answer

watching

views

11 Dec 2019

What is the difference between a primary and secondary coulemetric titration? Give an example of each.
B) A 0.1916 gram sample of an organic acid was reacted with the hydroxide ion produced coulometrically by the breakdown of water H2O + 2 e^-1 <--> 2 OH^-1 + H2(g) requiring 9 min and 24 seconds at a constant current of 0.333 amps to reach the indicator end point. If the weak acid is acetic acid, what % by weight is found in the unknown?

answer

watching

121

views

For unlimited access to Homework Help, a Homework+ subscription is required.

You have 0 free answers left.
Get unlimited access to 3.8 million step-by-step answers.

Get unlimited access

Already have an account? Log in

Hubert KochLv2

13 Dec 2019

Unlock all answers

Get 1 free homework help answer.

Already have an account? Log in

Join us on Discord

Chemistry Study Group

Join now

Related questions

Exp 25. pH Measurements- Buffers and Their Properties

I need help with finding kPa in the calculations part C for 1,2, and 3 Thanks!

Experiment 25 pH Measurements- Buffers and Their Properties ne of the more important propertics of an aqucous solution is its concentration of hydrogen ion. The H ion (or, more precisely, the HO ion) has a great effect on the solubility of many inorganic and organic species, on the nature of complex mctallic cations found in solutions, and on the rales of many chemical reac- tions. It is important that we know how to measure the concentration of hydrogen ion and understand its effect on solution properties. For convenience, the concentration of H' ion is frequendly expressed as the pH of the solution rather than as molarity. The pH of a solution is defined by the following equation: pH =-log where the logarithm is taken to the base 10. If [H] is I x 10-*moles per liter, the pH of the solution is, by the equation, 4. If the [H'1 is 5 x 10 M, the pH is 1.3, Basic solutions can also be described in terms of pH. In aqucous solutions the following equilibrium rela- tion will always be obeyed: In distilled water url equals [ OHÂ·], s, by Equation 2, [HÂ·l must be l x 10" M. Therefore, the pH of distilled water is 7. Solutions in which [H.] > 10é¨are sad to be acidic and will have a pH 7. A solution with a pH of 10 will have a [H"] of 1 x 10 0 M and a (OH] of I x 10M We measure the pH of a solution experimentally in two ways. In the first of these we use a chemical called an indicator, which is sensitive to pH. These substances have colors that change over a relatively short pH range (about two pH units) and can, when properly chosen, be used to delermine roughly the pH of a solu tion. Two very common indicators are litmus, usually used on paper, and phenolphthalein, the most common indicator in acid-base titrations. Litmus changes from red to blue as the pH of a solution goes from about 6 to about 8. Phenolphthalein changes from colorless to red as the pH goes from 8 to 10. A given indicator is useful for determining pH only in the region in which it changes color. Indicators are available for measure- ment of pH in all the important ranges of acidity and basicity. By matching the color of a suitable indicator in a solution of known pH with that in an unknown solution, one can determine the pH of the unknown to within about 0.3 pH units. The other method for finding pH is with a device called a pH meter. In this device two electrodes, one of which is sensitive to [H']. are immersed in a solution. The potential between the two electrodes is related to the pH. The pH meter is designed so that the scale will directly furnish the pH of the solution. A pH meter gives much more precise measurement of pH than does a typical indicator and is ordinarily used when an accurate determination of pH is needed Some acids and bases undergo substantial ionization in water, and are called strong because of their essentially complete ionization in reasonably dilute solutions. Other acids and bases, because of incomplete ionization (often far less than even 1% in 0.1 M solution), are called weak. Hydrochloric acid, HC, and sodium hydroxide, NaOH, are typical examples of a strong acid and a strong base. Acetic acid, HC,H,0, and ammonia, NH, are classic examples of a weak acid and a weak base. A weak acid will ionize according to the Law of Chemical Equilibrium: HBlaq)--H'(aq) + B-(aq)

I need help for my lab report please. It's only a-e Questions for study that I need help with, so any help or advice will be greatly appreciated! Underneath is more info on the lab if needed. Thank you!

Question for Study

1. A student performed this lab exercise, generating the following data:

Mass of H5IO6 used: 1.2320 g

Volume of thiosulfate used in blank titration: 47.93 mL

Volume of thiosulfate used in the unknown titration: 39.23 mL

a. What is the concentration of the sodium thiosulfate solution?

b. How many mmol of I3 â were present in the blank titration?

c. How many mmol of I3 â were present in the unknown titration?

d. How many mmol of glycol were present in the unknown titration?

e. What mass of ethylene glycol should be reported?

Introduction

The chemical analysis of organic compounds often presents major experimental difficulties. An accurate analysis depends upon the sample undergoing a single, well-defined reaction which quickly goes to completion. Often, organic compounds undergo several different reactions at the same time, and reaction rates are sometimes quite slow.

