CHEM 105b LAB FINAL

Significant figures of graduated cylinder?

3 sig figs.

Significant Figures of a burette?

4 sig figs.

Experiment #1 (Kinetics): What is the reaction?

2H2O2 -> 2H2O + O2

Experiment #1 (Kinetics): What is the reaction mechanism?

H2O2 + I- -> H2O + IO-

H2O2 + IO- -> H2O + O2 + I-

Experiment #1 (Kinetics): What is the rate law?

rate = k[H2O2]^m[I-]^n

Experiment #1 (Kinetics): Experiment Overview

Observe the amount of oxygen gas produced over time.

Experiment #1 (Kinetics): How to calculate the rate of relative volume produced?

Vol O2 (g) / (Vol sln * ∆t)

Experiment #1 (Kinetics): Why is it important to use relative volume of oxygen rather than absolute volume of oxygen?

We should use relative volume (Vol O2 / Vol Sln) in order to account for differences in total volume of each solution. Since concentration affects reaction rate, disparities in solution volume will affect our calculations, so we must use relative oxygen produced.

Experiment #1 (Kinetics): Why must temperature be held constant for trials 1-5?

Temperature affects the rate constant. If we have inconsistent temperatures, the reaction rate will vary and will make calculations inaccurate.

Experiment #1 (Kinetics): Objectives?

Study the reaction rate of a catalyzed reaction. How does changing concentration of H2O2 and KI affect reaction rate? Use initial concentrations and initial rates to determine rate law. How does changing temperature affect the reaction rate?

Experiment #1 (Kinetics): How to set up the apparatus?

unstoppered flask connected to Pipet A, which is connected to Pipet B. Fill pipet B with water until Pipet A is full and pipet B is empty. stopper flask and lower pipet B. There are no leaks if there is small water level changes but then becomes the same afterwards.

Experiment #1 (Kinetics): What temperatures will trials #1-5 and trial #6 be held in?

Trials #1-5 will be in room temperature (20 - 25˚C). Trial #6 will be be held INSIDE a warm water bath (30-35˚C).

Experiment #1 (Kinetics): What are the three constituents being used in each reaction and what is the total volume of each solution?

H2O2, KI, and DI water. Each solution should be ~15ml.

Experiment #1 (Kinetics): What should you be doing while the reaction is occurring?

One person will pour in H2O2 into the flask, stopper it, and swirl it. The other person should be making sure the water levels in Pipet A and Pipet B are even.

Experiment #1 (Kinetics): Why is it important to keep the water levels equal?

So that the pressure is even and that the volume of gas is measured at atmospheric pressure.

Experiment #1 (Kinetics): When should you record volume and time, and when should you stop recording data for that trial?

Record volume and time every 2ml of gas production (or when 2ml of water in Pipet A goes down).

Experiment #1 (Kinetics): What are the concentrations of H2O2 and KI we are using?

3% H2O2 (0.88 M) and 0.5M KI

Experiment #1 (Kinetics): What are the volumes of H2O2 and KI we are using for each trial?

Trial #1: 3 3

Trial #2: 4 3

Trial #3: 5 3

Trial #4: 3 4

Trial #5: 3 5

Trial #6: 3 3

Experiment #1 (Kinetics): Why must volumes of H2O2, KI, and H2O vary?

H2O must vary in order to keep the solution at 15ml. We must vary volumes of H2O2 and KI in order to see how different concentrations of each reactant affects the reaction. For some trials, H2O2 will be held constant, and other trials will have KI constant in order to make calculating the rate law easier.

Experiment #1 (Kinetics): What is the two point Arrhenius equation?

ln(k2/k1) = Ea/R(1/T1 - 1/T2)

where R = 8.314

Experiment #1 (Kinetics): According to the reaction mechanism, iodine is used and regenerated at the end of the reaction. How can we determine if this is true?

-We can use a metallic cation that produces a precipitate with iodine.

