactly like any other problem you have done. One small change can cause a problem that looks very similar to one you have done before to be quite different.When starting an acid-base problem, the wrong question to ask is: How can I use a problem whose solution I have memorized to solve this problem? The correct questin should be: What species are in solution nd what do they do? (think Chemistry!) Let the problem guide you.

The steps to follow are:

1. Write the major species in solution before any reaction takes place.

2. Look for any reactions taking place that can be assumed to go to completion.

Examples:OH- with acid (strong or weak)

H+ with a base (strong or weak)

3. If a reaction occurs that can be assumed complete:

a) do the stoichiometry problem

b) write the major species in solution after the reaction

4. Look at each major component of the solution and decide which are acids or bases.

5. Pick the equilibrium that will control the [H+]. Use the Ks for the various species to help decide.6. Do the equilibrium calculation.

a) write the equation for the reaction and the equilibrium expressions.

b) compute the initial concentrations (assuming the equilibrium of interest has not

yet accourred, ie no acid dissociation, etc)

c) define x

d) compute the equilibrium concentrations in terms of x

e) substitute in the equilibrium expression and solve for x, making approximations

to simplify the math if possible

f) check the validity of the approximations

g) calculate the pH

HOMEWORK

1. Calculate the pH of each of the following buffered solutions.

HC2H3O2

Ka = 1.8 x 10-5C2H5NH2

Kb = 5.6 x 10-4a. 0.10 M acetic acid/0.25 M sodium acetate

b. 0.25 M acetic acid/0.10 M sodium acetate

c. 0.50 M C2H5NH2/0.25 M C2H5NH3Cl

d. 0.25 M C2H5NH2/0.50 M C2H5NH3Cl

2. Calculate the ratio of [NH3]/[NH4+] in ammonia/ammonium chloride buffered solutions with the following pH values:

NH3

Kb = 1.8 x 10-5a. pH = 9.00

b. pH = 8.80

3. Choose which solution(s) (no calculations) will be a buffered solution and explain why.

a. A solution containing 0.1 M KNO3 and 0.1 M HNO3

b. A solution containing 0.1 M NaNO2 and 0.15 M HNO2

c. A solution containing 0.5 M HCl and 0.5 M KCl

d. A solution containing 0.1 M NaOH and 0.1 M NaCl

e. A solution containing 0.3 M HF and 0.5 M KF

4. Calculate the pH after 0.15 mol solid NaOH is added to 1.00 L of each of the following buffered solutions.

a. 0.050 M propanoic acid (HC3H5O2 Ka = 1.3 x 10-5)/0.080 M sodium propanoate

b. 0.050 M propanoic acid/0.80 M sodium propanoate

c. Is the solution in part a still a buffered solution after the NaOH has been added? Explain.

LAB - Buffers

Preparation of a Buffer Solution at a given pH

Purpose:

1) Practice making a buffer calculations

2) Choose and make a buffer of a given pH by dissolving known masses of solid salts in water

Introduction:

A buffer solution resists changes in pH by containing a component that will react with added acid or added base. Buffers are found naturally in blood, sea water and food.

Because the acid and base in the buffer must co-exist in solution and not react to neutralize each other, equilibrium between the acid and base part of the buffer must be established. So the buffer system must contain either a weak acid (HA) and its conjugate base (A-) or a weak base (B) and its conjugate base (HB+).

Write the equilibrium reaction and the Ka expression for each of these buffer solutions:

Possible solids to choose from:

NaHSO3

Na2SO3

NaCH3COO

Na3PO4(12H2ONa2HPO4(7H2OKH2PO4

NaHCO3

Na2CO3Na2SO4

KHSO4

KCl

NaCl

**these salts may be hydrated, so pay attention to the actual formulas on the stock bottlesProcedure:

1) Calculate the pKa and Ka for a solution with the following pH

a) 2.25b) 6.80c) 7.50d) 11.10e) 12.50

2) Determine which weak acid to use as a buffer in the region that you have selected

3) Determine many grams of each salt is needed to make the buffer solution

4) Make your buffer solution

5) Use the pH probe provided and measure the actual pH of your buffer solution

Include in conclusion*What is a buffer* How do you determine the best buffer to use

*Propose two reasons why the actual pH of your buffer solution might not be

identical to the theoretical pH. Explain in detail how these errors would lead to the pH you measured.

