Etiket Arşivleri: Titration

Acid-Base Titrations

Titration is an analytical method used to determine the exact amount of a substance by reacting that substance with a known amount of another substance. The completed reaction of a titration is usually indicated by a color change or an electrical measurement. An acid/base neutralization reaction will yield salt and water. In an acid-base titration, the neutralization reaction between the acid and base can be measured with either a color indicator or a pH meter.

Acid + Base  Salt + Water

In this experiment, a phenolphthalein color indicator will be used. Phenolphthalein is colorless in acidic solutions and pink in basic solutions. Phenolphthalein is also used in forensic crime scene analysis to detect the presence of blood, Kastle-Meyer test. In the Kastle-Meyer test, hemoglobin catalyzes the oxidation of the colorless form of phenolphthalein to its bright pink form. Four lab periods assigned for this experiment. In part I you will prepare an acid (HCl) solution and a base (NaOH) solution. These solutions will be used for all four periods so it is important to keep these solutions. These solutions will be titrated against each other to obtain a base/acid ratio. In part II you will find the normality of the base solution by titrating it against a solid acid standard. The normality of the acid can be calculated from the normality of the base and the base/acid ratio from part I. In part III the base will be titrated against an unknown acid to find the equivalent weight of the acid. In part IV the equivalent weight of an unknown base will be determined by reacting the unknown base with an excess of HCl and “back-titrating” the left-over acid with NaOH.

Equipment and Reagents (Part I)

6 NHCl 1 Liter plastic bottle 2 beakers (50 mL)
6N NaOH 2 burets 250 mL Erlenmeyer Flask
500 mL Florence Flask Iron stand wash bottle
Distilled water buret clamp phenolphthalein indicator
Stopper(or parafilm) 2 x 50 mL graduated cylinder

Procedure (Part I)

1. Rinse a clean 500 mL Florence flask with a small portion of DI water. Place about 16-17 mL of 6 M or 6 N HCl into the flask and dilute to 500 mL with distilled water. The 500 mL is approximated by bringing the level of the solution up to the point of constriction of the neck of the flask. Stopper the flask and shake to mix. The solution should be approximately 0.2 N HCl. Label the flask with tape.

2. Rinse a clean 1 L plastic bottle with distilled water. Place about 32-34 mL of 6 M or 6 N NaOH into the bottle and dilute to 1 liter with distilled water. Place the cap on the bottle and shake to mix. The solution should be approximately 0.2 N NaOH. Label the bottle with tape.

3. Obtain 2 burets from the stockroom and clamp them onto the ring stand using the buret clamp. Label the buret as acid or base. Label the 50 mL beakers as acid or base. These beakers will be used to transfer the solutions into the burets. Rinse each buret with about 5 mL of DI water and with about 3 x 5 mL of the solution to be used. Fill each buret with the correct solution and flush all of the air bubbles out of the buret tip.

4. Read the initial level of each buret to the nearest 0.02 mL and record this in your notebook. The proper reading is taken from the bottom of the meniscus (see Figure 1 below). If the initial reading is at exactly at zero, then report 0.00 mL.

5. Allow about 25 mL of the acid to run into an Erlenmeyer flask from the acid buret. Record the initial and final readings of this transfer. Calculate the volume of acid transferred by subtracting the final volume reading by the initial volume reading. Your final answer should be to the hundredth place.

6. Add about 20 mL of distilled water into the flask and add 2-3 drops of phenolphthalein indicator. The flask should remain colorless at this point.

7. Record the initial volume of base. Slowly add NaOH from the base buret into the flask with constant swirling. Continue adding base until a very faint color remains. If the color is too bright, add a few drops of acid so that the solution becomes colorless. Add base again to reach the faint end-point. Repeat this process until a faint pink end-point is reached. Record the final volume of base and the initial and final volume of extra acid added to this flask

8. Calculate the total final volume of acid and final volume of base added.

9. From these values, calculate the base to acid ratio:

Source: https://lahc.edu/classes/chemistry/arias/Exp%207%20-%20AcidBaseF11.pdf

Determination of Vitamin C Concentration by Titration

Determination of Vitamin C Concentration by Titration

(Redox Titration Using Iodine Solution)

Safety

Lab coats, safety glasses and enclosed footwear must
be worn at all times in the laboratory.