-we can add starch to the solution and if iodide is present, the solution will turn a dark purple/black color

-to determine final concentration of iodide at the end of reaction, use sodium thiosulfate and that will turn the reaction back to its clear color and use the amount of sodium thiosulfate to do stoichiometry

Experiment 1 common errors

-not fully balancing pipet A and pipet B; would cause measurement of volume of oxygen gas to not be read/measured accurately and would cause rate law to be determined wrong

-flask not stopped right away after the hydrogen peroxide was added, caused some of the oxygen gas to escape and not be measured in the collecting tube causing rate to be less than expected

Experiment #2 (Spectrophotometric Kinetics): What are the two solutions we are using?

Blue Dye #1 and OCl- (hypochlorite from bleach)

Experiment #2 (Spectrophotometric Kinetics): What happens when blue dye #1 and hypochlorite react?

Blue dye #1 loses its color and the solution becomes colorless.

Experiment #2 (Spectrophotometric Kinetics): How are we measuring the change in concentration of blue dye #1?

We are using spectrometers to measure the change in absorbance over time. We can then convert absorbance into concentration using Beer's Law.

Experiment #2 (Spectrophotometric Kinetics): What is Beer's Law?

Absorbance = (Concentration)(path of light usually 1 cm)(molar absorptivity constant)

Experiment #2 (Spectrophotometric Kinetics): After calculating concentration, how can we determine the rate order of blue dye #1?

We can graph concentration vs time using the zero, first, and second order integrated rate laws. Find which graph is most linear.

Experiment #2 (Spectrophotometric Kinetics): What is the overall reaction?

Blue Dye #1 (aq) + OCl- (aq) -> colorless product

Experiment #2 (Spectrophotometric Kinetics): What is the rate law?

rate = k[Blue Dye #1]^n[OCl-]^m

Experiment #2 Why do we use a large excess of [OCl-]?

When we use a large excess, that means with respect to the change in Blue #1, the concentration of hypochlorite is essentially constant.

Experiment #2 (Spectrophotometric Kinetics): What is a pseudo nth order reaction?

When we use very high concentrations of OCl-, it is essentially remaining constant throughout the reaction. Hence, the new rate law becomes:

rate = k'[Blue Dye #1]^n

Experiment #2 (Spectrophotometric Kinetics): What is a pseudo rate constant?

The pseudo rate constant is the new rate constant for a pseudo nth order reaction. It still depends on the concentration of NaOCl used, even though it is in large excess.

k' = k[OCl-]^m

Experiment #2 (Spectrophotometric Kinetics): For what absorbances is the relationship between Beer's law and concentration linear?

Between 0.1 - 1.0 absorbance.

Experiment #2 (Spectrophotometric Kinetics): Objectives

Using a spectrometer to measure the change in color of a solution. Learn how to measure absorbance vs time. Apply pseudo order reactions in order to separately calculate blue dye #1 order and OCl- order. Apply Beer's Law and integrated rate laws to absorbance vs time to determine blue dye order. Use calculated k' values to determine the value of k and OCl- concentration.

Experiment #2 (Spectrophotometric Kinetics): Why must we use clean and dry containers?

So that our solutions do not become diluted. Otherwise, absorbance and concentrations of Blue Dye #1 will go down.

Experiment #2 What is true for all trials?

-All trials will have a solution of 10.00 mL, should use volumetric flasks

-the initial concentrations of blue #1 will all be the same

-concentration of hypochlorite will be varied among trials

Experiment #2 (Spectrophotometric Kinetics): How will we know the reaction is complete?

When the solution is colorless.

Experiment #2 (Spectrophotometric Kinetics): What concentration of OCl- and Blue Dye #1 will we be using? How many trials?

We will choose a concentration and volume of blue dye #1 that yields an absorbance greater than 1. This is to account for the reaction starting right away and will give us time to transport the solution into the spectrophotometer. We can choose any volume of OCl- and we will increase it in every trial. There will be 4 trials.

Experiment #2 (Spectrophotometric Kinetics): What are the components of each mixture?

