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Message 1 of 4 in Discussion

From:

Abalo3 (Original Message)

Sent: 12/17/2007 5:09 PM

Hi Steve,

1. I need some help finishing up this lab.The net ionic equation of Copper and AgNO3 gives: 2Ag+(aq) + Cu(s)------2Ag(s) + Cu2+(aq)

Provide complete answers (showing calculation details) and discussion where appropriate. Use enough decimal points so that you do not create the situation where �? + 1 = 3�?(true for “large�?values of “one�?and “small�?values of “three�?) Always make sure your equations are balanced (same total number of atoms of each element on each side of the reaction arrow.) The lab instructions call for using about one gram of copper metal, but you used 5 grams. There is not enough silver in the amount of AgNO3 solution being used in this reaction to react with that much copper, so some solid copper metal would be left to contaminate your collected silver. The lab instructions also call for you to pour out the water and weigh just the beaker and the silver solids.

Thank you for your help and cooperation.

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Message 2 of 4 in Discussion

From:

·Steve·

Sent: 12/17/2007 7:55 PM

1. The first question concerns rounding. For example, what if two numbers to add together from the experiment were 1.4 g and 1.4 g. The sum is 2.8 g, almost 3 grams. However, what if you round 1.4 g to 1 g first and then add them together? You get 1 g + 1 g = 2 g, a totally different answer. This is why it is always risky to round numbers to only one significant digit, because roundoff error in the calculations can be very large.

2. Balanced reaction... always important! Either the calculations or the conclusions for the experiment are going to depend on the balanced reaction.

3. Well, the object seems to be to react all of the copper with excess AgNO3, and weigh the silver metal that forms in order to check the stoichiometry of the reaction to see if it agrees with the balanced equation. I assume that powdered copper metal is being used. If a large chunk of copper is used, the silver will coat the copper, preventing it from reacting further unless you continuously scrape off the silver as it forms. This is less of a problem if you use fine copper particles and stir them well during the reaction.

If you have a mixture of silver and unreacted copper at the end, you can still determine the stoichiometry of the reaction. The moles of silver metal formed equals the moles of AgNO3 you started with, since all of the Ag⁺ in the solution will have reacted. Moles times atomic weight of silver = grams of silver metal formed. Total weight of silver product and unreacted copper minus grams of silver formed equals grams of copper metal left over. Five grams of copper initially minus grams of copper left over equals grams of copper that reacted. You can convert grams of copper metal reacted to moles and compare to the number of moles of silver metal that formed to check the reaction stoichiometry (should be 2 moles of Ag formed per 1 mole Cu reacted).

Steve

Message 3 of 4 in Discussion

From: lackingchemboi

Sent: 3/22/2008 10:52 PM

Thanks you've been a big help thus far!

the following has to be calculated as well

T/oC	Msat/mol kg^-1	Xsat	10^3 T^-1/K^-1	ln Xsat

sample calculations would be much appreciated and th following questions areto be answered.

(a) it was assumed that delta Hfus0 is independent of temperature to obtain equation (2). When this assumption is valid what relationship exists between the heat capacities for the solid A and its liquid (fused)?

(b) Suggest a reason for taking the precaution of warming the pipette used to withdraw aliquots from the solution.

(c) comment on the differences between your values and those in the literature.( what is the literature)? can these differences be explained by approximations made in the theory underlying the experiment and by errors associated with your experimental measurements?

Message 4 of 4 in Discussion

From:

·Steve·

Sent: 3/23/2008 2:25 AM

Glad to help. Is this part of the other experiment (http://groups.msn.com/ChemistryCorner/general.msnw?action=get_message&mview=1&ID_Message=3222)? To keep everything in the same message thread, go to the last message and click "Reply" when you post a follow-up and everything will stay together. As it is, your questions are tacked on the ends of unrelated old messages, making it hard to keep track! BTW to start a new question entirely, click "New Discussion" when you are viewing the the list of questions, in the "General" board for example, or when viewing a message.

OK, looks like you are looking at the relationship of the ideal solubility (not solvent-specific) in terms of the mole fraction of solute, x_sat, when the solution is saturated. The formula is

ln (x_sat) = (ΔH_fus° / R) ( 1/T �?1/T_fus )

where ΔH_fus° is the heat of fusion of the solute

R is the gas constant in joule units, 8.314 J mol^�?K^�?

T is the temperature in Kelvins

Tfus is the melting point of the solute in Kelvins

(a) There are two approximations in use. ΔC_p can be considered to be negligible and therefore zero, or ΔC_p may be approximated to be equal to the entropy of fusion, ΔS_fus. The above equation results when ΔC_p is assumed to be zero.

(b) Well, if the warm solution is saturated and you withdraw some in a cool pipet, some of the solute would precipitate out if the solution in the pipet cools!

(c) I think they are talking about the literature or known values of the solubilities of the solute at different temperatures. The assumptions made in (a) are in part responsible for differences between the literature solubilities and those calculated from the formula. But mainly, the formula used here is independent of solvent, and in reality there will be solute-solvent interactions that will affect the solubility in a particular solvent. Such interactions are not accounted for in the formula.

Steve

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