Introduction

This method determines the vitamin C concentration in a solution by a redox titration using iodine. Vitamin C, more properly called ascorbic acid, is an essential antioxidant needed by the human body (see additional notes). As the iodine is added during the titration, the ascorbic acid is oxidised to dehydroascorbic acid, while the iodine is reduced to iodide ions.

ascorbic acid + I2 → 2 I− + dehydroascorbic acid

Due to this reaction, the iodine formed is immediately reduced to iodide as long as there is any ascorbic acid present. Once all the ascorbic acid has been oxidised, the excess iodine is free to react with the starch indicator, forming the blue-black starch-iodine complex. This is the endpoint of the titration. The method is suitable for use with vitamin C tablets, fresh or packaged fruit juices and solid fruits and vegetables.

NB: This method is more straight forward than the alternative method using potassium iodate, but as the potassium iodate solution is more stable than the iodine as a primary standard, the alternative method is more reliable.

Equipment Needed

burette and stand
100 mL or 200 mL volumetric flask
20 mL pipette
10 mL and 100 mL measuring cylinders
250 mL conical flasks

Solutions Needed

Iodine solution: (0.005 mol L−1). Weigh 2 g of potassium iodide into a 100 mL beaker. Weigh 1.3 g of iodine and add it into the same beaker. Add a few mL of distilled water and swirl for a few minutes until iodine is dissolved. Transfer iodine solution to a 1 L volumetric flask, making sure to rinse all traces of solution into the
volumetric flask using distilled water. Make the solution
up to the 1 L mark with distilled water.
Starch indicator solution: (0.5%). Weigh 0.25 g of soluble starch and add it to 50 mL of near boiling water in a 100 mL conical flask. Stir to dissolve and cool before using.

Source: https://www.canterbury.ac.nz/media/documents/science-outreach/vitaminc_iodine.pdf

Determination of Dissolved Oxygen By Winkler Titration

Lab 1:

DETERMINATION OF DISSOLVED OXYGEN BY WINKLER TITRATION

1. Background

Knowledge of the dissolved oxygen (O2) concentration in seawater is often necessary in environmental and marine science. It may be used by physical oceanographers to study water masses in the ocean. It provides the marine biologist with a means of measuring
primary production – particularly in laboratory cultures. For the marine chemist, it provides a measure of the redox potential of the water column. The concentration of dissolved oxygen can be readily, and accurately, measured by the method originally developed by Winkler in 1888 (Ber. Deutsch Chem. Gos., 21, 2843). Dissolved oxygen can also be determined with precision using oxygen sensitive electrodes; such electrodes require frequent standardization with waters containing known concentrations of oxygen. They are particularly useful in polluted waters where
oxygen concentrations may be quite high. In addition, their sensitivity can be exploited in environments with rapidly-changing oxygen concentrations. However, electrodes are less reliable when oxygen concentrations are very low. For these reasons, the Winkler titration is often employed for accurate determination of oxygen concentrations in aqueous samples.


Source: https://ocw.mit.edu/courses/earth-atmospheric-and-planetary-sciences/12-097-chemical-investigations-of-boston-harbor-january-iap-2006/labs/dissolved_oxygen.pdf

Principles of Titration and Errors ( Dr. A. Amsavel )

Principles of Titration and errors

By Dr. A. Amsavel

Introduction

Volumetric analysis
 Simple and easy
 Fast and can be done on site
 Less expensive
 Estimation of content or Assay
 Precise and accurate
 Depends on method and specificity

Requirements of a Titration Reaction

Reaction must complete by 99.9 % so that < 0.1 % (or 1 ppt) remains unreacted
Rxn must be rapid
Titration needs to be performed in a reasonable time period
The stoichiometry must be well defined, and known
It can be predicted from equilibrium constants
A method must be available to determine the equivalence point

Types of Titration

1) Precipitation
– A(aq) + B(aq) = AB(s)
2) Acid-Base rxn
– H+ + OH ̄ = H2O (strong acids or bases)
– HA + OH ̄ = H2O + A ̄ (weak acids)
– A ̄ + H+ = H2O + HA (weak bases)
3) Complexation rxn
– Zn2+ + 4NH3 = Zn(NH3)42+
4) Redox rxn (oxidation-reduction)
– Fe2+ + Ce4+ = Fe3+ + Ce3+

Standards

• Measurements are made with reference to standards
– The accuracy of a result is only as good as the quality and accuracy of the standards used
– A standard is a reference material whose purity and composition are well known and well defined
• Primary Standards
– Used as titrants or used to standardize titrants
– Requirements
• Usually solid to make it easier to weigh
• Easy to obtain, purify and store, and easy to dry
• Inert in the atmosphere
• High formula weight so that it can be weighed with high precision

