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Hello,
Would someone explain in "normal" terms the action of aspirin in terms of intermolecular forces that impart to the molecule? I have gotten myself confused. Thanks in advance. |
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| | From:  ·Steve· | Sent: 4/3/2008 7:06 PM |
Hi, well, for one thing the water solublity is greatly reduced when salicylic acid is converted to aspirin. This is because there is less opportunity for hydrogen bonding (dipole-dipole interactions) between the aspirin molecule and the solvent water molecules. Salicylic acid has two –OH groups, one in a carboxylic acid group (–COOH) and the other as a phenol group (–OH attached to the benzene ring). Both of these groups hydrogen bond with water molecules very well, enhancing the solubility. In aspirin, the phenol –OH group is "blocked" by replacing the H with an acetyl group (–COCH3). With the acetyl group now present, the –OH group is no longer available for hydrogen bonding, with the result that aspirin is less soluble in water. You may also see references to "intramolecular" hydrogen bonding in salicylic acid, which occurs between the –COOH group and the –OH group in the same molecule, which is no longer possible after the –OH hydrogen is replaced with the acetyl group in aspirin. Steve |
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| | | From: ·Steve· | Sent: 4/4/2008 7:10 AM |
You're welcome Meri, but I think I need to revise my statements about the solubilities of salicylic acid and aspirin! According to the Merck Index, salicylic acid has a water solubility of 1 gram per 460 mL of water, temperature not stated but presumably 25°C, while aspirin, which I thought was less soluble, is stated to have a water solubility of 1 gram per 300 mL at 25°C, which is actually more soluble. Converting to units of moles per liter, aspirin still beats salicylic acid by a small amount, 0.0185 M to 0.0157 M. We routinely prepare aspirin from salicylic acid in our college chemistry labs where I teach, and afterward the students take a small amount of their product and place it in a test tube with a few milliliters of water to check for the presence of the phenol group of salicylic acid with ferric chloride solution (it forms a dark purple color with salicylic acid, but does not give any color with pure aspirin). They also do the same thing with salicylic acid, but I have always observed that more of it dissolves compared to aspirin. That is, a small "spatula tip" amount of salicylic acid dissolves pretty well, but with aspirin, undissolved solid remains. But apparently that's not really true! Since the amounts were not actually weighed, the salicylic acid may just be giving the appearance of being more soluble when it actually isn't. Well, live and learn! We can also rationalize why aspirin can be more soluble that salicylic acid. From the standpoint of intramolecular hydrogen bonding, both –OH groups in salicylic acid are tied up to some extent, rendering them less available for hydrogen bonding with solvent water molecules. This would hinder water solubility. In aspirin, with the acetyl group present, one –OH group is no longer available, but the carboxylic acid –OH is still is present and available to hydrogen bond with water. But intramolecular hydrogen bonding is no longer possible in aspirin. So, rationalizing the relative solubilities, the intramolecular hydrogen bonding in salicylic acid must compete significantly with the solvent water molecules, hindering its solubility, while in aspirin, even with only the one carboxylic acid –OH group, is more soluble because its –OH group is not hindered by competing intramolecular hydrogen bonding. Steve |
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| | | From: ·Steve· | Sent: 4/5/2008 1:56 AM |
You're welcome again, and sorry for the confusion! I looked up the solubilities and melting points of the three isomers of hydroxybenzoic acid for comparison. In 3- and 4-hydroxybenzoic acid, intramolecular hydrogen bonding will not occur (because the two –OH groups are too far apart) and as a result they hydrogen bond more strongly with each other in the solid state (intermolecular hydrogen bonding instead of intramolecular), causing their melting points to be higher than that of 2-hydroxybenzoic acid (salicylic acid). The water solubility of 4-hydroxybenzoic acid is also much higher than the solubility of 2-hydroxybenzoic acid, again because the two –OH groups in 4-hydroxybenzoic acid, being too far apart to hydrogen bond with each other, hydrogen bond to water molecules instead, increasing the solubility. Sometimes it makes sense! Steve Compound Melting Point Water Solubility
2-hydroxybenzoic acid 159°C 1 g / 460 mL
(salicylic acid)
3-hydroxybenzoic acid 201-204°C (could not find)
4-hydroxybenzoic acid 214-217°C 1 g / 125 mL |
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| Steve, I have some other problems that I have worked through kinda that I need your help on, if possible. The instructions: Think of your body as a large molecule with your hand, feet, ears, etc as functional groups. Discuss how a utensil such as a fork is designed to work with one or 'functional' groups. Relate this to one of the reactions. So here is what I have: Parts of the human body can be regarded as functional groups. In this event, alcohol can be thought of as one of these groups. When you think of your body as a whole, different parts can do certain functions. Let’s take the feet. The feet can be compared to aldehydes, ketones, and carbolytic acids.There are things only the feet can do. Take the toes away and it can perform new functions, but loses a part of the old function. An oxidation reaction involves the loss of hydrogen compared to the foot losing toes. Oxidation of a primary alcohol produces an aldehyde, of a secondary alcohol produces a
ketone, and tertiary alcohols cannot be oxidized. Compared to the oxidation reaction, in order to occur, they must contain at least one C-H bond. Because tertiary alcohols contain three C-C bonds to the carbinol carbon, they cannot undergo oxidation. The second problem is this:Choose two functional groups and for each group find a household product that contains the functional groups you chose (e.g. fingernail polish remover contains acetone a ketone). Explain the action of the product in terms of the intermolecular forces that these groups impart to the molecule. I have: The hydroxyl group makes alcohols more soluable (polar) in solvents than the hydrocarbon from which they were derived. Alcohols are capable of hydrogen bonding between molecules. Being very polar, they are able to form intermolecular hydrogen bonds. The alcohol found in alcoholic beverages is ethanol. The oxygen and hydrogen group on the molecule is what gives
it the specific chemical properties of alcohol. The carboxylic acid group makes acetic acid hydrophilic and mixable in water, ethyl alcohol, and diethyl ether because of hydrogen bonding interactions. The hydrogen atom in the carboxyl group can be given off as an H+ ion giving vinager the acidic character. Any help you are willing to give me is greatly appreciated. Thanks, meri
Meri Salcedo
You rock. That's why Blockbuster's offering you one month of Blockbuster Total Access, No Cost. |
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| | | From: ·Steve· | Sent: 4/11/2008 6:06 AM |
I'm trying to think what a fork would be analogous to.... maybe a water molecule that binds reversibly to the functional group, such as OH. The parts of the body are not permanently changed when you pick things up, put on your shoes, etc., as is typically the case in a chemical reaction. Hehe... let your shoes represent H+ ions dissociating from your carboxylic acid "feet," maybe! After all, we take our shoes off, put them back on, back and forth, kind of like a weak acid dissociation. 
We often grasp our own hands together, and we shake hands with the hands of other "molecules" (people). Hydrogen bonding of OH groups can be like that too, intramolecular and intermolecular H-bonding.
I'm still trying to think of a good human analog to an oxidation reaction, other than chopping off the toes... sounds kinda painful!  Maybe trimming your fingernails or cutting your hair.
For the second part, those examples sound fine. You can also use rubbing alcohol, which is isopropyl alcohol. Acetic acid is water-soluble, due, as you say, to strong intermolecular hydrogen bonding between the acetic acid molecules and the water molecules, just as alcohols do.
Steve
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