...- Medicinal Chemistry - Question No. 3. Define the Hammett Constant and the hydrophobic (Hansch) substituent constant. Comment on how inductive and mesomeric polar effects are treated in substituted aromatic systems. A quantitative structure-activity relationship (QSAR) is an equation which correlates measurable or calculable physical or molecular properties to some specific biological activity. Once this relationship has been determined, it is possible to predict the biological activity of related drug candidates before they are put through expensive and time-consuming biological testing. The electronic effects of a substituent have an effect on the ionisation or polarity of a drug. This in turn may affect how easily a drug can pass through...
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...of “H”, or hydrogen. The lower something is on the pH scale, the higher its acidity. Rust can weaken a coin or any other metal material. The green patches on coins are called copper oxide. It is caused by oxidation. When an acid reacts with tarnish, it makes the coins shiny. In order for oxidation to occur, the air has to be damp. In my experiment I used specific solutions to clean coins. Lemon juice has a ph of 2.2. Baking Soda has a pH of 8.4. Water has a Ph of 7. Orange juice has a Ph of 3.70. Cola has a Ph of 3.18. Dishwashing liquid has a pH of 7.80. Solutions with a pH less than 7 are acidic and solutions with a pH greater than 7 are alkaline or basic. Pure water is neutral and isn’t an acid or a base. pH is an important measurement used in various medical, biological, chemical, environmental, and nutritional labs. pH standards are determined using a concentration cell with transference, by measuring the potential difference between a hydrogen electrode and a standard...
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...Acid Rain --------- Acid Rain is caused by pollution containing sulfur dioxide, nitrogen oxide, and ozone ( SOý, NOx, and Oý ) is released into the air. These chemicals are absorbed into clouds and results in Acid Presipitation ( Acid Rain, Acid Snow, Acid Hail, Acid Sleet ). When the chemicals aren't absorbed into clouds, they can drift for miles and fall to the ground, resulting in Acid Deposition, or dry deposition. When Acid Rain falls into water it is mixed in with the normal water and causes the pH of the entire body to be raised. Measurments on the pH ( potential Hydrogen ) scale, rise exponentialy, thus, a lake with a pH of 4 is ten times as acidic as a lake with a pH of 5, and a lake with a pH of 3 is 100 times as acidic, After many rain falls of Acid rain, the pH of a normal lake ( 5.8 ) to 4. Acid Rain has been known to reach the acidicy of pH 2, ( battery acid has a pH if 1 ) this is a drastic change, as normal rain is average pH 5.2. Acid Rain can dissolve limestone and chalk, and corrodes outdoor structures. Statues and monuments that are left unprotected can fall victim to the unpredjudiced destruction of acid rain. Acid Rain reacts to different types of soil and rocks in two ways. 1) Acid rain will dissolve alkaline rocks and soil, or will neutralize the alkalinity. 2) Acid rain will increase the acidicy of already acidic rocks and soil, such as granite...
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...primary standards and give two examples. primary standard should fulfill the following requirements: 1) It should be100.00% pure, although 0.01-0.02% impurity is tolerable. 2) It should be stable to drying temperatures, and it should be stable indefinitely at room temperature. The primary standard is always dried before weighing. 3) It should have a high formula weight. A high formula weight will reduce experimental error since a relatively large amount of it will have to be weighed in order to get enough moles to perform the titration and, the relative error in weighing a greater amount of material is smaller than for a small amount. * Potassium hydrogen phthalate (usually called KHP) for standardisation of aqueous base and perchloric acid in acetic acid solutions * Sodium carbonate for standardisation of aqueous acids: hydrochloric, sulfuric...
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...Jonte Berry LAB 3 REPORT SHEET – ACIDS, BASES, INDICATORS, pH Procedure Number 3 Estimated pH with pH paper Vinegar (Ph 2) Soap + H2O (Ph 6) Tap water (Ph 8) Baking soda + H2O (Ph 9) Ammonia (Ph 13) 4 What color is your “red cabbage solution” when diluted with tap water? (The water turns ruby red) Do you think we will all have exactly the same color? Explain your answer. (No) 5 Solution color Estimated pH with cabbage indicator with cabbage indicator Vinegar (Light pink) (Ph 4) Soap + H2O (Clear Pink) (Ph 4) Tap water (Light blue) (Ph 10) Baking soda+H2O (Light light blue) (Ph 10) Ammonia (Clear) (Ph 10) 6 Describe what happened to the color of the solution when you mixed the vinegar and ammonia solutions. What do you estimate the pH of the solution to be with pH paper? (The pink color from the vinegar changes to light blue when the ammonia was added.) What do you estimate the pH of the solution to be with the cabbage indicator? (Ph 10) 7 What happened when you added the baking soda solution to the mixture of vinegar/ammonia? Describe your observations. (When I added the baking soda solution nothing happen at all.) What do you estimate the pH of the solution to...