Blue Dye #1, NaOCl-, and H2O to make 10ml solutions.

Experiment #2 (Spectrophotometric Kinetics): How can we calculate OCl- rate order and rate constant from

k' = k[OCl-]^m

We can derive the equation

log(k') = m * log[OCl-] + log(k).

Where m is the slope and log(k) is the y-int.

Expriment #2 Common errors

• transferring the solution into the cuvette took too long and gave a smaller absorbance than expected, which leads to a smaller concentration, and thus the results are not as accurate

• not mixing the solution enough before putting it in the spectrophotometer leads to absorbance being off

Experiment #2 (Spectrophotometric Kinetics): if your absorption spectrum for blue dye #1 stock solution has a flattened peak at the max absorption, what should you do?

You should decrease the concentration of blue dye #1.

Experiment #3 (Le Chatelier's Principle): What are some signs of a reaction occurring?

Temperature change, color change, gas production, precipitation, turbidity, etc.

Experiment #3 (Le Chatelier's Principle): Define Le Chatelier's Principle

A system at equilibrium, when distressed, will shift in order to revert the stress.

Experiment #3 (Le Chatelier's Principle): What are some precautions to take when using a hot plate?

Make sure the glass beaker being placed on the hot plate is not cracked. We must also pick up the hot plate using tongs.

Experiment #3 (Le Chatelier's Principle): Objectives

Apply stress to chemical systems to observe how the systems react. Make qualitative observations about chemical and physical changes. Use pH paper to test acidity/basicity. Explain Le chatelier's principle using observations of reaction changes.

Experiment #3 (Le Chatelier's Principle): What are acid-base indicators?

Solutions that can be added to a solution and tells the acidity/basicity of a solution based on what color it makes the mixture. If there are a lot of H+ ions, it will be one color. If there are a lot of OH- ions, it will be another color.

Experiment #3 (Le Chatelier's Principle): Why is it necessary to use dry test tubes?

We must use dry test tubes because it could otherwise affect concentrations of solutions.

Experiment #3 Is HIn or In- responsible for the orange color of the methyl orange indicator?

HIn, adding sodium hydroxide (the base) makes the solution more yellow

Experiment #3 (Le Chatelier's Principle): Is the formation of [CoCl4]2- endothermic or exothermic?

It is exothermic as placing the test tube in heat caused the solution to turn blue, or toward the reactants side. This indicates that heat is a product. When the test tube is placed in cold temperatures, it becomes cold, indicating that heat was taken away.

Experiment #3 (Le Chatelier's Principle): Do magnesium ions react to form insoluble hydroxides?

The reaction creates a precipitate MgOH, indicating that insoluble hydroxides are formed.

Experiment #3 (Le Chatelier's Principle): Do magnesium ions form complex ions with hydroxide and ammonia?

Complex ions are not created from magnesium ions.

Experiment #3 (Le Chatelier's Principle): For the reaction Fe3+ + SCN- ⇌ FeSCN2+, the solution goes from yellow to dark red. Adding Fe(NO3)3 causes the color to change to what? Which way does the reaction shift?

Adding Fe(NO3)3 causes the solution to become darker since we're adding more reactants. It shifts toward the product side.

Experiment #4 (Spectrometers and K): What pH indicator are we using? What colors can it be?

Bromothymol blue. In acidic solutions it is yellow (HB). In basic solutions it is blue (B-).

Experiment #4 (Spectrometers and K): What is the reaction between bromothymol blue and water?

HB + H2O ⇌ B- + H3O+

Experiment #4 (Spectrometers and K): What is the equilibrium constant?

K = [B-][H3O+]/[HB]

Experiment #4 (Spectrometers and K): What are the peak absorbances when BTB is yellow (acidic) and when BTB is blue (basic).

432nm when yellow and 616nm when blue.

Experiment #4 (Spectrometers and K): What are the concentrations of [HB] and [B-]?

Concentrations of [HB] and [B-] will always add up to [BTB]. Their own individual concentrations depend on the acidity/basicity of the mixture.