A functional-group analysis is one in which a particular grouping of atoms in a compound is made to react in a specific well-defined way with a particular reagent.1 In this experiment use is made of the reaction between an oxidizing reagent and an organic compound with the functional group COH-COH.

Analysis of Ethylene Glycol

Ethylene glycol reacts with periodate ion to form formaldehyde and iodate:

+ ð¼ð4 â â 2 ð»2ð¶ð +ð¼ð3 â + ð»2ð (8-1)

This oxidation-reduction reaction will proceed to completion with negligible side reactions under the appropriate set of conditions. The reaction is slow, but increases with increasing concentration of periodate. Here a known, fixed amount of periodate in excess of that needed to completely react with the glycol is added to the unknown. After the glycol has reacted, an excess of iodide is added to the solution. The iodide reacts with unreacted periodate according to:

ð¼ð4 â + 11 ð¼ â + 8 ð» + â 4 ð¼3 â +4 ð»2ð (8-2)

The iodide also reacts with the iodate produced in reaction 8-1 according to:

ð¼ð3 â + 8 ð¼ â + 6 ð» + â 3 ð¼3 â + 3 ð»2ð (8-3)

The triiodide produced in these reactions is then determined by titration with thiosulfate:

ð¼3 â +2 ð2ð3 2â â 3 ð¼ â + ð4ð6 2â (8-4)

To find the glycol content, a blank solution containing the same fixed amount of periodate but no glycol is treated similarly. The difference in the number of moles of ð¼3 â produced in the blank solution, and that produced in the sample solution, is equal to the number of moles of glycol in the sample solution.

Procedure

1. Prepare a solution from your unknown: a. Your unknown may have recently been refrigerated â make sure it is at room temperature before use. b. Transfer the vial contents quantitatively to a 250 mL volumetric flask. c. Dilute to volume with water and mix well.

2. Prepare a Periodic Acid standard solution (pronounced âpurr-eye-oh-dic acidâ): a. Weigh out 1.16-1.24 g of H5IO6 on an analytical balance. Periodic acid is very corrosive â to your skin and to metals. Do not spill any on the analytical balances. b. Quantitatively transfer the solid to a 100 mL volumetric flask, and dilute to volume with water.

3. Prepare a sodium thiosulfate solution: a. Weigh ~11.2 g of sodium thiosulfate pentahydrate. b. Transfer the solid to a 1L Pyrex bottle and dilute to 500 mL to produce an approximately 0.09 M solution. c. Add 0.10 g Na2CO3 and swirl to mix well.

4. Prepare a ~0.5 M solution of sulfuric acid: a. In a beaker (250 or 400 mL), combine 25 mL of 3 M sulfuric acid and 125 mL of water. Mix. Remember: âAcid Into Waterâ. What glassware is appropriate to measure these volumes?

5. Prepare samples for titration: a. Pipet 20 mL of your diluted unknown and 10 mL of the periodic acid solution into a 250 mL conical flask and swirl to mix. b. Wait a minimum of 45 minutes to allow complete reaction before titrating. Do the blank titrations while you wait! c. Prepare a total of 3 samples in this way.

6. Perform the sample titrations: a. Add ~1.5 g of KI to your conical flask and swirl to dissolve. b. Add 10 mL of your 0.5 M sulfuric acid (use a graduated cylinder). c. Titrate from a dark red initial colour to the first change to yellow. d. Add 2.0 mL of thiodene indicator solution. The solution will become a dark grey/violet colour. e. Finish the titration (fractional drops) to a sharp disappearance of the violet colour. 7. Perform the blank titrations: a. Pipet 10 mL of periodic acid solution into a 250 mL conical flask. b. Add 20 mL of water with a graduated cylinder. c. Add reagents and perform the titration as in Step 6 above. Calculations Report the total mass (in mg) of ethylene glycol in your unknown.

Experiment 1: Standardize an NaOH Solution Using Benzoic Acid as Primary Standard

Part 1: Prepare the NaOH Solution

Take a 250 mL beaker from the Containers shelf and place it on the workbench.

Take sodium hydroxide from the Materials shelf and add 0.800 g to the beaker.

Take water from the Materials shelf and add 200 mL to the beaker.

Calculate the expected concentration of the sodium hydroxide solution. The molar mass of NaOH is 39.997 g/mol.

Calculate and record the expected mass of benzoic acid required to react with 20.00 mL of a 0.100 M sodium hydroxide solution. Record the mass of benzoic acid using three significant digits

Part 2: Perform a Coarse Titration

Take a clean Erlenmeyer flask from the Containers shelf and place it on the workbench.

Take water from the Materials shelf and add 100 mL to the Erlenmeyer flask.