*What would happen to the pH of your buffer solution if it were diluted with 100 mL of distilled water? Explain TITRATIONS!!!

A titration is commonly used to determine the amount of acid or base in a solution.

Need: a) a solution of known concentration (titrant)

b) solution of unknown concentration

The titrant is delivered by a buret into the unknown solution the substance being analyzed is just consumed. The stoichiometric equivalence point is indicated by the color change of an indicator.

Progress is monitored by plotting the pH of the solution being analyzied as a function of the amount of titrant added. This is called a pH or titration curve.

Strong Acid-Strong Base

The net ionic reaction for a strong acid-strong base titration is

Molarity = moles solute = mmoles solute

L solution mL solution

Problem #1 Strong Acid Strong Base Titration

Consider the titration of : 50.0 mL of 0.200 M HNO3

With 0.100 M NaOHWe will calculate the pH of the solution at various selected points in the course of the titration where specific volumes have been added.

POINT A: No NaOH has been added

POINT B: 10.0 mL of 0.100 M NaOH has been added

POINT C: 20.0 mL (total) of 0.100 M NaOH has been added

POINT D: 50.0 mL (total) of 0.100 M NaOH has been addedPOINT E: 100.0 mL (total) of 0.100 M NaOH has been added

POINT F: 150.0 mL (total) of 0.100 M NaOH has been added

POINT G: 200.0 mL (total) of 0.100 M NaOH has been added

Draw the pH curve for the above results.

REMEMBER:

1. Before the equivalence point [H+] (and

hence the pH) can be calculated by dividing

the number of moles of H+ remaining by the

total volume of the solution in liters

2. At the equivalence point, the pH is 7

3. After the equivalence point [OH-] can be determined by dividing the number of moles of OH- remaining by the total volume of the solution in liters. Then [H+] is obtained from Kw

HOMEWORK:

Problem #2 Strong Acid Strong Base Titration

Calculate the pH after the following total volumes of 0.250 M HCl have been added to 50.00 mL of 0.1500 M NaOH

A. 0.00 mL

B. 4.00 mL

C. 29.50 mL

D. 30.00 mL

E. 30.50 mL

F. 40.00 mL

Titration of Weak Acids with Strong Bases

Strong acids and strong bases are fairly straightforward. However, when we are using a weak acid, there is a major difference: to calculate [H+] after a certain amount of strong base has been added, we must deal with the weak acid dissociation equilibrium. This is similar to a buffered solution. It is important to remember that even though the acid is weak, it reacts to completion with the OH-.Problem #1: Weak Acid - Strong Base Titration

We will consider the titration of 50.0 mL of 0.10 M acetic acid (Ka = 1.8 x 10-5) with 0.10 M NaOH

A. No NaOH has been added

B. 10.0 mL of 0.10 M NaOH has been added

C. 25.0 mL (total) of 0.10 M NaOH has been added

D. 40.0 mL(total) of 0.10 M NaOH has been added

E. 50.0 mL (total) of 0.10 M NaOH has been added

F. 60.0 mL (total) of 0.10 M NaOH has been added

G. 75.0 mL (total) of 0.10 M NaOH has been added

Draw the pH curve for the above results

Problem # 2 Weak Acid Strong Base Titration

Calculate the pH after the following total volumes of 0.4000M NaOH are added to 50.00 mL of 0.2000 M HCOOH (Ka = 1.8 x 10-4)A. 0.00 mL

