3.1.2 Standardization Standardization of Sodium Hydroxide Standard Solution Reference: Standard Solution of Sodium Hydroxide (936.16) Official Methods of Analysis. 1990. Association of Official Analytical Chemists. 15th Edition. Scope: This method is applicable for the preparation and standardization of standard sodium hydroxide solution. Basic Principle: A basic solution is titrated with a standardized acidic solution to determine normality. Equipment: Buret, 50 mL, graduated to 0.1 mL Illuminated magnetic stirrer Volumetric pipet, 40 mL, class A Analytical balance, sensitive to 0.1 mg Container of alkali resistant glass pH meter with glass electrode (alternate to using phenolphthalein indicator) Reagents: Distilled water, carbon dioxide (CO2) free fr ee prepared either 1) by boiling for 20 min and cooling with soda-lime protection or 2) by bubbling air, freed from CO2 by passing through tower of soda lime, through water for 12 hr. Sodium h ydroxide solution, to 1 part reagent grade NaOH add 1 part distilled, carbon dioxide-free water by weight. Acid A cid potassium phthalate, NIST SRM for Acidimetry 84, dry for 2 hr at 120oC and cool in desiccator. Buffer solution, pH 8.6, 12.00 mL 0.2 N NaOH added to 50 mL 0.2M boric acid/potassium chloride solution made as follows: Boric acid-potassium chloride solution - dry boric acid (H3BO3) to constant weight in desiccator over CaCl2. Dry potassium chloride (KCl) 2 days in oven at 115 to 120oC. Dissolve 12.405 g H3BO3 and 14.912 g KCl in water and dilute to 1 L. Phenolphthalein indicator, 1%, dissolve 1 g phenolphthalein in 100 m L 95% ethanol. Safety Precautions:
Alkalis can burn skin, eyes and respiratory tract severely. Wear heavy rubber gloves and face shield to protect against concentrated alkali; if spilled on skin wash with copius amounts of water. Use effective fume removal device to protect against alkali dusts or vapors. Always add sodium hydroxide pellets to water, not vice versa.
Procedure: Preparation
Add appropriate volume of NaOH solution (1 to 1) to CO2-free distil led water necessary to make 10 L of solution: mLNaOH to be diluted to 10 L Desired Normality
0.01
5.4
0.02
10.8
0.10
54.0
0.20
108.0
0.50
270.0
1.0
540.0
Standardize
1. Accurately weigh enough dried acid potassium phthalate (ca. 0.4 g) t o titrate about 40 mL of NaOH solution and transfer to 300 mL flask. 2. Add 50 mL CO2-free water, stopper flask and swirl until sample dissolves. 3. Titrate to pH 8.6 with solution being standardized, taking precautions to exclude CO2 and using as indicator either glass-electrode or phenolphthalein. If using indicator, add 3 drops phenolphthalein to a flask containing 50 mL of pH 8.6 buffer and stopper. This flask is used as the reference endpoint for a pH 8.6 titration. 4. Determine volume NaOH required to produce endpoint of blank by matching color in another flask containing 3 drops phenolphthalein and same volume (50 mL) CO2-free water. 5. Subtract volume required to titrate blank from that used to titrate the potassium acid phthalate and calculate normality. Normality should be slightly high. 6. Adjust to desired concentration, mix well, and recheck standardization. 7. Record final standardization in logbook. Calculations: Normality = g KHC8H4O4 X 1000/mL NaOH X 204.229
Adjust to desired concentration by following formula: V1 = V2 X N2 / N1 Where N2 and V2 represent normality and volume of stock solution and V1 equals volume to which stock solution should be diluted to obtain desired normality, N1.
Preparing Standard Sodium Hydroxide Solution* By Dr. Murli Dharmadhikari and Tavis Harris Note: This article has been written at the request of the industry. It is written for wine lab workers with no background in chemistry.
In a wine laboratory, analyzing wine for TA, VA and S02 involves the use of a sodium hydroxide (NaOH) reagent. Winemakers usually buy sodium hydroxide solution of a known concentration (usually 0.1 Normal). This reagent is r elatively unstable and its concentration changes over time. To ensure the accuracy of analytical results it is important to periodically check the concentration (Normality) of sodium hydroxide. If the concentration has changed then it must be readjusted to the original concentration or the new concentration (Normality) value needs to be used in calculations. Sometimes a winemaker may wish to make his/her own NaOH solution instead of buying it. Whether making a new solution or checking the normality of an old solution it i s important to know the procedure for making a standard (known concentration) solution of NaOH reagent.