Endpoint Detection

It is critical, to know the completion of reaction / determination
1) Visual indicators:
• Observe a colour change or precipitation at the endpoint.
– Rxn progress checked by addition of external or self indicator
2) Photometry:
• Use an instrument to follow the colour change or precipitation
3) Electrochemistry:
• Potentiometry – measure voltage change ( pH electrode)
• Amperometry – measure change in current between electrodes in solution
• Conductance – measure conductivity changes of solution
Later two used for coloured, turbid, end point accurate

Acid-base titration

 Neutralization titration
 Neutralization Indicators
 Indicators & mixed indicators
 Neutralization curve
 Non-aqueous titration


Lab: Oxidation – Reduction Titrations ( Permanganometry )

Potassium permanganate, KMnO4, is probably the most widely used of all volumetric oxidizing agents. It is a powerful oxidant and readily available at modest cost. The intense color of the permanganate ion, MnO4- , is sufficient to detect the end point in most titrations.

Depending upon reaction conditions permanganate ion is reduced to manganese in the 2+, 3+, 4+ or 6+ state.

In solutions that are 0.1 M or greater in mineral acid the common reduction product is manganese (II) ion

MnO4-+ 8H++ 5e-↔ Mn2++ 4H2O E0 = 1.51 V

This is the most widely used of the permanganate reactions.

In solutions that are weakly acidic (above pH 4) neutral, or weakly alkaline manganese dioxideis the most common reduction product

MnO4-+ 4H++ 3e-↔ MnO2(s) + 2H2O E0 = 1.70 V

Titration in which manganese dioxide is the product suffer from the disadvantage that the slightly soluble brown oxide obscures the end point; time must be allowed for the solid to settle before an excess of the permanganate can be detected. Some important volumetric analyses based on permanganate involve reduction to manganese ion according to the half reaction given below;

MnO4-+ e-↔ MnO42-   E0 = 0.56 V

This stoichiometry tends to predominate in solutions that are greater than 1 M in sodiumhydroxide. Alkaline oxidations with permanganate have proved to be most useful in the determination of organic compounds.


Source: http://users.metu.edu.tr/chem223/permanganometry.pdf

Acid – Base Titration Lab. Reports

PURPOSE:

The purpose of the titration is to determine the amount of acid it contains by measuring the number of mL of the standard NaOH need to neutralize it. The technique of titration will use to determine the concentration of solutions of acids and bases.

THEORY:

Titration is a laboratory technique designed to use the reaction of two solutions to determine the concentration of one or the other.  The titration with the HCl will be used to determine our NaOH solution concentration.

The technique of titration can be applied to different materials and different types of reactions. In this experiment, an acid-base neutralization between two solutions is used. The standard solution is 0.1 M NaOH, and the unknown solution is the mixture resulting from dissolving the antacid in excess “stomach acid”. This solution is still acidic; The reaction is:

NaOH + HCl à H2O + NaCl

An indicator is used to show when the solution has been neutralized. In this case, the indicator is methyl red, an organic compound which is yellow if it is dissolved in a basic solution, and red in an acidic one. As the excess stomach acid is neutralized by the NaOH, the solution changes from red to yellow; the exact neutralization point is orange.

At the neutralization point (or endpoint), the number of moles of OH-1 added exactly equals the number of moles of H+1 that remained after the antacid tablet “worked”. This number of moles of OH-1 can be found by multiplying the number of liters delivered by the burette times the molarity of the NaOH:

liters NaOH sol’n x moles NaOH = moles NaOH 1 liter sol’n.

The number of moles of HCl remaining after treatment with the antacid tablet is equal to the number of moles of NaOH calculated above. The total number of moles of HCl added to the flask can be found by multiplying the total volume of HCl used (in liters) by the Molarity of the HCl solution. The number of moles of HCl neutralized by the stomach acid is just the difference between these two values.

moles HCl neutralized by tablet = total moles HCl in flask – moles HCl neutralized by NaOH

The effectiveness of the antacid may be expressed in terms of mL of stomach acid neutralized, or moles of HCl neutralized, and may be calculated per tablet or per gram of antacid. See calculation page.

Burette Use and Titration Technique:

Typically, a special piece of glassware is used to measure out one or both of the solutions used in a titration. It is called a burette, and quickly and accurately measures the volume of the solutions delivered. A diagram of a burette, and instructions on how to use it are given below.

