Calorimetry: Crash Course Chemistry #19

Calorimetry: Crash Course Chemistry #19

October 3, 2019 100 By Ewald Bahringer


Hydrochloric Acid: every chemist’s frenemy,
as terribly dangerous as it is terribly useful. It’ll burn your skin, your eyes, even your mucus membranes if you breathe in its fumes for too long. But HCL as an acid gives up its hydrogen
pretty easily, which makes it good for making things like
fertilizers and dyes and even table salt. Then, there’s sodium hydroxide, another substance
that I wouldn’t wish to be on my worst foe, although I’m glad we have it. You may know it as lye, an extremely caustic
substance that’s used for everything from clearing clogged pipes to
purifying drinking water. It’s a base. It readily accept the protons
that acids release. So what do you think will happen when I mix
solutions of these two things together? Will they just cancel each other out and do nothing, or will they explode, or maybe they’ll travel through time? Well, if you’ve been paying attention, you
already know what’s going to happen. They’re going to undergo a neutralization
reaction, which we’ve talked about before. These two potentially deadly substances will
form harmless salt and water. But the reaction will also have an effect
that you can actually feel. It will release heat, and not just a little
heat. Mixing concentrated acids and bases releases
so much heat that it can result in an explosion. But I will show you how to produce a safe,
but noticeable amount of heat with this reaction. To me, the coolest part of this is where the
heat actually comes from. The energy used to exist as part of chemical
bonds in the acid and the base. Just like a ball at the top of a hill, the molecules always move towards a lower energy state if they can, and that’s just what they’ll do. High energy bonds will break and lower energy
bonds will form. The change in energy between those states you can actually feel the effects of, and that’s pretty dang cool. And what’s even more awesome, if you ask me, is that we can actually figure out exactly
how much heat will be released by this reaction. [Theme Music] Remember that measuring heat change is closely
related to enthalpy, which we defined as the internal energy of
a system plus the energy it uses to push the surroundings back and make room for its own pressure and volume. And in a constant pressure, like we have here
at the surface of the Earth, that works out to be exactly the same as the
heat that’s absorbed or released by a reaction. Naturally, it can be very useful to know how
much heat a chemical reaction absorbs or releases. In addition to the exothermic hand-warmers
that we have out there, there are also endothermic chemical ice packs
for treating injuries. The ability to calculate change in enthalpy
is also what tells pilots how far the fuel in an airplane’s tank will allow it to fly, which I personally am very interested in making
sure they get right. One of the ways we can calculate the change
in enthalpy of a system is with Hess’s Law, which you’ll recall states that the total
enthalpy change for a chemical reaction doesn’t depend on the pathway it takes, but only on its initial and final states. It’s often expressed in terms of Standard
Enthalpy of Formation, that is, the amount of heat lost or gained when one mole of a compound is formed from its elements. That’s how we figured out exactly how much
heat my hand-warmers release. But that’s not the only way that Hess’s Law
can be used. The law itself says nothing about the standard
enthalpy of formation. Any way that we can figure out the change
of heat between the products and the reactants will work just as well, and that’s where calorimetry
comes in. Calorimetry is the science of measuring the change in heat associated with a chemical reaction. And this may look like a plastic bottle inside
a koozie, but it’s actually a calorimeter. A calorimeter can be fancy and an expensive
piece of hardware, or it can be simple. But no matter what it looks like, it’s basically
an insulated container that contains a thermometer. And it can be made out of stainless steel
or Styrofoam cups, but there really are no fundamental differences
in how they work. And you know the general setup by now: the chemicals in the calorimeter make up the
thermodynamic system and everything else is the surroundings. The insulation minimizes the amount of heat
that leaks in or out of the system, so that we can be fairly confident that any heat transfer is part of the system, not the surroundings. The thermometer tracks the temperature change, which is part of the calculation we have to do. And there’s usually some way to stir the solution
to make sure that the reaction occurs fully. Alright everybody, safety first, though I
really should be wearing gloves… I’m gonna put 100 mL, also 100 grams, of HCL’s one mole of HCL solution into my calorimeter here… (mumbling) And now I’m going to put the same amount of
sodium hydroxide solution. Before I do the reaction, I have to know our
starting temperature, so I’m going to stick my thermometer in there
and wait for a second to see what it does. It should be room temperature, it’s been in
the room for a long time. So, we are currently at, like, 20.8 degrees
Celsius. So that’s like, 294 Kelvin, and now I shall
add my 100 mL of sodium hydroxide. The temperature, unsurprisingly, is rising
