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### Course: Physical Chemistry (Essentials) - Class 11 > Unit 6

Lesson 4: Measurement of U and H- Calorimetry# Constant-pressure calorimetry

Constant-pressure calorimetry is used to measure the change in enthalpy, Δ

*H*, for a physical or chemical process. In this technique, a process is carried out in solution in a coffee cup calorimeter, an inexpensive device composed of two Styrofoam cups. The amount of heat transferred in the process (*q*) can be calculated from the mass, specific heat, and temperature change of the solution. Because the calorimeter is at constant (atmospheric) pressure,*q*is equal to Δ*H*for the process. Created by Jay.## Want to join the conversation?

- how did they come up with the name calorimeter?(2 votes)
- “Calor” is Latin for heat, and “metry” is Greek for measurement. So it literally means heat measurements.(5 votes)

- At7:29, due to the law of conservation of energy,should the energy in the system stay the same? Assuming that the system is reffering to the block and the water, the thermal energy in both objects should be the same before and after equillibrium, right?(2 votes)
- So the law of conservations of energy states that the energy of the universe is constant. And when we're talking about thermodynamics the universe is the system and the surroundings combined. So energy can leave the system and be transferred to the surroundings. So individually the system loses energy and the surroundings gain energy, but together the energy content is constant because they both constitute the universe.

In a chemistry context, the system refers to the reaction; specifically just the chemicals reacting. The water enveloping the block (the reaction) would be part of the surroundings. So the energies (and therefore temperatures) of the reaction and water (and technically the rest of the surroundings) will change over the course of the reaction, and this is still allowed under the law of conservation of energy because the sum of their energies remained constant.

Hope that helps.(5 votes)

## Video transcript

- [Instructor] Calorimetry
refers to the measurement of heat flow. And a device that's used
to measure heat flow is called a calorimeter. An easy way to make a calorimeter
is to use two coffee cups. So at the base here,
we have one coffee cup, and then we can also
use another coffee cup as a loose fitting lid. And since this top coffee
cup is loose fitting, our calorimeter is exposed
to the constant pressure of the atmosphere. Therefore, we could use
this coffee cup calorimeter for constant-pressure calorimetry. Other components of our
calorimeter include some water and then we also have a stir
bar to stir up the water, and a thermometer to measure
the temperature change of the water. Let's say we have 150.0 grams of water at an initial temperature
of 25.0 degrees Celsius. Next, let's take a block of
copper, 120.0 grams of it, and let's heat up that block of copper to 100.0 degrees Celsius. Once the copper has
reached that temperature, we add the copper block
to our calorimeter. Here, we can see the
copper block has been added to the calorimeter, and since the copper is at a higher temperature than the water, heat flows from the
copper block to the water, and therefore the temperature
of the water will increase, which we will see on the thermometer. So we'll see the temperature
increase on the thermometer. Heat is transferred from the
copper block to the water until thermal equilibrium
has been reached. And we know when thermal
equilibrium has been reached by looking at the thermometer and measuring the highest
temperature that's reached. Let's say the final temperature is equal to 30.0 degrees Celsius. So at thermal equilibrium,
both the pieces of copper, both the copper block and the water are at the same final temperature. Next, let's calculate the
heat gained by the water by using the equation q
is equal to mc delta T. So q is what we're trying to calculate, the heat gained by the water. M is the mass of the water,
which is 150.0 grams. So we can write in 150.0 grams. C is the specific heat of
water, which is 4.18 joules per gram degrees Celsius. And delta T is the change
in the temperature, which would be the final temperature, so Tf minus the initial temperature Ti. The final temperature of the
water is 30.0 degrees Celsius, and the initial temperature of the water was 25.0 degrees Celsius. So 30.0 minus 25.0 is equal
to 5.0 degrees Celsius. So we can write that in. And next, we look at units
and see what cancels out here. So the grams cancel out,
degrees Celsius cancels out, and we're left with joules as our unit. So q is equal to, when we go
to two significant figures, this is positive 3.1 times
10 to the third joules. The positive sign means that
this was the energy gained by the water. Next, let's do the same
calculation for copper. So we're trying to find q. The mass of the copper was 120.0 grams. So we can plug that in. The specific heat of copper is .39 joules per gram degrees Celsius. And let's think about the change in the temperature of the copper. The final temperature of the copper was 30.0 degrees Celsius, and the initial temperature of the copper was 100.0 degrees Celsius. So the change in the temperature
would be 30.0 minus 100.0, which of course is negative 70.0. So let's plug in negative
70.0 degrees Celsius. Once again, we see what
cancels for our units. Grams will cancel, degrees
Celsius will cancel, and our answer will be in joules. So q is equal to, using
two significant figures, negative 3.3 times 10 to the third joules. And the negative sign, so this negative sign
means this is the energy that was lost by the copper. Next, let's look at these
two numbers that we got from our calculations. Let's think about the
magnitude of these two numbers. If our coffee cup calorimeter
were a perfect insulator, the magnitude of these two
numbers would be the same. So it could be something like 3.3 times 10 to the third joules for both of them. But since these two numbers
are not the same, right? We can see that we've lost
more heat from the copper than we've gained in terms
of energy for the water, which means we could have
lost some of the energy to the environment. So not all of the heat
was transferred directly to the water. Some of it could have escaped
our coffee cup calorimeter. Next, let's think about calorimetry
for a chemical reaction. So before we do that, let's review some terms
for thermodynamics. So the system is the part of the universe that we are studying. So in the case of a chemical reaction, the reactants and the
products make up the system. The surroundings are everything else, which would include the
water in the calorimeter, the coffee cup itself, the thermometer, the environment outside, so the surroundings are everything else. And finally, the universe
would be the system plus the surroundings. So the reactants and the
products make up the system. So that's what the S stands
for here in our calorimeter. That's our system. And let's say we run a reaction, and in the reaction, heat is given off. So in that case, heat would flow from the
system to the surroundings, and so the temperature of
the water would increase. So we would see that as
the temperature increases on the thermometer. Next, we could calculate
the heat gained by the water by using our q is equal
to mc delta T equation. And let's say q is equal
to positive 1.0 times 10 to the second joules. The positive sign means that
the water gained energy. If we assume a perfect transfer of heat from the system to the surroundings, if the surroundings gained
positive 1.0 times 10 to the second joules, that means the system must
have lost negative 1.0 times 10 to the second joules. So the same magnitude, but
we changed the sign here, because if we're talking about the energy lost by the system, it's the same in magnitude,
but opposite in sign. Next, remember that our lid
over here is loose fitting, which makes this constant
pressure calorimetry, and therefore this heat that
was transferred is the heat that's transferred at constant pressure. So we can write a
subscript p in here, so qp. The heat transfer at a constant
pressure is the definition for the change in the enthalpy delta H, so we can write that
qp is equal to delta H. And when delta H is negative, we're talking about an
exothermic reaction. So when a reaction is exothermic, heat is transferred from the
system to the surroundings, and therefore we see an increase in the temperature of the water. Finally, let's think about
an endothermic reaction. In an endothermic reaction, heat is transferred from the
surroundings to the system. So here we can show heat flowing from the surroundings to the system. Since energy is leaving the surroundings, the temperature of the water will decrease for an endothermic reaction. And since heat is being
transferred to the system, we can go ahead and write heat over here, we can go ahead and write
heat on the reactants side, and delta H would be positive
for an endothermic reaction. So for an endothermic reaction, energy is transferred from the
surroundings to the system, and therefore the temperature
of the water will decrease.