In the present article the standardization procedure along with the basic concept behind the titration procedure are explained. Expressing concentration in solution
A solution consists of a solute and the solvent. Solute is the dissolved substance and solvent is the substance in which the solute is dissolved. A solute can be a solid or a l iquid. In NaOH solution, sodium hydroxide (solid) is the solute and water (liquid) is the solvent. Note that the solute being a solid is measured in terms of weight (in grams) and the solvent water is measured in terms of volume. This is an example of expressing solution in weight per volume (w/v) basis. In a solution consisting of two liquids the concentration is expressed in a volumes per volumes basis. For example the concentration of alcohol in wine is expressed as volume per volume. A 12% alcohol wine means it contains 12 ml of alcohol per 100 ml of wine. Generally, in many solutions, the weight is given in grams and volume is given in milliliters or liters. At this point, it is important to establish the relation between the units of weight and volume. One kilogram (weight) of water at a temperature of maximum density and under normal atmospheric pressure has the volume of one liter. This means that one kilogram (weight) of water equals one liter of volume, and one gram of water by weight equals one milliliter of water by volume. Thus the units of weight (gram) and volume (ml) are similar and interchangeable. The chemist expresses the concentration of a solution i n various ways. The common expressions include Percent, Parts per million (ppm), Molar and Normal. It is important to have a clear understanding of these terms. Percent
One of the simplest forms of concentration is the percent. This simply means units per 100 units, or parts per 100 parts. The percent concentration can be used in three ways. It can be weight per weight, volume per volume or weight per volume basis. When winemakers use °Brix hydrometer to measure sugars in grape juice they are essentially measuring grams of sugar per 100 grams of juice. A juice sample of 18 °Brix means 18 grams of sugar per 100 grams of juice or commonly referred as 18%. In describing the alcohol content of a wine, percent alcohol content is expressed in terms of a volume per volume basis. In many cases, including in a laboratory, a solution is made by dissolving a solid in a li quid, usually water. In such a case the concentration is expressed in a weight per volume basis. Parts per million
When dealing with a very small amount of a substance in solution, the concentration is often expressed in terms of parts per million. A 20 ppm concentration means 20 parts of solute dissolved for every 1,000,000 parts of solution. The unit of measurement can be weight or volume. Generally the ppm concentration is used to indicate milligrams of solute per liter of solution. Molar solution
A molar solution implies concentration in terms of moles/liter. One molar (I M) solution means one mole of a substance (solute) per liter of solution. A mole means gram molecular weight or molecular weight of a substance in grams. So the molecular weight of a chemical is also its molar weight. To calculate the molecular weight one needs to add the atomic weights of all the atoms in the molecular formula unit. For example the molecule of NaOH consists of one atom each of sodium (Na), oxygen (0), and hydrogen (H). Their respective atomic weights are: Na - 23,0 - 16 and H - 1, so the molecular weight, is 23 + 16 + I = 40. Thus 40 grams of NaOH equals one mole of NaOH, and a 1 molar solution of NaOH will contain 40 grams of NaOH chemical. Normal solution/Normality
The other form of concentration used relatively frequently is normality, or N. Normality is expressed in terms of equivalents per liter, which means the number of equivalent weights of a solute per liter of a solution. The term normality is often used in acid-base chemistry. The equivalent weight of an acid is defined as the molecular weight divided by the number of reacting hydrogens of one molecule of acid in the reaction. Understanding equivalents requires knowing something about how a reaction works, so let's start there. Below is a basic equation for an acid and a base. HCI + NaOH ------> NaCl + H20 or Acid + base -------> salt + water In our simple equation above you can see we have the acid and base reacting to form a salt and water, and that they react equally. The acid gives 1 H+ for every -OH given by the base. So for every mole of H+ one needs a mole of -OH. This reaction is one-to-one reaction on a molar basis. One mole of acid has one reacting unit and one mole of base also has one reacting unit thus both acid and base has, in the above example, equal 1:1 reacting units. As stated above, for acids we define an equivalent weight as the molecular weight divided by the number of H+ donated per molecule. Above, the HCI gave up 1 H+ (proton) to the reaction.
