CHEM 177L
2017_10_13_notebook_7822
12
CHEM 177L Spring 2017 - Angie Burke
Weekly Notes/Week 8
PDF Version generated by
Angie Burke
on
Oct 12, 2017 @09:21 PM CDT
Table of Contents
Table of Contents
We
...
CHEM 177L
2017_10_13_notebook_7822
12
CHEM 177L Spring 2017 - Angie Burke
Weekly Notes/Week 8
PDF Version generated by
Angie Burke
on
Oct 12, 2017 @09:21 PM CDT
Table of Contents
Table of Contents
Week 8
Week 8
Week 8 - Heat Exchange in Chemical Process
Angie Burke, Lane Vaselaar, Emma Sass, Emma Arenholtz
Prelab Writing
Investigation Questions:
1. What kind of reactions procedure (exothermic)/ absorbs (endothermic) heat?
2. What is the sign of enthalpy for exothermic/ endothermic reactions?
3. What does the Hess's Law state?
4. Why is the Hess's Law useful?
General Experimental Procedure for Part 1-3:
1. Use a solution calorimeter for making all of your measurements. Do not allow the total volume of the solution being studied to excess
80% of the total volume of the calorimeter.
2. Suspend the thermometer from a clamp attached to the ring stand. Place the calorimeter on top of the magnetic stirrer and place the
stir bar in the calorimeter to keep the solution well stirred.
3. Add the first reagent to the calorimeter. Then screw the cap onto the bottle.
4. Start the stirrer at a moderate speed. Observe and record the temperature at 30-seconds intervals for several minutes to establish the
initial temperature of the first reagent. Make sure not to turn on heat to avoid interfering with the reaction.
5. Note the time, then loosen the screw cap and quickly and carefully add the pre-weight sample of interest, making certain the entire
sample drops directly into the first reagent.
6. Screw the cap on again and begin to record the temperature reading 30 seconds after the reactants were mixed.
7. Temperature readings continue at 30-second intervals until 10 minutes after the maximum temperature was reached. Use the
temperature value that holds steady for five minutes.
8. Use Google sheets to plot graphs of your "temperature versus time" data for each reaction.
9. Gather materials:
1. Calorimeter
2. Thermometer (calibrated as outlined above)
3. Graduated cylinder
4. Hot plate
5. Magnetic stir bar
6. Automatic stirrer
7. Ice
Part 1 Procedure: Accounting for other heat transfer (optional, check with TA)
1. You need at least 150 mL for thermometer to reach the water in the calorimeter and maximum volume of the calorimeter is about 400 mL.
2. Pour the water down into the drain.
3. Equations for finding the amount of heat transferred
1. q = mCΔT, where ΔT = T final - Tinitial
2. qgained = - qlost
3. q gained by cold water + q calorimeter = -q lost by warm water
1. q gained = [mCΔT]cold water
2. q calorimeter = KΔT calorimeter
3. q lost = [mCΔT] hot water
4. After substitutions: [mCΔT] cold water + kΔT calorimeter = -[mCΔT] hot water
5. Therefore: K = (-[mCΔT] hot water - [mCΔT] cool water) /ΔT calorimeter
Part 2 Procedure: An acid-base reaction
1. You need at least 200 mL for thermometer to reach the solution in the calorimeter and the maximum volume of the calorimeter is about 400 mL.
2. Pour the waste into proper waste containers.
3. Equations for finding the heat of reaction from temperature change data.
Week 8/ 2 of 8
1. q rxn + q solution + q calorimeter = 0
2. q calorimeter = C x ΔT = 0
3. q rxn + q solution = 0
4. ΔH rxn = qrxn / moles of limiting reagent reacted
Part 3 Procedure: Dissolving salts
1. Use a mass range of 5.0 g to 8.0 g of the salts and a volume of about 250 mL
2. Dispose the waste into proper waste containers
Safety Precautions:
1. Sodium hydroxide NaOH (aq) - May cause eye and skin irritation. Inhalation will produce respiratory tract irritation.
Wear PPE.Use proper ventilation during reaction
2. Hydrochloric acid Hal (aq) - Corrosive. Causes eye, skin, respiratory and digestive tract burns
Wear PPE. Use proper ventilation
3. Cooper (II) sulfate pentahydrate CuSO 4*5H2O(s) - May cause eye and skin irritation. Inhalation may cause respiratory tract irritation.
Avoid skin and eye contact. Wash thoroughly after handling
4. Cooper (II) sulfate, anhydrous CuSO 4(s) - Hazardous in case of skin contact, eye contact, ingestion and inhalation.
Avoid skin and eye contact. Wash throughly after handling.
Procedure Part 2 - An acid and base solution
Sodium hydroxide: 100 ml.
Hydrochloric Acid: 100 ml
Ti = 20.9 C in 2 mins 50 secs.
Tf = 31.5 C in additional 1 mins 10 secs
After the two solutions are mixed together, after 7 mins: 31.3 C; 7mins 30 secs: 31.2 C; 10 mins 31.2 C
Procedure Part 3 - Dissolving salts
1. Cooper (II) sulfate pentahydrate: 5.870 g
Water: T i = 21.1 C after 1 min 57 secs
Solution1: T f = 21.0 C (21.0 C after 30 secs, 21.0 C after 1 mins, 21.0 C at 1 mins 30 secs. After 5 mins, the temperature stayed at 21.0 C.)