Add the required mass of benzoic acid calculated in step 5 of part 1 to the Erlenmeyer flask. The 100 mL of water should be enough to dissolve the solid benzoic acid.

Take phenolphthalein from the Materials shelf and add 2 drops to the Erlenmeyer flask.

Take a pH meter from the Instruments shelf and place it into the flask. Record the initial pH of the solution.

Take a burette from the Containers shelf and place it on the workbench.

Take the prepared sodium hydroxide solution from the workbench and add 50 mL to the burette. Pass the mouse cursor over the burette and a gray tool tip will briefly display the total volume in the burette as well as the volume dispensed. Record these initial volumes.

Move the Erlenmeyer flask onto the base of the burette.

Perform a coarse titration, adding large increments of titrant (~ 2 mL) by pressing and holding the stopcock at the bottom of the burette for 1-2 seconds. Pause after each dispensation. Record the volume dispensed.

Check if the end point has passed. When the reaction reaches the end point, the solution changes color. Also, as sodium hydroxide is added to the benzoic acid solution, the pH increases.

Stop once you reach the end point. Record the volume dispensed and pH at which this occurs. Record both the last dispensed volume where the solution was colorless (right before the end point) and the first dispensed volume where the solution changed color. You will use the volume dispensed before the end point in the fine titration.

Discard the Erlenmeyer flask by dragging it to the recycling bin beneath the workbench.

Part 3: Perform Fine Titrations

Prepare an Erlenmeyer flask as described in steps 1 - 5 in part 2 of experiment 1.

Take sodium hydroxide solution from the workbench and refill the burette to a total volume of 50 mL. Record this initial volume.

Move the Erlenmeyer flask onto the base of the burette.

Click and hold the stopcock of the burette to quickly add the before the end point volume of sodium hydroxide determined in the coarse titration.

Add titrant in small increments, down to one drop at a time. This can be tedious, but if you click and hold you might miss the exact end point of the titration.

When the solution changes color, stop adding titrant. Record the pH and volume dispensed.

Clear your workstation by dragging the Erlenmeyer flask to the recycling bin.

Repeat the fine titration two more times, and record the results.

Calculate the concentration of the sodium hydroxide solution. The rest of the sodium hydroxide solution can now be used in further lab work as a secondary standard with a reliably known concentration equal to the average of the three titrations.

Experiment 2: Use the Standardized NaOH Solution to Determine the Concentration of an Acid

Part 1: Perform a Coarse Titration

Take an Erlenmeyer flask from the Containers shelf and place it on the workbench.

Take acetic acid #1 (unknown concentration) from the Materials shelf and add 25 mL to the flask.

Take phenolphthalein from the Materials shelf and add two drops to the flask.

Take a pH meter from the Instruments shelf and place it into the Erlenmeyer flask. Record the pH.

Take the NaOH solution from the workbench and refill the burette to a volume of 50 mL. Pass the mouse cursor over the burette and a gray tool tip will briefly display the total volume in the burette as well as the volume dispensed. Record these volumes.

Move the Erlenmeyer flask onto the base of the burette.

Perform a coarse titration, adding large increments of titrant (~ 2 mL) from the burette by pressing and holding the stopcock at the bottom of the burette. Pause after each dispensation. Record the volume dispensed.

Check if the end point has passed. When the reaction reaches the end point, the solution changes color. Also, as more sodium hydroxide is added to the acetic acid solution, the pH increases.

Stop once you reach the end point. Record the dispensed volume and pH at which this occurs. Record both the last dispensed volume where the solution was colorless (right before the end point) and the first dispensed volume where the solution changed color. You will use the volume dispensed right before the end point in the fine titration.

Clear your station by dragging the Erlenmeyer flask to the recycling bin.

Part 2: Perform Fine Titrations

Set up the titration as described in steps 1 â 6 in part 1 of experiment 2.

Click and hold the stopcock of the burette to quickly add the before the end point volume of sodium hydroxide determined in the coarse titration.

Add sodium hydroxide in small increments, down to one drop at a time. This can be tedious, but if you click and hold you might miss the exact end point of the titration.

When the solution changes color, stop adding titrant. Record both the pH and volume dispensed.

Clear your station by dragging the Erlenmeyer flask to the recycling bin.

Repeat the fine titration twice more

Weekly leaderboard

Home

Homework Help 3,900,000

Chemistry 830,000

Unlock all answers

Related textbook solutions

Basic Chemistry

Principles of Chemistry Molecular Approach

Principles of Chemistry Molecular Approach

Chemistry: Structure and Properties

Chemistry: A Molecular Approach

Chemistry: A Molecular Approach

Principles of Chemistry: A Molecular Approach

Chemistry: The Central Science

Related questions

Weekly leaderboard