Reshmi Nair Title: Determination of Aspirin through back titration. Aim: To determine the concentration of Aspirin in a tablet using NaOH and Hcl. Research Question: What is the concentration of Aspirin in a normal tablet? Background: Aspirin is the general name for acetylsalicylic acid (ASA); it is also the trademark of the drug produced by Bayer in Germany. In eighty countries‚ aspirin is a registered trademark‚ but in other places the term aspirin refers to ASA by itself or as an ingredient
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measure the heats of reaction for each of them. Hess’s Law states that the heat of reaction of the one reaction should equal to the sum of the heats of reaction for the other two. The three reactions used in this experiment are: (1) NaOH(s) Na+(aq) + OH-(aq) (2) NaOH(s) + H+(aq) + Cl-(aq) H2O(l) + Na+(aq) + Cl-(aq) (3) Na+(aq) + OH-(aq) + H+(aq) + Cl-(aq) H2O(l) + Na+(aq) + Cl-(aq) In
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(ChemicalsIndicators) or the context menu. Step 3: Fill buret with NaOH‚ obtain a 50 ml buret and fill with .100M NaOH solution. Step 4: Titrate NaOH into HCl until end point‚ record initial buret volume and add NaOH (quickly at first then slowly) until the HCl solution turns pink and record the final buret volume of NaOH in buret. Step 5 repeat steps 1-4 using pH meter‚ add a pH meter to the acid solution. Record numerous points of pH and NaOH added (especially near equivalence point) to be used later
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Chemistry102 5/7/2013 Lecture Presentation Chapter 17 Additional Aspects of Aqueous Equilibria John D. Bookstaver St. Charles Community College Cottleville‚ MO © 2012 Pearson Education‚ Inc. Common Ion Effect HA(aq) + H2O(l) ⇔ A−(aq) + H3O+(aq) • Adding a salt containing the anion NaA‚ which • is the conjugate base of the acid (the common ion)‚ shifts the position of equilibrium to the left This causes the pH to be higher than the pH of the acid solution 9lowering the H3O+ ion concentration
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Investigating Chemical Equilibrium Date: 30th April 2013 Due Date: 15th May 2013 Prepared For: M. Seraji Prepared by: Andrea Odunze Abstract Many reactions proceed to a state of equilibrium. A chemical reaction at equilibrium‚ where the rates of the forward reaction and reverse reaction are equal‚ looks like this: A + B AB There are three factors‚ according to Le Chatelier’s principle‚ that affect the equilibrium position and equilibrium constant. These are the concentrations of products
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(ChemicalsIndicators) or the context menu. Step 3: Fill buret with NaOH‚ obtain a 50 ml buret and fill with .100M NaOH solution. Step 4: Titrate NaOH into HCl until end point‚ record initial buret volume and add NaOH (quickly at first then slowly) until the HCl solution turns pink and record the final buret volume of NaOH in buret. Step 5 repeat steps 1-4 using pH meter‚ add a pH meter to the acid solution. Record numerous points of pH and NaOH added (especially near equivalence point) to be used
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buffer solution with 0.1 M HCl and 0.1 M NaOH. The volume of 0.1 M HCl is about double of the amount 0.1 M NaOH used to lower/raise the pH of a blood buffer. In this experiment‚ HCl (a strong acid) and NaOH (a strong base) are used as examples of strong acids/bases‚ and the titration with H2PO4 shows the effect on a buffer solution. The assumption was the addition of large amounts of HCl will lower the pH‚ while the addition of large amounts of NaOH will increase the pH‚ while small amounts
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Abstract By using acid-base titration‚ we determined the suitability of phenolphthalein and methyl red as acid base indicators. We found that the equivalence point of the titration of hydrochloric acid with sodium hydroxide was not within the ph range of phenolphthalein’s color range. The titration of acetic acid with sodium hydroxide resulted in an equivalence point out of the range of methyl red. And the titration of ammonia with hydrochloric acid had an equivalence point that was also out of
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base)‚ and its (conjugate acid) Carbonic acid was used. The purpose of the experiment was to test the capacity of an undiluted and diluted buffer solution as well as examining the buffering capacity of distilled water using measured concentrations of NaOH and HCL. These replacements of strong acids and bases for weaker ones give buffers their ability to moderate pH. (Stoker 2013). Part 2 of the experiment‚ the buffering capacity of lake water was tested. The ions naturally present in rivers are
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using 50.0 mL of 2.0 M HCl. Using Equation 3 (M=mol/L) from above‚ calculate the number of moles of acid used. M=mol/L 2.0M = x / 0.0500L x = 0.10 mol HCl 3. You will also be using 50.0 mL of 2.0 M NaOH. How many moles of base are used? M=mol/L 2.0M = x / 0.0500L x = 0.10 mol NaOH 4. The specific heat of a solution is 4.18 J/g ºC. The solution is formed by combining 25.0 g of solution A with 25.0 g of solution B‚ with each solution initially at 21.4ºC. The final temperature of the combined
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