I just need a step by step on how to do these. Like one calculation explained per section at least. I'm just not sure exactly how to do it. I think I need to use the m1v1=m2v2 but not sure how the tap water affects that

isk and add 3 ml of pH 10 eBl indicator (caution! EBT is made up in organic solvent sidered toxic, add the EBT in the hood). 2. Titrate with EDTA solution recording the titration beginning and endpoint volumes Save the flask with the appropriate endpoint color for comparison. 3. Knowing the approximate titration volume, you are ready t titration. Pipet 10.00 mL of the calcium unknown into a 125-ml. Erlenmeyer 3 mL of pH 10 buffer and about 6 drops of EBT indicator. Titrate to the e carry out a quantitative ration volumes. Repeat the titration for a total of at least three runs (not counting the "practice" trial).ã TA titrations require a blank. Perform one blank titration with 10 mL of deionized er. It should only require a small amount of titrant. Subtract this volume from each result.- 5. Calculate the average concentration of Ca2 in molarity and in ppm Caco. Titration of Tap Water A. Determination of Total lons (Ca2 and Mg2) 1. Pipet (or measure carefully with a graduated cylinder) 50.0 mL of tap water into a 250-mL Erlenmeyer flask and add 3 mL of pH 10 buffer and about 6 drops of EBT indicator (caution! EBT is made up in organic EBT in the hood). solvent and is considered toxic, add the Titrate with EDTA solution recor ding the titration beginning and endpoint volumes. 2. ou may need several tries to be able to see the endpoint color (wine red to blue). Save the flask with the appropriate endpoint color for comparison. 3. Repeat the titration a minimum of three times (not counting any "practice" runs). 4. EDTA titrations require a blank. Perform one blank titration with 50.0 mL of deionized water. It should only require a small amount of titrant. Subtract this volume from each result. 76 TEXAS INSTRUMENTS TI-8

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Your unknown is an aqueous solution containing an amount of a magnesium salt and hydrochloric acid. Quantitatively transfer a 10.00 mL aliquot of the unknown solution to a 250 mL Erlenmeyer flask. Add 2-3 drops of methyl red. and titrate with your base. Express the results as the molarity of HCI in the unknown. This step is completed to determine the amount of titratable acid that may be present in your sample. You may repeal replicate titrations if you wish, though you must make sure you retain enough solution for the ion exchange portion of the experiment. Add 10.00 mL aliquot of your unknown solution to the top of the prepared column. Start collecting your aliquots in a clean Erlenmeyer flask. After the unknown has eluted nearly to the top of the resin bed, wash with about one bed volume of distilled water. Repeat the washing two more times, allowing each volume of wash to elute to the top of the resin bed before adding the next volume. All of the effluent should be collected. Add 3-4 drops of methyl red to the effluent and titrate with your 0.1 M NaOH. Do not discard the solution. After the endpoint is attained, place the titration flask under the column, and elute another one-half bed volume of distilled water into the solution you previously titrated. If the indicator changes to red (acidic solution), continue the titration with your standardized base to a yellow endpoint. Then repeat the process of eluting one-half bed volumes of distilled water and titrate if necessary, until the eluted distilled water does not require more than one drop of the base to change the indicator to yellow. Repeat the procedure above for a total of three unknown determinations. Recharge the column by converting the column to its fully protonated form (see instructions above) between each sample. Use the total amount of base used in this process to calculate, the millimoles of H^+ displaced from the resin, the millimoles of Mg^2+ remaining in the column on exchange sites, and finally the concentration of Mg^2+ in terms of mg of Mg^2+ per mL of sample. For your final answer, report the average concentration of Mg^2+ in your unknown sample in mg/mL and the confidence limits at the 95% level. Post lab question. Experiments 9 and 10 are different methods for the determination of alkali earth metals. Use your data, calculate and compare the %relative standard deviation for experiments 9 and 10. Discuss the difference (if any) and suggest reasons why the values are different, or if they are similar, suggest reasons why values are similar.

I am unable to solve these questions due to a time constraint. I received help on part 1 of the unknown acid. Need help with the rest! I really appreciate it!