Molecular weight of H2SO4 = 98.08 g = 49.04 grams per equivalent # of protons given 2 protons
Normality is the molecular weight divided by the grams per equivalent (all this results in the number of equivalents) in a given volume. For an 1 N solution we need 1 equivalent/liter. For hydrochloric acid (HCl) the equivalent weight is 36.46 grams. Therefore, for making an 1 Normal solution, 36.46 g/liter of HC1 is needed. Note that a 1 M solution is also 36.46 g/L. For molecules that can give off or accept only one proton per molecule, the Normality is equal to the Molarity. Table 1. Molecular and Equivalent weights of some common compounds.
Molecular Equivalent Chemical name
Formula weight g/mol
weight g/equiv
hydrochloric acid
HCI
36.46
36.46
nitric acid
HNO3
63.01
63.01
sulfuric acid
H2SO4
98.08
49.04
phosphoric acid
H3PO4
98.00
32.67
tartaric acid
C4H6O6 150.09
75.05
malic acid
C4H6O5 134.09
67.05
citric acid
C6H6O7 192.12
64.04
lactic acid
C3H6O3
90.08
90.08
acetic acid
C2H4O2
60.05
60.05
water
H2O
18.02
18.02
sodium hydroxide
NaOH
40.00
40.00
56.11
56.11
potassium hydroxide KOH
In the case where a molecule can give off or accept more than one proton, you need to adjust your calculation. For example, sulfuric acid with a formula of H2SO4 donates 2 separate protons. Using the molar mass of sulfuric acid, and knowing that one molecule can donate 2 protons we can find the equivalent weight. With a molar mass of 98.08 grams, a solution containing 98.08 g in 1 li ter would have a Molarity of 1 M and a Normality of 2 N. This is because every I mole of sulfuric acid (H2SO4) has 2 moles of H+ atoms. Table 1 lists the molecular weights and equivalent weights of important acids and bases used in a wine laboratory. Making 1 N solution of NaOH
From the discussion above, it should be clear that to make 1 Normal solution we need to know the, equivalent of NaOH, which is calculated by dividing Molecular weight by 1, that is 40 divided by 1= 40. So the equivalent weight of NaOH is 40. To make 1 N solution, dissolve 40.00 g of sodium hydroxide in water to make volume 1 liter. For a 0.1 N solution (used for wine analysis) 4.00 g of NaOH per liter is needed. Standardization
Before we begin titrating that wine sample we have one more important step, standardization of NaOH solution. Standardization simply is a way of checking our work, and determining the exact concentration of our NaOH (or other) reagent. Maybe our dilution was inaccurate, or maybe the balance was not calibrated and as a result the normality of our sodium hydroxide solution is not exactly 1 N as we intended. So we need to check it. This is achieved by titrating the NaOH solution with an acid of known strength (Normality). Generally 0.1 N HCI is used to titrate the base. The reagent, 0.1 N HCI solution is purchased from a chemical supplier that is certified in concentration. That means it was standardized to a base of known concentration. "But isn't that going in circles?" you ask. No, because acids are standardized to a powdered base called KHP, or potassium hydrogen phthalate. This can be very accurately weighed out because it is a fine powder, and then is titrated with the acid. To standardize NaOH, start by pipetting 10.0 ml of 0.1 N hydrochloric acid (HC1) into a flask. Add approximately 50 ml of water (remember, not tap water) and three drops of methyl red indicator. Fill a 25 ml buret with the 0.1 N sodium hydroxide solution and record the initial volume. Titrate the hydrochloric acid to the point at which a lemon yellow color appears and stays constant. Record the final volume. Subtract the initial volume from the final to yield the volume of NaOH used, and plug that into the equation below. Normality of NaOH = Volume of HCI x Normality of HCI Volume of NaOH used Titration Techniques
Before conquering volumetric analysis totally, we need to discuss some tit ration techniques. First of all, handle the buret with care. Avoid damaging the tip and petcock assembly because damage and leaks in these areas can and will alter performance. Also, be sure to always record your final and initial volume readings accurately by reading the bottom of the meniscus of the solution. Don't try to squeeze in that last sample and drain the buret past i ts lowest mark; take the time to refill it properly. For help in reading a buret, take a white index card and color a black square on it as shown. Hold this behind the buret scale when taking readings to aid in seeing the meniscus. Some burets actually come with a stripe painted on them for this reason. Next, remember to stir your sample as you titrate. Whether using a stir plate (r ecommended) or stirring by swirling the flask manually, it is imperative that the solution be mixed. Be sure not to slosh the sample outside of the beaker/flask and don't allow the buret's contents to fall outside of the beaker. Also, lower your buret enough so that splatter from the sample does not exit the flask as you titrate. This is not only bad lab practice but can also be dangerous.