2. Cooper (II) sulfate anhydrous: 5.256 g
Water: T i = 21.6 C after 3 mins 21.6 C
Solution2: T f = 23.7 C ( 24.2 C after 30 secs, 24.0 C after 2 mins, 23.9 C after 2.5 mins, 23.8 C after 3 mins, 23.7 C after 7.5 mins, the
temperature stayed at 23.7 C after 10 mins)
Analysis Questions
Week 8/ 3 of 8
Part 2 result
Part 3 result
Image 1
Week 8/ 4 of 8
Picture 1 - Calculations
Analysis Questions Part 1
1. Calculate the calorimeter constant for the calorimeter you are using. If this part was discussed only, explain what the
calorimeter constant is and what assumption you made.
1. The calorimeter constant is assumed to be zero, because it's too small to make an effect.
2. Do you think the calorimeter constant in this experiment will affect your results for the later investigations? Why or why
not?
1. No. Because it's too small to cause an effect.
Analysis Questions Part 2
1. Write a balanced equation for the reaction of sodium hydroxide with hydrochloric acid. Calculate the number of moles of
each reagent you used in the reaction. Use the balanced equation to determine which reagent was limiting, and how
many moles of it reacted (show all your calculations for full credit). Was the acid or the base the limiting reagent?
. ( calculations are shown in the picture 1)
2. Calculate the value of the enthalpy change (ΔHrxn) in units of kJ/mol for the reaction of NaOH(aq) + HCl(aq). Is the
reaction an endothermic or exothermic process?
1. Exothermic. ( calculations are shown in the picture 1)
Analysis Questions Part 3
Discuss with your team (show your group data table) and the whole class (show class data table) and enter the answers of the following
questions in your ELN.
1. Calculate the value of the enthalpy change (ΔHdiss) in units of kJ/mol for dissolving both forms of the salt (show all your
calculations for full credit). Is dissolving an endothermic or exothermic
1. ( calculations are shown in the picture 1)
2. Write balanced reactions for both dissolving processes. ( calculations are shown in the picture 1)
a) How do the ΔHdiss values compare for the two forms of the salt?
1. One is enothermic with a small value, Another one is exothermic with a much
higher value.
b) Does the difference in ΔHdiss values tell you anything about how the two salts compare? For
example, which is the more stable form of the salt?
1. The one that's enthermoic is more stable, because it's a lower-energy compound that has
Week 8/ 6 of 8
Angie Burke Sep 30, 2017 @03:10 PM CDT
a smaller value of ΔHdiss.
c) Is there a way to determine how much heat you would need to add to get from one salt to the other?
1. Heat can be determined by subtracting the ΔHdiss from both reactions.
3. How does having waters of hydration affect the enthalpy for dissolving?
1. Water changes how stable the solution is.
Reflective Writing
Week 8/ 7 of 8
Angie Burke Oct 12, 2017 @08:43 PM CDT
Explain:
Investigation Questions:
1. What kind of reactions procedure (exothermic)/ absorbs (endothermic) heat?
1. Endothermic reaction absorbs heat. Exothermic reaction produces heat.
2. What is the sign of enthalpy for exothermic/ endothermic reactions?
1. exothermic - negative; endothermic - positive
3. What does the Hess's Law state?
1. No matter how many stages of steps of a reaction takes, the total enthalpy change is the sum of all changes.
4. Why is the Hess's Law useful?
1. It allows the enthalpy to be calculated even if it can't be measured directly.
Post Discussion Questions
q cal is approximately = 0: Why is this a good assumption?
Because it's not big enough to make any effect
Which q did you actually measure? qrxn or qsolution?
q reaction
Can you measure the other q? Why or why not? If yes, how?
No. Because only calculation can be done on q solution.
Why is the qrxn= - qsolution? Define what the system and the surroundings are for each experiment.
You can neither create or destory the energy. The water is the surrounding. The NaOH and HCl are the
system
How do volume and mass influence ΔH e q?
It affects how much compacity you have to measure the heat.
How did you decide which was the limiting reagent in each experiment?
Compare the moles of the reagents.
How did you applied the Hess’s Law?
We measured how much heat the system lost/ absorbed with different types of solutions, and how much
heat that's been absorbed/ lost by surrounding.
What type of errors might affect your results? (Please avoid mentioning “Human error” and try to be as specific as
possible)
There are water leftover, spills, an assumption we made on q (states that it doesn't matter).
Evaluate:
1. Do your answers to investigation questions make sense? Are they consistent with what is already known about chemistry? How do your
conclusions tie in with what you are learning in lecture?
Yes. They are consistent with what we have been learning. From lectures, we learned that endothermic reactions are positive and exothermic
reactions are negative.
2. Are your results consistent with your entire class? If not, why not? What would you do differently in order to get even better results? Saying
anything that has to do with human error will not receive any points.
They are not all consistent with class data. The reasons might due to water leftover, spills and inaccurate measurements that caused the errors
to happen.
Extend:
1. How does what you have learned in lab today tie in with anything outside your chemistry lab and lecture?
Exothermic and endothermic happens everywhere in normal life, such as cooking, light a candle, and freeze something in the refrigerator.
Week 8/ 8 of 8
[Show More]