Lab: ACID-BASE TITRATION

Objective: To use titration to determine in Part 1, the concentration (molarity) of an unknown acid solution and in Part 2, the purity of a sample of KHP acid. Background: Titration is a volumetric technique used to determine the concentrations of solutions, molar masses of solids, and purity of samples. A titration inolves the addition of a titrant (a solution of known concentration) to an analyte ( a solution of unknown concentration), or vice versa, using a piece of glassware called a burette. The titration is carried out until it reaches an equivalence point (the exact point at which the reaction between the two solutions is complete). A chemical indicator is often used to aid in the identification of the equivalence point. An indicator changes color upon reaching the equivalence point of endpoint. Since we will be carrying out an acid-base titration, the indicator must change color upon reaching the endpoint at a specific pH. An example of such an indicator is phenolphthalein, which changes from colorless to pink near the equivalence point when pH approaches 7. It should be noted that the indicator selected is dependent upon a given titration, so that the observed color change is close to the ideal equivalence point. Once the titration is completed, we can use the volumes measured for each solution, as well as the concentration of the titrant, to determine the concentration, molar mass, or purity of the unknown analyte. Procedure Note: During lab, record all measurements and data in a clearly labeled table. Part 1: 1. Obtain ~80mL of ~0.1M of base (sodium hydroxide, NaOH) into a labeled beaker. Make sure to report the actual concentration from the bottle. 2. Obtain ~25mL of unknown acid into a second labeled beaker. Record whether your acid is monoprotic, diprotic, or triprotic. 3. Using a funnel, or carefully pouring, fill the burette with sodium hydroxide NaOH. Make sure a labeled waste beaker is under the buretter to catch any drippings. Rinse your burette by filling it with ~5-10mL of NaOH. Turn the knob to let ~2 mL out the tip into the waste beaker. Make sure your waste beaker is under the burette to catch any drips. Next, turn the Burette upside down and pour the remaining NaOH into your waste beaker while twisting the burette to coat the inside with NaOH. 4. Fill your burette with sodium hydroxide. Turn the knob to fill the tip of the burette with NaOH. Record the initial volume to the correct number of significant figures. 5. Using a volumentric pipette, transfer 10.00 mL of your unknown acid into the Erlenmeyer flask. 6. Add 3 drops of pH indicator and magnetic stirrer to the Erlenmeyer. 7. Under the burette, place the Erlenmeyer on a hot plate and turn on the stir knob. Heat should be off. 8. Titrate the unknown acid until the solution turns pink and remains pink. 9. Record the final volume on the burette to the correct number of significant figures. 10. Repeat (4) to (7). Pre-plan the whole lab and who will do which task. MY RESULTS: PART 1 The actual concentration of NAOH is 0.5939, 80 mL Initial Final 24.9 mL 44.5mL 26.1 mL 45.57mL PART 2 Impure KHP (SAMPLE B ) Only Paper: 0.280 g Impure KHP Mass: 0.526 g Initial volume: 22.30 mL + water +26.10 mL Final Volume: 26.80 mL QUESTIONS: 1. Calculate the concentration (molarity) of your unknown acid. 2. Calculate how many grams in your impure KHP sample was indeed KHP. 3. Calculate the percent (by mass) KHP in the sample. Part 2: 1. What would happen to the final result( state what the final result is ) if the tip of the burette was not filled when reading the initial volume? 2. What would happen to the final answer if you used Ca(OH)2 instead of NaOH and DID NOT know it? 3. What would happen to the final result if the Erlenmeyer flask was contaminated with other types of acid? 4. Recalculate the concentration of your unknown acid if it was given as a diprotic acid. 5. Draw a molecular level picture (showing Na+, H+, OH-, H2O, etc...) of what is in your beaker before, at half point, at equivalent point, and at end point of the titration. Part 3: 1. What would have happened to the final result if the impurity in the sample was also acidic? 2. What would have happend to your final result if you used ~25mL of water to dissolve the solid acid instead of 20mL? Summary data: Create a clearly labeled table of all your final data results for each run, and athe average of multiple runs. Part 1: Unknown #, Vol base, M base, Vol acid, M acid. Part 2: Sample #, % KHP (purity)

If you were to over-titrate your calcium solutions in Part B of this lab:a.Would the calculated molarity of EDTA be too high or too low? Explain.

b.Would the calculated molarity of calcium unknown in part C be too high or too low? Explain.

Part B Setup: Complete the following steps, after reading page 6 on how to use volumetric glassware:1.You should have already used a 500 mLvolumetric flaskto prepare an approximately 0.010 M solution of Na2EDTA (Part 1).2.Clean the buret carefully and rinse it two or three times with small amounts (~5mL) of your EDTA solution. Hint; use a clean, dry, labeled beaker to pour solution into the buret. Make sure you rinse through the tip and that no air bubbles are in the tip. Mount the buret on a ring stand using buret clamps. Fill the buret with your EDTA solution nearly to the top mark (to somewhere between the 0 and 5 mL marks for every trialâDO NOT waste time filling the buret to exactly 0.00 mL!). The EDTA is the titrant,the solution added during the titration. Chemistry 161 Lab5Bozlee -Based on S. Critchlow Winter 2018page 4of 11Green River Community College Standardization of the EDTA solution:4.Record the exact reading of the liquid level (bottom of the meniscus) in the buret, recording to the nearest 0.01 mL. This is your initial volume. Note:The burets can be read to a precision of at least +/-0.02 mL. Be sure that all of your readings are precise and consistent.Check with your instructor if you are unclear whether your buret readings are correct.5.In a clean dry beaker, take about 50 mL of standard 0.01000 MCa2+solution; do not take more than this at a time. This is enough to do two trials; if only doing one trial take no more than 30 mL. After proper rinsing, pipet 20.00 mLof this standard Ca2+solution into a 125 mL Erlenmeyer flask. The volumetric pipet tip should NOT be âblown out.â Add about 25 mL of deionized water to the flask.The volume of water is not critical.6.Add a few drops (~6) of indicatorand one full dropperfulof the buffer solutionto the solution in the Erlenmeyer flask.7.Place the stir bar in the flask, place the flask on the stir plate, and set the control for gentle stirring. Or, if you have a good swirling technique with the Erlenmeyer flask, the stir bar and plate is unnecessary. It is essential that the solutions be continually stirred or swirled while titrating.8.What net volume of titrant should be required for the standard trials? You should be able to predict this by calculation!Slowly add your EDTA solution from the buret until the endpoint is reached. The endpointin this titration is the point at which the solution in the Erlenmeyer flask first turns a distinct blue. The color change of the indicatorat the endpoint signals that the equivalence pointhas been reached, where the reactants are in the correct stoichiometric ratio to completely react. As you near the endpoint, add EDTA drop by drop to avoid âovershooting the endpoint.â Record to the nearest 0.01 mL the exact liquid level of EDTA in the buret (your final volume), and calculate the net volume of EDTA used in the titration by subtracting the initial volume from the final volume of EDTA.9.Discard your waste in the appropriate container, using only a very small amount of DI water as the first rinse. Then rinse out your Erlenmeyer flask thoroughlywith several portions of DI water at your sink. Refill the buret (always to between 0 and 5 mL) if you will do another trial. 10.Repeat steps 4-9 two more times, using a fresh aliquot(measured sample) of standard Ca+2solutioneach time. Ideally your replicate titrations should agree within 0.2 mL

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