CHEM 318 Chapter Notes - Chapter 1: Gravimetry, Detection Limit, Reagent

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roseweasel705

21 May 2017

School

University of Victoria

Department

Chemistry

Course

CHEM 318

Professor

Hore Dennis

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Document Summary

Qualitative analysis: identifying the components of a material. Quantitative analysis: measuring the total amount or concentration of speci c compounds in the sample. Standard deviation: measure of the variability of results from a sample. Systemic error: an error in the measured value that introduces a bias to all values due to an error within the instrument. This is a repeatable error that a ects the accuracy of the value. Random error: an error introduced by limitations of our measuring devices and technique. Primary standard - reagent chemical of high purity (>99%) and known formula can be weighed out to make a solution of known concentration. Secondary standard reagent chemical of low purity or uncertain formula make a solution and analyze it to get concentration. Standard solution a solution with an accurately known concentration made from a reagent chemical, not from an unknown sample. Detection limit: is the smallest quantity of analyte that is signi cantly di erent from the blank.

Related Questions

Internal Standards

Some analytical methods are subject to one or more common sources of variability in the instrument response. One method aimed at overcoming this variability is the internal standard method. In this method, a known amount of a compound similar to the analyte is added to each standard and sample. Since we know the concentration of the internal standard (IS), we can use the IS response as an indicator of variability. In this type of calibration curve, the quantity plotted on the y-axis is not simply the instrument signal for the analyte, but rather the ratio of the analyte signal to the IS signal.

The IS method is commonly used in gas chromatography, because of the variability inherent in the small volumes injected into the instrument. In the forensic lab, blood alcohol measurements are one of the most common analyses performed. This dry lab example will use blood alcohol data.

You will use ethanol standards that you have purchased. You have 5 of these, with concentrations of 0.050, 0.080, 0.100, 0.200, and 0.300 %w/v of ethanol. You also have a blank, which doesnât contain any ethanol.

You will use n-propanol as the internal standard. You have made up a 0.0300% (v/v) solution of propanol.

You have one sample to run.

You add 2 mL of each of the standards and the sample to GC vials. Then you add 0.450 mL of the internal standard n-propanol to each vial. You get the data below.

Ethanol concentration, % (w/v)	Peak area for ethanol (pA*sec)	Peak area for n-propanol pA*sec)	Ratio of ethanol area to propanol area

First, make a calibration curve using only the area of the ethanol peak and the concentrations of the standards. As in Part 1 a and b, prepare and print a calibration curve including the linear least squares equation of the line equation and the R² value, and print a copy for your report. Also, calculate the concentration of your sample and the uncertainty (95% confidence interval), as well as the uncertainty of the unknown determination on a percentage basis, and include these values in your report. .

Next, make a calibration curve based on the ratio of ethanol area to propanol area and the concentrations of the standards. As in the previous question, prepare and print a calibration curve including the linear least squares equation of the line equation and the R² value, and print a copy for your report. Also, calculate the concentration of your sample and the uncertainty (95% confidence interval), as well as the uncertainty of the unknown determination on a percentage basis, and include these values in your report. .

At a blood alcohol concentration of 0.08% (w/v) or more a person is considered impaired. Briefly discuss whether the calculated results for the sample in e and f above support criminal charges of impairment.

Report Requirements:

Spreadsheet with graphs and calculations. Prepare separate sheets in one excel file for each of the following:

Data Set 1

Data Set 2

Calibration with external standard

Standard addition with one standard (optional, could submit hand or type written calculation only)

Standard addition with multiple additions

Internal standard peak areas

Internal standard peak area ratios

Print-outs of graphs

Data Set 1

Data Set 2

Calibration with external standard

Standard addition with multiple additions

Internal standard peak areas

Internal standard peak area ratios

Results

Fill in the table below with results from the spreadsheet. Remember to include units.