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...Experiment 1 Determination of the Acetic Acid Content of Vinegar. Goal: During this experiment you will gain experience in working with volumetric glassware, and gain knowledge on how to determine the uncertainties in the volumes delivered by these. In addition, you will use a suitable titration to determine the concentration of acetic (ethanoic) acid in vinegar samples. Objectives: On completion of the laboratory you should be able to: Demonstrate proper techniques for using `pipette, burette and volumetric flask. Assess the random error in the volume delivered from the pipette and a burette. Accurately dilute a sample. Use a suitable titration to determine the concentration of ethanoic acid in vinegar. Theory: All measurements (volumes, lengths, weights etc.) have associated errors. Some, called gross errors, arise from mistakes (writing down the wrong number, recording the wrong units, etc.) but these can be easily avoided by working carefully. Others, called systematic errors, arise from equipment or instruments not operating according to their specifications (for example a pipette delivers 4.96 cm3 rather than the stated 5.00 cm3) or something goes wrong with the measurement procedure (for example there is something unexpected in the sample being studied (called an interferant) that results in the measurement being different from what it would be if the interferant was not there). Systematic errors lead to inaccurate results but can be avoided if the cause...
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...Warm water bath (not boiling).After 3 minutes, pour each tube of H2O2 into the corresponding tube of liver and observe the reaction. The boiling water bath will be left for 5 minutes, after which the liquid will be poured off and the liver will be placed into a 2ml solution of hydrogen peroxide. This is done to see the effect that boiling has one the enzyme. Record the rate of each reaction from 0(no reaction)-5(very fast). Test the effect of pH on catalase activity using basic, acidic and neutral solutions. 2 ml hydrogen peroxide will be added to each of 5 clean test tubes. Tube 1--add 3 drops of HCl (acid) Tube 2 - dilute HCl (1 drop / 3 ml) – add 3 drops Use pH paper to determine the exact pH Tube 3 – add 3 drops of NaOH (Base) Tube 4 – dilute NaOH (1 drop / 3 ml) - add 3...
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...with a base. Titration is the process of adding a known amount of a solution of known concentration to a known amount of solution of unknown concentration. The more accurately the concentration of the solution of known concentration is known, the more accurately the concentration of the unknown solution can be determined. Some chemicals can be purchased in a pure form and remain pure over a long period or time. Other chemicals are easily contaminated by the absorption of carbon dioxide or water from the air. Sodium hydroxide absorbs moisture from the air and often appears wet. Thus if a solution of sodium hydroxide is prepared by weighing the sodium hydroxide, the concentration of the solution may not be precisely the intended concentration. Potassium hydrogen phthalate on the other hand, has a lesser tendency to absorb water from the air and when dried will remain dry for a reasonable period of time. Potassium hydrogen phthalate may be purchased in pure form at reasonable cost. Potassium hydrogen phthalate is a primary standard. This means that carefully prepared solutions of known concentration of potassium hydrogen phthalate may be used to determine, by titration, the concentration of another solution such as sodium hydroxide. The equation for the reaction of potassium hydrogen phthalate with sodium hydroxide is: KCO2C6H4CO2H + NaOH ( KCO2C6H4CO2Na + H2O The equivalence point of a titration occurs when chemically equivalent amounts of acid and base are present...