Safety is an important consideration when working with burets, acids and bases. Realize that you are handling corrosive chemicals and delicate glassware, treat it like an irreplaceable wine in the daintiest glass. That means deliberately and with respect. Wear safety glasses and a labcoat at least, and gloves are also recommended. When filling a buret, take it out of the stand and hold it at an angle with the tip above the sink. That way any spills will drain into the sink and you can stand safely on the floor, not a stool. Leaning over the buret while it is on the benchtop is dangerous. Be sure to have access to an eyewash station or something that can supply a stream of water to your body and/or eyes for 15 minutes, the OSHA recommended treatment for chemical spills to the eyes and body. Remember you will have sodium hydroxide in the buret at and above eye level so make sure your equipment is attached to a steady base. Good laboratory practices can help you monitor the quality of your wines more accurately and efficiently. Volumetric analysis by titration is one of the most common techniques the winemaker employs to analyze his product. Improving your skills in this area is important in the quest for excellent wines on a consistent basis. *Previously published in Vineyard and Vintage View , Mountain Grove, MO.
Determining the Molar Concentration of Vinegar by Titration Objective: Determine the concentration of acetic acid in a vinegar sample
Expressing solution concentration. Using volumetric glassware: pipet and buret. Performing a titrimetric analysis.
Background
In a titration, the analyte (the substance whose concentration is unknown and sought i n the analysis) is reacted with a standard (a substance that reacts with the analyte but whose concentration is known). The analysis uses just enough of the standard to react with all of the analyte, thereby allowing the amount of analyte present to be determined. In this experiment, acetic acid (CH3COOH) is the analyte and sodium hydroxide (NaOH) is the standard. The reaction is: CH3COOH(aq) + NaOH(aq) --> CH 3COONa(aq) + H2O(l) Titration: an analytical procedure involving a chemical reaction in which the quantity of at least one reactant is determined volumetrically. Standard solution: a solution in which the concentration of a solute is precisely known. Usually it is the volume of the standard solution required to react with a given quantity of an analyte that is precisely determined during a titration. Titration endpoint: The quantity of reactant in the standard solution added during the titration is stoichiometrically equivalent to the quantity of reactant in the analyte at the titration endpoint.
Expressing Solution Concentration mols (solute)
molarity = ------------------ = M (mol/L) Liter (solution) 0.493 M NaOH means 0.493 mol NaOH/L mol s mols = ------ x L L mols = M
xV
In a titration procedure, 40.57 mL of 0.493 M NaOH solution was used. How many mols NaOH did this volume of NaOH solution contain? mols = M x V 0.493 mols NaOH mols = ----------------------- x 0.04057 L L mols = 0.0200 mols NaOH
Volumetric glassware: buret and pipet Reading the buret
Using the pipet
Buret reading = 0.76 mL
Determining the Volume of Titrant Delivered in a Titration Final buret reading: 49.37 mL Initial buret reading: 0.74 mL volume delivered:
Procedure You may work in groups:
48.63 mL
Each person must perform a titration. For those working in groups, each group member will contribute the results of one determination to the group effort. Your reported result will be the average of at least two titrations.
Sample Calculation Calculating the concentration (M) of CH 3COOH in commercial vinegar. 1. From the balanced chemical equation: mols CH3COOH(vinegar) = mols NaOH(titrant) 2. mols NaOH = MNaOH x VNaOH,L (from titration) 0.493 mols NaOH mols Na OH = ------------------------ x 0.04863 L L
3. mols NaOH = 0.240 = mols CH 3COOH(vinegar) 4. mols CH3COOH(vinegar) M CH3COOH(vinegar) = ----------------------------volume(vinegar) 0.0240 mols M CH3COOH(vinegar) = ---------------- = 0.96 M 0.0250 L
Chemical Analysis by Acid-Base Titration Introduction:
Titration is a common method of determining the amount or concentration of an unknown substance. The method is easy to use if the quantitative relationship between two reacting solutions is known. The method is particularly well-suited to acid-base and oxidationreduction reactions. Titrations are routinely used in industry to analyze products to be sold. Many manufacturers are under strict standards of quality control because their products are sold for public consumption. In this experiment, we will analyze a number of commercial products and, in some cases, test the validity of the information given on their labels and/or
the claims made in television commercials. The products to be tested include antacid tablets, vinegar, fruit juice, and household ammonia.