	Concentration of unknown	Uncertainty of unknown	Uncertainty of unknown in terms of %

	Concentration of unknown	Uncertainty of unknown	Uncertainty of unknown in terms of %

2d	Concentration of unknown	Uncertainty of unknown	Uncertainty of unknown in terms of %

	Concentration of unknown	Uncertainty of unknown	Uncertainty of unknown in terms of %

	Concentration of unknown	Uncertainty of unknown	Uncertainty of unknown in terms of %

Short answers/additional tables

Section 1c:

Section 1d:

Section 1e2:

Section 2a Table:

Section 3b Table:

Section 3b4 calculation

Section 3b5

Section 4d Table with calculated ratios

Section 4g

INTRODUCTION

Many transition metal cations can have more than one possible charge (Worksheet 1). Differently charged species of the same element can have different chemical and physical properties. Knowing the charge can allow separation of the different ions based on their properties. For example, Cu²⁺ is blue and stable in solution, while Cu⁺ is colorless and reacts with oxygen. Hemoglobin contains iron as Fe²⁺, but it is inactive at binding oxygen when the iron gets oxidized to Fe³⁺. There are many similar situations in which it is important to always properly indicate charges of ions.

This laboratory experiment uses the difference in chemical properties of Fe²⁺, iron(II), and Fe³⁺, iron(III), to quantify the amount of each present in a solution of unknown concentrations of each. You will first determine the amount of iron(II) in your sample by titrating with a reagent that does not react with iron(III). You will then titrate a second sample after converting all the iron(III) to iron(II) in order to determine the total iron present. The percent composition of Fe²⁺ in the unknown mixture will be calculated.

BACKGROUND

The concentrations of both iron(II) and iron(III) ions in a solution can be determined by titration with potassium permanganate. The reaction is an oxidation-reduction reaction, or redox, for short. A brief introduction to oxidation-reduction reactions is given in the background of the âAcids and Basesâ experiment. The permanganate ion effectively oxidizes iron(II) to iron(III) in an acidic solution while being reduced to manganese(II) ion. The balanced net ionic reaction is:

5 Fe²⁺(aq) + MnO4^â(aq) + 8 H⁺(aq) â 5 Fe³⁺(aq) + Mn²⁺(aq) + 4 H₂O(l)

Iron(III) is not oxidized by the permanganate ion, and thus does not react. Thus in order to determine the concentration of iron(III) ions, they must first be reduced to iron(II) ions. Zinc metal is used as the reducing agent, also under acidic conditions.

Zn(s) + 2 Fe³⁺(aq) â 2 Fe²⁺(aq) + Zn²⁺(aq) (reduction of the Fe3+)

Zn(s) + 2 H⁺(aq) â Zn²⁺(aq) + H₂(g) (a side reaction that oxidizes excess Zn)

The Zn2+ formed does not interfere with the subsequent titration because Zn2+ is not oxidized by permanganate. Notice that converting all the iron(III) to iron(II) means that the total iron concentration ([Fe2+] + [Fe3+]) is titrated, not just the iron(II) originally present.

Techniques

This is a lab will utilize several of the techniques introduced in the experiment âLaboratory Techniques and Measurements,â the procedures for which are given in Appendix B. Poor laboratory technique and misuse of the glassware will result in poor data results, which will affect the grade you receive on the laboratory report. In order to use your time efficiently in lab, review how to properly:

rinse glassware prepare a standard solution use a volumetric pipette

weigh by difference set up and use a burette use a graduated cylinder

heat samples on a hot plate

Analysis of Iron by Oxidation-Reduction Titration

This lab uses the common quantitative technique of titration. The only difference between the titration you will be doing in this lab and that done in the Molar Mass of a Known Acid lab is the type of reaction used: this lab will use a redox reaction in place of the acid-base reaction performed in the previous experiment.

In order to determine the concentrations of iron(II) and iron(III) ions in a solution containing both ions, two portions of the same sample must be titrated separately. This first titration will determine only the iron(II) concentration. A second sample will be treated with zinc metal in order to reduce the iron(III) to iron(II). The resulting solution will be titrated with potassium permanganate. The result of this titration will give the total iron concentration in the sample. The difference between the first and second titration is the equal to the iron(III) concentration.

In this titration, potassium permanganate serves as its own indicator. The intense purple color of the permanganate solution becomes colorless when the permanganate is reduced to manganese(II). Thus, as iron(II) is oxidized with the addition of the permanganate solution from the burette it will impart a purple color that will fade. However, once there is no iron(II) left in solution a faint purple color will remain, marking the endpoint of the titration. This color can be slightly altered by the presence of iron(III), which is yellow in solution. Remember that iron(III) is generated by the reaction between iron(II) with permanganate, increasing the yellow appearance of the solution and which can interfere with the detection of the endpoint. Phosphoric acid is added to help minimize this effect.

pre lab questions:

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.

CHEM 318 Chapter Notes - Chapter 1: Gravimetry, Detection Limit, Reagent

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