B. 5.00 mL

C. 12.50 mL

D. 24.50 mL

E. 25.00 mLF. 25.50 mL

G. 40.00 mL

Problem # 3 Weak Base Strong Acid Titration

Calculate the pH at each of the following points in the titration of 50.00 mL of 0.0100M sodium phenoalate (NaOC6H5 ) solution with 1.000 M HCl solution (Ka = 1.05 x 10 -10)

A. Initially

B. Midpoint

C. Equivalence Point

In Class Review Problems:

1. A 100. mL sample of 0.100 M NH4Cl solution was added to 80.0 mL of a 0.200 M NH3 solution. The value of Kc for ammonia is 1.79 x 10 -5.

A. What is the value of pKb for ammonia

B. What is the pH of the solution described in the question?

C. If 0.200 grams of NaOH were added to the solution, what would be the new pH of the solution? (assume that the volume of the solution does not change)

D. If equal molar quantities of NH3 and NH4+ were mixed in the solution, what would be the pH of the solution?

2. H3PO4 (( H+ + H2PO4-

K1 = 7.5 x 10-3 H2PO4- (( H+ + HPO42-

K2 = 6.2 x 10-8 HPO42- (( H+ + PO43-

K3 = 2.2 x 10-13A. Choose an amphoteric species from the reactions listed above and give its conjugate acid and conjugate base.

B. Explain why the dissociation constant decreases with each hydrogen ion lost

C. Of the acids listed above, which would be the most useful in creating a buffer solution with a pH of 7.5

D. Sketch the titration curve that results when H3PO4 is titrated with excess NaOH and label the two axes

3. Use the principles of acid-base theory to answer the following questions.A. Predict whether a 0.1 M solution of sodium acetate, NaC2H3O2, will be acidic or basic and give a reaction occurring with water that supports your conclusion

B. B. Predict whether a 0.1 M solution of ammonium chloride, NH4Cl, will be acidic or basic and give a reaction occurring with water that supports your conclusion

C. Explain why buffer solutions are made with weak acids instead of strong acids.

NO MATH!!

Simply write any chemical equations associated with the question and explain the chemistry

1. Acetic acid, CH3COOH, added to water

2. CH3COOH added to water in the presence of NaCH3COO

3. Formic acid, HCOOH (Ka=1.8x10-4) added to water in the presence of sodium formate NaCOOH.

4. Consider a solution containing NH3 and NH4Cl.

5. NaOH to 500 mL of water.

6. HCOOH/NaCOOH buffer with NaOH added to it.

7. A solution is prepared by adding NaCN and HCN water (Kb for HCN 6.2 x 10-10)

8. NH3 added to water

9. N2H4 in solution

10. propionic acid, HC3H5O2 in solution

LAB -

Preparation of a Buffer Solution at a given pH

Purpose:

1) Practice making a buffer calculations

2) Choose and make a buffer of a given pH by dissolving known masses of solid salts in water

Introduction:

A buffer solution resists changes in pH by containing a component that will react with added acid or added base. Buffers are found naturally in blood, sea water and food.

Because the acid and base in the buffer must co-exist in solution and not react to neutralize each other, equilibrium between the acid and base part of the buffer must be established. So the buffer system must contain either a weak acid (HA) and its conjugate base (A-) or a weak base (B) and its conjugate base (HB+).

Write the equilibrium reaction and the Ka expression for each of these buffer solutions:

Possible solids to choose from:

NaHSO3

Na2SO3

NaCH3COO

Na3PO4Na2HPO4

KH2PO4

NaHCO3

Na2CO3Na2SO4

KHSO4

KCl

NaCl

**these salts may be hydrated, so pay attention to the actual formulas on the stock bottlesProcedure:

1. Calculate the pKa and Ka for a solution with the following pH

a) 2.25b) 6.80c) 7.20d) 11.10e) 12.50

2. Determine which weak acid to use as a buffer in the region that you have selected

3. Determine many grams of each salt is needed to make the buffer solution

4. Make your buffer solution

5. Use the pH probe provided and measure the actual pH of your buffer solution

Include in conclusion*What is a buffer* How do you determine the best buffer to use*Propose two reasons why the actual pH of your buffer solution might not be

identical to the theoretical pH. Explain in detail how these errors would lead to the pH you measured.