Outline:
I. Standardization of a base (NaOH) using a primary standard (KHP) II. Standardization of an acid (HCl) with the standard base III. Titration analysis of unknown acids and bases: A. Antacid tablets (Tums vs. Rolaids vs. Maalox) B. Vinegar C. Fruit juice (apple or grape) D. Household ammonia
I. Standardization of NaOH
Titrations permit the concentrations of unknown acids/bases to be determined with a high degree of accuracy. In order to analyze unknown acids/bases, we must have a "standard" solution to react with the unknowns. A standard solution is one in which the concentration is known accurately. We will first prepare a standard solution of NaOH. One way to prepare a standard solution is to dissolve an accurately massed amount of the substance and dilute it to a measured volume (like we did with the MnO 4- solution in Expt. #1.6). In this way, the concentration can be calculated exactly. However, it is usually impossible to obtain NaOH of sufficient purity to use it as a primary standard. An indirect method is more practical for obtaining a standard solution of NaOH. We will prepare a solution of a approximate molarity and standardize it against a primary standard of known purity.
A primary standard should:
* be of high purity * remain unchanged in air during massing and remain stable during storage * have a high molar mass to reduce massing errors * react with the solution to be standardized in a direct, well-defined reaction
Potassium acid phthalate will serve as our primary standard. This is a large molecule (KHC8H4O4) with a molar mass of 204.2 g/mol. Instead of writing the whole formula, we abbreviate it as KHP, where "P" stands for the phthalate ion, C 8H4O42-, not for phosphorus. KHP is an acidic substance, with the ionizing hydrogen being set forward in the formula for emphasis. Therefore, KHP is monoprotic and will react with NaOH in a simple 1 to 1 relationship according to the following equation:
NaOH(aq) + KHC8H4O4(aq) ---> KNaC8H4O4(aq) + H2O(l)
If we write this reaction showing the complete structure for KHP and indicating the ionization which occurs, we have:
After we standardize our base with KHP, we will use this standard base to standardize our acid so that we will know the exact concentrations of both solutions and can then easily titrate and analyze the unknowns.
Procedure:
1. Dilute enough 1 M NaOH to make 1 L of 0.1xx M NaOH. This will be your titrant. Rinse a buret with water and then with a small amount of the NaOH solution. Fill the buret with NaOH solution. Fill the buret tip by momentarily opening the stopcock. Label this buret BASE. Read the initial volume. 2. Accurately mass approx. 0.8xx g of KHP into a 250 mL Erlenmeyer flask. Add about 100 mL of water and swirl the flask until the sample is dissolved. Add 3 drops of phenolphthalein indicator (colorless in acidic solution; pink in basic solution).
3. Titrate the KHP solution with the base solution to be standardized. Titration should proceed until the faintest pink persists for 30 sec. after swirli ng. The color will fade upon exposure to the air (WHY?). After completing a trial, breathe into the flask and swirl. What has happened? 4. Make duplicate determinations and calculate the average molarity of the NaOH. For excellent work, the molarities need to be within 1% of one another.
Data (for each trial):
* mass of KHP in sample flask * initial buret reading * final buret reading * volume of NaOH required to neutralize the KHP
Calculations:
* molarity of NaOH for each trial * avg. concentration
Question:
1. Explain why the pink color of a phenolpthalein endpoint will disappear after exposure to atmospheric air (or the air we exhale). Include 2 balanced equations. 2. Explain why the high molar mass of KHP is an advantage to its being used as a primary standard. Make up a scenario with numbers and % errors.
Now you are ready to standardize your acid with your known concentration of base.
II. Standardization of HCl Solution
Procedure:
1. Obtain about 50 mL of 0.1xx M HCl solution in a beaker. Rinse a second buret with water and then a small amount of HCl. Fill the buret with HCl solution. Fill the buret tip by momentarily opening the stopcock. Label this buret ACID. Record the initial volume and then deliver approx. 10.xx mL (as long as you know accurately how much you used) into a clean 250 mL Erlenmeyer flask. Add about 100 mL of water and 3 drops of phenolphtalein indicator to this sample flask. 2. Read the initial level in the standard NaOH buret. Titrate the acid with the NaOH standard to the faint pink endpoint. 3. Repeat the titration with a second sample of HCl. The calculated molarities of HCl should be within 1% of one another.
Data (for each trial):
* volume of acid sample * initial buret reading * final buret reading * volume of NaOH used * molarity of standard NaOH Calculations:
* molarity of HCl for each trial * avg. concentration of standardized HCl
We are now ready to titrate any and every unknown acid and base with our known concentrations of acid and base!