Before use, the burette must be rinsed with the solution it is to contain. Close the stopcock and use a small beaker to pour about 10 mL of solution into the burette. Tip the burette sideways and rotate it until all of the inside surfaces are coated with solution. Then open the stopcock and allow the remaining solution to run out. Again close the stopcock, and pour enough solution into the burette to fill it above the “0” mark. With the burette clamped in a vertical position, open the stopcock and allow the liquid level to drop to “0” or below. Check the burette tip. It should not contain air bubbles! If it does, see your instructor. Adjust the burette so that the liquid surface is at eye level, and take the initial burette reading as shown below:

M Na 2 CO 3 , 0.1 M HCl , 0.1 M NaOH and methyl orange indicator. Two flasks, burette, burette clamp, pipet, Bunsen burner

PROCEDURE:

1.  Determination of Concentration of HCl Solution:

Firstly, 10 ml of 0.1 M of Na 2 CO 3 was place in each of two flasks using a pipet. Then two drops of methyl orange indicator was added to each of the flasks and the solution was mixed with a gentle shaking. The burette was filled until full with HCl. One of the Erlenmeyer flasks was placed under the tip of the burette. The HCl was run from the burette into the Na 2 CO 3 solution until the end point is reached. Then the mixture was boiled for a few minutes. After the solution was cooled and added a few drops of HCl until the end point is observed again. When the end point is obtained, the volume of HCl added was recorded.

The solution of Na 2 CO 3 in the second flask was titrated following exactly the same procedure.

The molarity of HCl was calculated.

2. Quantitative Determination of NaOH:

Firstly, 10 ml of 0.1 M NaOH was placed in each of two flasks using a pipet. Then two drops of methyl orange indicator was added to each of the flasks and it was titrated as we did with Na 2 CO 3 . But the solution did not boil. The solution of NaOH in the second flask was titrated following exactly the same procedure.

The molarity of NaOH solution was calculated.

3. Unknown

Given an unknown solution of NaOH was titrated similarly above.

Titration of Sodium Carbonate

  • Carbonate Chemistry

  • CO2 in atmosphere and dissolved in water

  • Major global buffering system

  • Industrial sources

–limestone:  CaCO3(s) + heat à CaO(s) + CO2(g)

–trona (Na2CO3) deposits

  • Sodium Carbonate As a Base

  • Commercially important

–source or base or carbonate for industrial processes

–washing soda (automatic dishwashers)

  • Carbonate is a moderately strong base

  • Titration of Na2CO3 with HCl

  • Titration of Sodium Carbonate

  • Derivative Plots

  • Determination of Carbonate in a Sample

  • Effects of Carbonate Equilibria on Titration of Carbonate

  • Shift in Phenolphthalein Endpoint

  • Boiling to Enhance Second Endpoint

  • Effect of CO2 Absorption on Phenophthalein Endpoint

  • Reaction of CO2 with carbonate solutions

CO2(g) + H2O(l) + CO32-(aq) à 2HCO3-(aq)

–Let 1 mmol CO2 react with 1 mmol CO32-

–For titration with HCl to BCG endpoint to form H2CO3

  • Initial 1 mmol CO32- would require 2 mmol HCl

  • Resultant 2 mmol HCO3- requires 2 mmol HCl

  • Net effects

–Volume of HCl require to reach phenolphthalein end decreases

  • This effect may be avoided by excluding air from the titration system.

–Volume required to reach bromcresol green endpoint is not affected.

  • Recommendation

–Use BCG endpoint to calculate amount of carbonate in your sample

  • Effect of Absorbed CO2 on Titration of Na2CO3 with HCl

  • Boiling to Enhance Visual Endpoint in Titration of Na2CO3 with HCl

  • Summary: Titration of Sodium Carbonate

  • Titration of Na2CO3 with strong acid yields two equivalence points

  • At phenolphthalein end point

–Na2CO3 à NaHCO3

  • At bromcresol green (or methyl red) end point

–All carbonate is converted to H2CO3

  • pH of carbonate solutions may be unstable due to exchange of CO2 with atmosphere.


Laboratory‎ > ‎Acid – Base Titration

An acids-base titration is used to determine the unknown concentration of an acid or base by monitoring its neutralization reaction with a standard of known concentration in the presence of an indicator. For example, a base of unknown concentration may be titrated with an acid standard of known concentration or the reverse. The endpoint is the point of the acid-base titration in which just enough of the standard solution has been added to completely neutralize the solution of unknown concentration. Indicators are used to signal this point of the titration, and are generally selected according to the type of titration (strength of acid and based used in the titration). When the endpoint of the titration has been reached, one drop of standard solution past the endpoint will result in a color change of the indicator.

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