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...out into the virtual field for additional research. Please type your answers on this form. When your lab report is complete, submit it to the Submitted Assignments area of the Virtual Classroom. Part I: Answer the following questions while in the Phase 1 lab environment. Section 1: You will be testing 4 known solutions for pH levels using a standard wide-range indicator. Based off of the results obtained in the lab room, fill in the following table: |Solution Number |pH from Lab |Acid, Base or Neutral? |Solution Name (what was in the test tube?) | |Solution 1 |6 |Acid |Pure Water | |Solution 2 |1 |Acid |Lemon Juice | |Solution 3 |12 |Base |Bleach | |Solution 4 |5 |Acid |Coffee | 1. How many drops of wide range indicator will you use for each test, based on industry standards such as the LaMotte field test? 10 Section 2: Now that you understand how to read pH measurements, go out into the field to gather pH samples from 3 different lakes to take back to the lab...
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...In table one, KHP, a salt was used to titrate NaOH to find the molarity for the following two titrations. In table two and three, the known molarity of NaOH, a base, was used to find the molarity of acetic acid and sulfuric acid. The phenolphthalein indicator was used in all three titrations to make it possible to visually see when the solution reached the endpoint point. In the lab, it is impossible to record the equivalence point, when the moles of base equals moles of acid, so the endpoint is recorded. The endpoint is the point when the solutions turns from clear to light pink. The endpoint is very close to the equivalence point so it is assumed that they are same in the lab. This allows the moles of the acid to be found and the molarity to be calculated. The accuracy of this technique could be improved if it was possible to measure the equivalence point instead of the endpoint. Overall, the precision and accuracy of this test is good, but could be improved by performing more tests and averaging them...
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...Table 1 displays how each solution was determined their acidity, alkalinity or neutrality based on the color of blue and red litmus paper. When red and blue litmus paper was dipped into sodium chloride, both papers remained the same color indicating that the solution is neutral. The ammonium chloride is revealed as acidic due to the blue litmus paper changing to a red color. The red litmus paper turned blue when it was dipped into sodium acetate demonstrating that it is a basic solution. Finally, the red litmus paper continued to be red in the presence of calcium chloride for all four trials. However, when the blue litmus paper was dipped into the solution, it turned red for one trial and blue for the other three trials. Due to the majority of the trials having a red color, calcium chloride was established as acidic. Table 2 demonstrates the pH levels that were determined using the pH probe and LabQuest. The sodium chloride’s pH levels for all four trials were 5.82, 5.38, 5.34, and 5.46, indicating that sodium chloride is acidic. The pH levels for ammonium chloride were 14.5, 5, 4.94, and 4.95. Most of the trials for ammonium chloride were below seven, therefore making the solution acidic. Sodium acetate’s pH levels were 9.55, 14.71, 10.72, and 11.3 presenting the solution as basic. Lastly, the calcium chloride’s pH levels were 9.92, 12.41, 4.59, and 4.79. The calcium chloride’s first two trials were above seven due to error, thus establishing calcium chloride...
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...release and H2PO4 was very close. The 1st equivalence point for unknown solution was 1.920 pH with half-equivalence points of 0.563 pH. The Ka was 0.274 and the pKa was 0.563. While The 2nd equivalence point for unknown solution 8.811 pH with half-equivalence point of 6.347 pH. The pKa was 6.347 and the Ka was 4.5 x 10-7. The data fits extremely well with the rest of the results we collected. To determine the concentration of the unknown, we took first equivalence point, 2.375mL of NaOH multiplied by concentration of 0.0804M NaOH divided by 5mL, then divided by 1000 to achieve concentration of unknown acid as 0.0382M. Figure 1. Titration of weak acid (H3PO4) with a strong base (NaOH). Figure 2. Titration of weak acid (H2PO4) with a strong base (NaOH). Figure 2. Titration of unknown acid with a strong base (NaOH). Conclusion: The unknown solution is H3PO4 with a concentration of 0.0382M. Reason for the unknown being a H3PO4 solution because it has similarities between Figure 3 and Figure 1, also because there are two titration curves of Figure 3, which means there is a release of two H+ ions. I am fairly confident in my results because I had relatively clear equivalence and half-equivalence points that made sense on the graph, as well as relatively clear titration curves. Reasons I am not extremely confident is because the unknowns solution (Figure 3) only took around 200 drops of NaOH to finish the titration curves, but H2PO4 (Figure 2) took over 400 drops of NaOH to finish...