III. Titration Analysis of Unknown Acids and Bases
A. Analysis of Antacid Tablets
*You will be sign up for either Rolaids, Tums, or Maalox
Procedure:
1. Obtain an antacid tablet and mass the whole tablet. Grind the tablet with mortar and pestle. Mass two portions of approx. 0.6xx-0.7xx g each. 2. Add the samples to two 250 mL Erlenmeyer flasks. From your acid buret, add approx. 60.xx mL (must be known accurately) of the standard HCl to the antacid powder. Swirl the flask to dissolve the solid. All of it may not dissolve, since there may be some insoluble components. 3. Add 4-5 drops of methyl orange indicator to the sample flask. The solution should turn pink-red since the HCl was added in excess. If the solution is yellow, add approx. 10.xx mL more acid, measuring the volume carefully. 4. The antacid has neutralized a portion of the acid. Our job is to figure out exactly how much HCl was neutralized by titrating the acid left over with our standard base. "Back-titrate" the excess acid in the flask with the standard NaOH. As you approach the endpoint, add the base drop by drop. The endpoint of the titration occurs when the solution turns from red to pale orange. 5. Repeat with the other antacid sample.
Data (for each trial):
* mass of whole antacid tablet * mass of antacid sample * volume of HCl added * concentration of standard HCl * initial buret reading of NaOH * final buret reading of NaOH * volume of NaOH required * concentration of standard NaOH solution
Calculations (for each trial):
* mmol of HCl added
* mmol of NaOH added * mmol of acid neutralized by the antacid * mmol of acid neutralized per gram antacid * avg. mmol acid neutralized/g antacid
Data and Calculations (based on class data):
* class avg. mass of tablet for each brand * class avg. mmol acid neutralized per gram antacid for each brand * cost of bottle for each brand * cost of antacid per gram (cents/g) for each brand * best buy (cents/mmol acid neutralized)
Questions:
1. Rolaids once ran a national advertising campaign in which it was claimed that Rolaids consumes "forty-seven times its own weight of excess stomach acid". Can this be true?!! *Assume that stomach acid is 0.100 M HCl and that its density is 1.00 g/mL. 2. Tums once ran a national advertising campaign in which it touted that its product "neutralizes one-third more acid than Rolaids". Are they lying?!
B. Analysis of Vinegar
Percent acidity is the basis for determining the legality of vinegar. The percent acid in a sample is calculated as g/mL x 100. Vinegar is regulated by the Food and Drug Administration (FDA) and must have a minimum of 5% acidity. The acid involved is acetic acid.
Procedure:
1. Pipette 5.00 mL of an unknown vinegar sample into a 250 mL Erlenmeyer flask. Add about 100 mL of water and 3 drops of phenolphthalein indicator. Titrate with your standard base. Complete a duplicate determination.
Data (for each trial):
* volume of vinegar sample * initial buret reading * final buret reading * volume of base required * avg. volume of base
Calculations:
* the acidity of the vinegar expressed as molarity * the acidity of the vinegar expressed as a percent (is the vinegar legal?)
C. Analysis of Fruit Juice (Apple or Grape Juice) Procedure:
1. Pipette 10.00 mL of fruit juice into a 250 mL Erlenmeyer flask. Add about 100 mL of water and 3 drops of phenolphthalein. Titrate with your standard base. Due to the yellow color of the juice, the endpoint will appear peach instead of pink. Complete a duplicate determination.
Data (for each trial):
* volume of fruit juice sample * initial buret reading * final buret reading * volume of base required
* avg. volume of base * concentration of standard base
Calculations:
1. The acidity of the juice expressed as molarity 2. The acidity of the juice expressed as a percent 3. The juices are fortified with vitamin C. Determine the percent acidity from ascorbic acid versus the native acid. Extension:
build the carboxylic acids (acetic, malic, tartaric, and ascorbic) in ChemSite (see separate instructions).
D. Analysis of Household Ammonia
Procedure:
1. Pipette 5.00 mL of the ammonia solution into a 250 mL Erlenmeyer flask. Add about 100 mL of water and 3 drops of methyl orange indicator. The solution will t urn yellow. Titrate the sample with your standard HCl solution. The endpoint is orange (if you go too far, the solution will become red). Complete a duplicate determination.
Data (for each trial):
* volume of ammonia sample * initial buret reading * final buret reading * volume of acid required * avg. volume of acid required * concentration of standard acid
Calculations:
1. The molarity of ammonia