*What would happen to the pH of your buffer solution if it were diluted with 100 mL of distilled water? Explain LAB Aqueous Stoichiometry

Modified from Laboratory Investigations AP* Chemistry (Hostage/Fossett)

Purpose:

1) Determine the stoichiometric of a chemical reaction experimentally

2) Determine the chemical formula of a precipitate

3) Determine the oxidation state of an ion in solution

Day 1 Setup

1. You will be assigned one or two test tubes to set up for this lab.

2. Use a clean and rinsed buret to deliver the exact volume of copper ion solution to your test tube(s). Record the exact volume added in your data table.

3. Use a clean and rinsed buret to deliver the exact volume of hydroxide ion solution to your test tube(s). Record the exact volume added in your data table.

Test Tube #Volume copper ion (ml)Volume OH- (ml)

10.0010.00

21.009.00

32.008.00

42.507.50

53.007.00

63.336.67

74.006.00

85.005.00

96.004.00

107.003.00

118.002.00

129.001.00

1310.000.00

4. Stir the mixture carefully with a stirring rod and then cover the test tube(s) with parafilm

5. Arrange the test tubes side by side in the rack in order of increasing volume of hydroxide ion added.

6. Let the test tubes sit until day 2

Day 2 - Height of Precipitate and Qualitative Observations1. Placing a sheet of white paper behind the samples, carefully observe both the precipitate and the supernatant (that is, the solution on top of the solid). Record you observations.

2. Remove your test tube(s) from the rack and rest its bottom on the lab bench. Hold the test tube straight up.

3. Measure the height of the precipitate.

4. Add your data to the table on the computer.

Day 3 pH

1. Without disturbing the precipitate, carefully place the pH probe in your test tube(s) so the glass bulb on the bottom is covered by the solution.

2. Allow the pH reading to become constant. Then record the pH reading in the data table on the computer.

Day 4 - Absorbance at 635 nm

1. Be sure the spectrophotometer is warmed up. Follow the directions to zero the colorimeter at 635nm.

2. Without disturbing the precipitate, carefully remove a sample from the test tube, using a transfer pipet, and place it in a cuvette.

3. Be sure no solid is in the sample, then remove all air bubbles. Wipe the outside of the cuvette with a damp paper towel to remove any fingerprints or dirt.

4. Place the cuvette in the colorimeter and close the lid.

5. Record the absorbance at 635 nm in the data table on the computer.

Day 5 Results1. Graph height precipitate vs. mole fraction Cu ion

2. Graph pH vs. mole fraction Cu ion

3. Graph pH vs. mole fraction OH ion

4. Graph absorbance vs. mole fraction Cu ionData:Test TubeXCuXOHHeight of pptpH of solutionAbsorbance of solution

1

2

3

4

5

6

7

8

9

10

11

12

13

Include in conclusion

*discuss the ppt reaction

*discuss mole fraction

*discuss how the height is used to determine the stoichiometry of the reaction

*discuss pH as it pertains to this reaction

*discuss how the pH is used to determine the stoichiometry of the reaction

*discuss absorbance as it pertains to this reaction

*discuss how the absorbance is used to determine the stoichiometry of the reaction

*discuss results/sources of error

LABTitration-Part 1Standardization of NaOH Solution

Purpose:

To determine the concentration of a sodium hydroxide solution by titrating against a primary standard acidic substance, potassium hydrogen phthalate KHC8H4O4 (KHP, 204.22 g/mol)Procedure:

Preparation of Buret

1. Rinse the buret with a few milliliters of distilled water three times

2. Rinse the buret with NaOH solution three times

3. Fill the buret with fresh NaOH solution over the 0.00 mL mark and then drain until the bottom of the meniscus is on the 0.00 mL mark

Standardization of NaOH Solution

1. Measure 1.4 1.6 g of KHP into a weigh dish on an analytical balance and transfer to a 50 mL Erlenmeyer flask

2. Add about 5 mL of distilled water and 2 3 drops of phenolphthalein indicator solution to the flask

3. Make sure all the KHP is dissolved in the solution. If it is not, slightly heat the solution until all the KHP is dissolved

4. Begin adding the NaOH solution from the buret to the sample in the Erlenmeyer flask, swirling the flask constantly during the addition

5. As the NaOH solution enters the solution in the Erlenmeyer flask, streaks of pink will be visible. They will fade as the flask is swirled.

6. Eventually the pink streaks will persist for longer periods of time. This indicates the approach of the endpoint of the titration 7. Begin adding the NaOH dropwise, with constant swirling, until a single drop of NaOH causes a permanent pale pink color that does not fade

8. Record the reading on the buret to the nearest 0.02 mL

9. Repeat steps 1-4 two more times, for a total of three trials

10. Calculate the moles of KHP used

11. Calculate the moles of NaOH in the solution (The reaction of NaOH and KHP is of 1:1 stoichiometry)

12. Calculate the molarity of the NaOH solution 13. Find the average molarity of NaOH

Titration-Part 2

Analysis of a Solid Acid

Purpose:

To determine the molecular weight of an unknown solid acid by titration with standardized NaOH solution

Procedure:

Preparation of Buret

1. Rinse the buret with a few milliliters of distilled water three times

2. Rinse the buret with NaOH solution three times

3. Fill the buret with fresh NaOH solution over the 0.00 mL mark and then

drain until the bottom of the meniscus is on the 0.00 mL markAnalysis of a Solid Acid

1. Weigh 0.5 0.7 g of unknown solid acid in a weigh dish on an analytical balance and transfer to a 50 mL Erlenmeyer flask

2. Dissolve the unknown in about 5 mL of distilled weater and add 2 3 drops of phenolphthalein indictor solution. 3. Make sure all the unknown acid is dissolved in the solution. If it is not, slightly heat the solution until all the unknown acid is dissolved

4. Begin adding the NaOH solution from the buret to the sample in the Erlenmeyer flask, swirling the flask constantly during the addition

5. As the NaOH solution enters the solution in the Erlenmeyer flask, streaks of pink will be visible. They will fade as the flask is swirled.

6. Eventually the pink streaks will persist for longer periods of time. This indicates the approach of the endpoint of the titration

7. Begin adding the NaOH dropwise, with constant swirling, until a single drop of NaOH causes a permanent pale pink color that does not fade

8. Record the final volume of NaOH used to reach the endpoint

9. Repeat steps 1-8 two more times for a total of three trials

10. Calculate the moles of NaOH used

11. Calculate the moles of unknown acid used (the acid is a triprotic acid)

12. Calculate the molecular mass of the unknown acid solid (g/mol)

13. Average the three molecular masses and compare to the actual

Data:

Part 1-Standardization

Trial 1Trail 2Trial 3

Mass KHP

Moles KHP

Moles NaOH

Volume NaOH

Molarity NaOH

Part 2-Analysis of Solid Acid

Trial 1Trial 2Trial 3

Mass unknown acid

Volume NaOH

Moles NaOH

Moles unknown acid

MW unknown acid

Include in conclusion:

*Define titration

*Discuss standardization of NaOH

*Discuss titration of solid acid with NaOH

*Discuss triprotic acids

*Discuss results/sources of error

packet chapter 15 - aqueous equilibrium 2010

Documents

common monoprotic acids

buffer page

fri35packet strong base

strong basetitration

strong baseproblem

upbuffer problem page

equilibrium shift

lab report wed310titration