Experiment #4 (Spectrometers and K): How can we calculate concentration of B-.

Absorbance (@616nm) = Molar Absorptivity l [B-]

Experiment #4 (Spectrometers and K): What is the purpose of part 1 of the experiment?

To determine the molar absorptivity of B- at 616nm.

Experiment #4 (Spectrometers and K): For part 1, what are concentrations of the components we are using. What is the total volume of the mixture?

20mL solutions of water, 2.5 * 10 ^-4 M BTB, and 0.25M K2HPO4.

Experiment #4 (Spectrometers and K): For part 1, what volumes of K2HPO4 and BTB are we using.

We are keeping K2HPO4 constant at 4mL, and increasing volumes of BTB by 0.5mL for every 5 mL. Start with 0.5mL of BTB end with 2.5 mL of BTB.

Experiment #4 (Spectrometers and K): For part 1, what do we expect the color of each solution to be based on the concentration of BTB.

Since the solution of K2HPO4 is basic, we expect the solution to be blue due to the presence of BTB. The blue color will become darker as we add more BTB.

Experiment #4 (Spectrometers and K): What is the purpose of part 2 of the experiment?

The purpose is to determine the the equilibrium constant of the indicator.

Experiment #4 (Spectrometers and K): For part 2, what solution from part 1 will you use?

The solution whose absorbance was closest but less than 1.

Experiment #4 (Spectrometers and K): For part 2, what new component are we using?

0.25M KH2PO4

Experiment #4 (Spectrometers and K): How will we create solutions for trials 7-10.

Use mL of BTB that yielded absorbance closest to 1 but less than 1. Use varying volumes of K2HPO4 and KH2PO4, where at least one trial has equal parts and one trial has 4mL KH2PO4 and no K2HPO4. All solutions should be 20mL.

Experiment #4 (Spectrometers and K): Do you expect solution 11 to be acidic, basic, or neutral?

Solution 11 should be acidic.

Experiment #4 (Spectrometers and K): What is the correlation between solution color and pH?

Solutions with blue color are basic pH, and solutions with yellowish color are acidic pH.

Experiment #4 (Spectrometers and K): How to calculate [H3O+]?

We can use

-log[H3O+] = pH and solve for [H3O+].

Experiment #4 (Spectrometers and K): How to find molar absorptivity?

Graph absorbance vs concentration of B- (we assume concentration of BTB = concentration of B-). The slope of this plot will be equal to molar absorptivity.

Experiment #4 (Spectrometers and K): How to calculate [HB]

Use [BTB] = [HB] + [B-]

Experiment #4 (Spectrometers and K): What is the trend in color change of the solutions?

The more basic the solution, the more blue it is. The more acidic, the darker yellow.

Experiment #4 (Spectrometers and K): How is this change observable in the absorption spectra?

Solutions that had deeper blue and more basic had a greater absorbance at 616nm, whereas solutions that were more acidic and yellow had greater absorbance at 434nm.

Experiment #4 Common Errors

• Transferring solutions from one beaker to another. Small amounts of solution may have been lost which causes less accurate dilution calculations

Experiment #5 (Buffer Capacity): How do you calculate Ka?

Ka = [H3O+][A-]/[HA]

Experiment #5 (Buffer Capacity): What is the Henderson equation?

pH = pKa + log([base]/[acid])

• only valid within 1 pH unit of the pKa of the acid

experiment #5 what is a buffer solution?

• contain appreciable amounts of both a weak acid and its conjugate base

• buffer is able to resist changes in pH

• capacity of buffer depends on ratio of conjugate base to weak acid and on the total absolute concentrations of the conjugate acid-base pair

Experiment #5 How to prepare a buffer?

• add weak acid and then add its conjugate base to get the correct mole ratio

• add weak acid and then add a strong base causing some of the weak acid to form its conjugate base

experiment #5 to prepare a solution by the reaction of a weak base with a strong acid, which reactant must be limiting?