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...EXPERIMENTAL For Part A of the experiment, 50 mL of 0.05 M acetic acid solution was prepared with the use of the 0.5M acetic acid the stock solution. To star, 45 mL of deionized water were poured into a 150 mL beaker that was labelled “0.05 M acetic acid.” A 10 mL graduated cylinder was used to add 5 mL of 0.5 M acetic acid to the beaker, and the solution was stirred with a glass stirring rod, which was clean each time after use to . Then, a 100 mL of 0.5 M NaOH solution were prepared with the help of a 50 mL graduated cylinder to add 90 mL deionized water to a 250 mL beaker labelled “0.05 M NaOH.” A 10 mL graduated cylinder was used to add 10 mL of 0.5 M NaOH to the beaker, and the solution was stirred. After, the LabQuest and pH probe were prepared for data collection. “Sensors” and “Events with Entry” were selected after the pH probe was plugged into Channel 1. After that, in the Name box, “Volume” was typed into the designated box, and the corresponding units, “mL,” were also added. All changes made were saved. To proceed, the burette was “conditioned” after it was secured the ring stand, and the stopcock was closed in order to add 10 mL of deionized water to clean and rinse the inside of the burette. This procedure was done twice with water, and another two times with the prepared NaOH solution. A plastic funnel was used each time to pour any given liquid into the burette, and after, before recording the volume, the funnel was removed. Then, the burette was filled with the...
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... POTENTIOMETRIC TITRATIONS & SOLUBILITY EQUILIBRIA Tianyi Hu Partner: Yiteng Zhang, Teaching assistant: Nicholas Vizenor Lab section: Thursday 6 pm Date: May 7, 2015 Abstract In this experiment, the burette was used to titrate sodium chloride solution by adding different amount of silver nitrate solution. The electrochemical potential was measured by a reference electrode in order to understand the Nernst equation and calculate out experimental result of concentration of silver nitrate solution, Ksp of Ag ion, and the concentration of chloride ion in the unknown sodium chloride solution. There are four titration performed in this experiment, 2 rough titration of both 150 ppm and unknown sodium chloride solution and 2 careful titration. After collecting data from 8 groups, the equilibrium point are approximated. Since the mole of chloride ion and silver ion are equal, the concentration of silver nitrate solution was calculated. After that, the Ksp of Ag which is equal to [Ag+]^2 at equilibrium point from both known and unknown solution careful titration was computed by using the Nernst equation and was compared with the literature value. Finally, the concentration of chloride ion was also calculate by using two different methods, volume of titration and Nernst equation. Introduction The applications of the titration will performed in this experiments are: 1) rough titration (added 2 mL titrant each time) to determine roughly the equivalent point; 2) careful titration(when...
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...Title Acid-base titration: Determination of the percentages (%) of sodium carbonate (Na2CO3) and sodium hydroxide (NaOH) in a mixture Objective To determine the respective weight per cent of sodium carbonate and sodium hydroxide in a mixture by acid-base titration. Result and calculation Part A Titration 1 Titration number 1 2 3 Initial volume of burette( cm3) 5.10 2.70 9.70 Final volume of burette (cm3) 34.40 31.80 39.20 Total volume of HCl used (cm3) 29.30 29.10 29.50 Average volume of HCl required for titration =(29.30+29.10+29.50)/3 cm3 = 29.30 cm3 Titration 2 Titration number 1 2 3 Initial volume of burette( cm3) 4.50 14.00 2.70 Final volume of burette (cm3) 25.00 21.70 22.80 Total volume of HCl used (cm3) 20.50 20.30 20.10 Average volume of HCl required for titration =(20.50+20.30+20.10)/3 cm3 = 20.30 cm3 Part B Titration number Rough 1 2 3 Initial volume of burette( cm3) 4.9 4.80 3.60 2.20 Final volume of burette (cm3) 28.3 28.90 27.70 26.20 Total volume of HCl used until phenolphthalein decolourised (cm3) , x 23.4 24.10 24.10 24.00 Initial volume of burette after adding methyl orange indicator ( cm3) 28.3 28.90 27.70 26.20 Final volume of burette (cm3) 34.1 33.40 32.10 30.70 Total volume of HCl used until phenolphthalein decolourised (cm3) , y 5.8 4.50 4.40 4.50 Average volume of HCl required to react with Na2CO3 (2y) =2(4.50+4.40+4.50)/3 cm3 = 2(4.4667) cm3 ...
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