>> host: Welcome back
to Chemistry Matters!
By now, you performed
the coffee cup calorimetry lab
and carefully analyzed
your data, looking
for trends and patterns.
So based on your data,
you're prepared to address
the original experimental
challenge, to design
an effective hot pack
and cold pack.
Your conclusion,
which in this case,
will be a recommendation on
which ionic compounds to use,
should be supported
by your data.
With all of that in mind,
let's rejoin our classroom
so you can compare your findings
with those of our students.
>> professor: I'm looking
forward to seeing your results
and hearing your conclusions
from the calorimetry experiment.
Who'd like to present
their data first?
>> student: We will.
We transferred our data table
into a spreadsheet and used it
to make graphs.
>> student: We figured graphs
would let us see trends
more easily.
>> student: Our graphs showed
us that two of the ionic
compounds increased
the temperature of the water,
and two of them decreased
the temperature.
>> student: Calcium chloride
increased the temperature
the most, by 6.1 degrees.
Sodium carbonate increased
the temperature just a little,
by 1.4 degrees.
>> student: And potassium
chloride lowered the temperature
of water the most,
by 4.2 degrees.
Sodium bicarbonate cooled
the water down too,
but only by 1.1 degree.
>> professor: Great presentation
of your data, students.
The trends are easy
to see on your graph.
And it's also very helpful
that you combined
all four data sets
onto one graph.
So what conclusions did you
make based upon your data?
>> student: We concluded
that calcium chloride
would make the best hot pack
because it increased
the temperature of the water
the most, and that
potassium chloride would make
the best cold pack
because it cooled the best.
>> professor: And based on
your data, I agree
with your conclusion.
Nice job.
>> student: And we actually
made our own hand warmer!
So we don't have to spend
money on hand warmers anymore.
We can just put
calcium chloride and water
into plastic baggies!
>> professor: Well, that's
a great application,
just be sure that you're
very careful!
How about you two?
Are you ready to present
your results and conclusions?
>> student: Yes, we are,
Our results were really
close to theirs.
We calculated
the heat transferred
from each ionic compound using
the equation, q=mC delta T.
>> student: Right.
we found that dissolving
potassium chloride and
sodium bicarbonate
are endothermic,
they absorbed heat.
>> student: We worked it out
this way: q=mC delta T.
Since we're actually measuring
the temperature change
of the aqueous solution,
we set m equal to the mass
of the entire solution,
105 grams.
And since the solution
is mostly water, we used
the specific heat capacity of
water, which is 4.184 Joules
per gram degree Celsius.
And we measured delta T,
the change in temperature,
from the starting point
to the biggest change
in temperature.
>> student: The solution
of potassium chloride
went down in temperature
by 4.1 degrees Celsius,
which means the solution
actually got colder!
So the potassium chloride
absorbed heat from the water
when it dissolved.
We did the calculation and
found that potassium chloride
actually absorbed 1,799 Joules
of heat when it dissolved.
>> student: The same process
showed us that
sodium bicarbonate absorbed
1,097 Joules of heat.
>> student: When calcium
chloride and sodium carbonate
dissolved, the water warmed up.
So those are exothermic
processes.
>> student: Calcium chloride
released 2,897 Joules of heat
and sodium carbonate
released 1,536 Joules.
>> professor: Excellent job,
students,
You used your data to
calculate energy absorbed
or released.
If we were developing
a commercial product, we would
need this information.
What are your conclusions
based on your data?
>> student: We had the same
conclusions as them.
Calcium chloride produced
more heat than sodium carbonate,
so we chose it to make
our hot pack.
>> student: And potassium
chloride absorbed more heat
than sodium bicarbonate,
so it's our choice
for the cold pack.
>> professor: Great job
by both groups.
All of your data was presented
in a way that made it
easy to detect the trends
and draw conclusions.
And both groups' conclusions
were based directly upon data.
All in all, I'd say this was
a very successful experiment!
>> host: I hope this engineering
and design challenge
went as well for you
as it did for our students.
If you'd like,
you can even determine
what it would cost you
to make hot and cold packs
based on the current cost
of the materials.
We covered a lot in this unit
on thermodynamics.
But if want to further explore
thermodynamics and energy
in reactions,
you can check out
our Closer Look video on
Graphing Gibbs Free Energy,
which combines enthalpy and
entropy into a single value.
You'll be one step closer
to becoming a professional
chemical engineer!
And with that in mind,
you can move on to Unit 9,
Kinetic Energy and Gases.
I'll see you there!