• strong acid must be limiting

Experiment #5 (Buffer Capacity): What is the equation for buffer capacity.

Buffer Capacity = ∆V added / ∆pH Solution

Experiment #5 (Buffer Capacity): Objectives

Create solutions of a desired pH. Making two buffers, one with strong base + weak acid and the other with weak acid + conjugate base. Titrate both buffers with strong acids and bases. Determine capacity of each buffer. Determine how capacity changes with pH.

Experiment #5 (Buffer Capacity): What are the components we are using?

0.50M acetic acid and 0.50M sodium acetate, and 0.50 Sodium Hydroxide. We will make 40mL of the buffer.

Experiment #5 (Buffer Capacity): Is it important to use exactly 40mL of the buffer?

It is important to get as close to 40mL of buffer. This is because buffers are stronger when there's an overall greater quantity of buffer present.

Experiment #5 (Buffer Capacity): What are the titrants being used?

0.50M NaOH and 0.50M HCl

Experiment #5 (Buffer Capacity): How many times are we titrating?

We are titrating each buffer twice, once for each titrant.

Experiment #5 (Buffer Capacity): What pH solution calibrations should be used?

For two point calibration, always use pH 7. For acidic solutions, calibrate using pH 4 as well. For basic solutions, calibrate using pH 10 as well.

Experiment 5 At what pH do the buffers have maximum capacity? What is the significance of this pH?

• Both the buffers have a maximum capacity around pH= 5. This is significant because it is near the pKa of the weak acid, acetic acid. Acetic acid has a pKa of 4.76. This aligns with the concept that buffers are at their highest capacity when they are within one unit of the acid’s pKa value.

Experiment 5 common errors

• not swirling the solutions after titrating either strong base or strong acid enough causing the pH meter to give an inaccurate reading

Experiment #6 (Titration): What is a polyprotic acid?

Acid with multiple H+ ions to donate.

Experiment 6 Is it necessary to use exactly 40 ml of acid? Would the results change if you used a different volume of acid?

• no it is not necessary to use exactly 40 ml of acid

• since this is a titration with a strong base, moles will be calculated

• The mole ratio stays the same even though the volume increases

Experiment 6 Could you complete this experiment if the volume of acid did not completely cover the pH electrode?

• no you couldn’t because the initial pH reading would be inaccurate and that would mess up the rest of the readings

Experiment #6 (Titration): For the titration of a weak base with a strong acid, four regions can be distinguished.

1. A dilute solution of a weak base is prior to adding titrant.

2. A buffer solution is until the equivalence point is reached.

3. At the equivalence point, we have a mixture of conjugate acid of a weak base.

4. Beyond equivalence point, we have a mixture of strong acid and a weak acid.

Experiment #6 (Titration): Objectives

Understand the titration of a weak monoprotic acid using a strong base. Understand the four regions of the titration curve: prior to adding titrant, until before equivalence point, at equivalence point, and beyond equivalence point. Determine equivalence points using gran plots. Understand the titration curve of a polyprotic acid.

Experiment #6 (Titration): How to titrate?

For the first 1mL, titrate in 0.2-0.3mL aliquots. There are three ranges of aliquots: 1-2mL, 0.1-0.2mL, and 0.05mL aliquots. Switch between the three when you observe large pH jumps that are greater than 0.3.

Experiment #6 (Titration): What are the components of this experiment?

0.1M acetic Acid and 0.35M NaOH. 0.1M H3PO4.

Experiment #6 (Titration): what is the graph of pH vs mL base added.

The titration curve.

Experiment #6 (Titration): How many equivalence points are observed for acetic and phosphoric acid?

acetic acid will have one equivalence point. phosphoric acid will should have 3 equivalence points, but we observe 2 of them.

Experiment #6 (Titration): Why is it hard to find the third equivalence point for phosphoric acid.

It will require an immensely large amount of NaOH- since its Ka3 is so small.

Experiment #6 (Titration): What is the modified gran plot?

∆V/∆pH vs. average volume of OH- added.