very rapidly. And I’m doing something right now that you should never ever do, which is stir with the thermometer, because if this happens in schools across
the world, then there will be a million billion broken thermometers and the stuff inside these thermometers is not good. So never do what I’m doing. All right, the temperature should be stable
by now, we have 28 point like 2 degrees Celsius. Now there’s a simple formula that allows us
to calculate the heat change of a reaction, simply by measuring the change in temperature
that occurs in a calorimeter. It says that the change in heat equals the
specific heat capacity of the substance times its total mass times the change in temperature. Let’s examine the parts of this. First of all, the heat change in the calorimeter
is normally represented by a lowercase “q,” but it can also be represented by change in
enthalpy, or delta H, because remember that constant pressure (delta
H) equals q, and constant pressure is almost always a good
assumption for the duration of an experiment, or at least as long as we stay at the surface
of the earth. For reasons that will become clear later, we’ll sure delta H to represent the heat change for this experiment. Specific heat capacity, represented by
a lowercase “s”, is the amount of heat required tp raise the
temperature of one mass unit, like a gram or kilogram, of a substance by
1 degree Celsius. So it turns out that different amounts of
heat create different temperature changes, like metals get hot really easily and cool
down really easily. Others like water require a lot of thermal
energy to raise the temperature, and therefore have to release a lot of heat
to cool down. I’m always wondering though, like, what does
that really mean? Like, physically in the molecules, shouldn’t heat raise the temperature of all substances equally? And why does water in particular have such
a high specific heat capacity? Heat energy can do a lot of things besides
just increase temperatures. Temperature, or the speed at which molecules
bounce around, is just one way that atoms or molecules can
absorb energy. Heat energy can also be absorbed by the breaking
and formation of bonds between molecules, and as we’ll learn in another episode, the
extremely high specific heat capacity of water is due to the breaking and formation of hydrogen bonds that are associated with relatively small changes in temperature. And how do we know the specific heat capacity? Well, I am happy to report that some noble
chemists have worked hard to determine the specific heat capacities of hundreds of substances
so that we don’t have to. We just have to look up the numbers in a table. Okay, so specific heat capacity times mass
times the change in temperature. The mass is important because the more mass of a substance we have, the more chemical bonds are present, and because energy is contained in chemical
bonds, they have a big effect on how much energy
we’re able to absorb or release. And finally, there’s the change in temperature. When doing calorimetry, we calculate a change
in heat by measuring a change in temperature, but as we’ve said a billion times before,
heat and temperature are not the same thing. But please do not think that this thing is
measuring heat because it’s not! It’s just that luckily, in this specific case, they are related by our handy little calorimeter formula. Now you might not have noticed, but we are right at the interface between chemistry and physics here. Each science could claim ownership over this
phenomenon, but the truth is humans made up the difference
between chemistry and physics anyway. Thermodynamics, the study of heat, energy,
and work, doesn’t care about our little rules. Thermodynamics itself makes the rules of the
universe. It is the ultimate law. So now you know, even though you might not
have cared, but you should! Because it’s good! It’s all wiggly-wobbly
bondy-wondy. All right! Enough talk, let’s get out there,
actually do some math here. Remember that the formula is delta H, s, m,
delta T. The solutions we’re using here are so dilute
that almost all of their mass consists of water. Therefore, we can use the specific heat capacity
of water. If we look that up on our table, we’ll see
that it is 4.184 Joules per gram degrees Celsius; I used 100 grams of each chemical for a total
mass of 200 grams. And finally, we need the temperature change. If you remember, the temperature rose from
294.0 Kelvin to 301.4 Kelvin; the difference between these two is 7.4 Kelvin. It’s a positive value because the temperature
increased. Cancel out all the appropriate units and then bang on the calculator to get a final release of 6192.32 J, or 6.2 kiloJoules of heat from the reaction. Because this formula is based on temperature
change, and since the temperature increased, we end
up with a positive result. But most importantly, it tells us the magnitude
of the change in heat energy. So, I wonder how that compares to the amount
we would predict using Hess’s Law and the standard enthalpy of formation? Remember that we can look up the standard
enthalpy of formation for all the products and reactants in the back of a textbook or
online, the chemical reaction between hydrochloric acid and sodium hydroxide produces liquid water and sodium chloride. The standard enthalpy of formation for hydrochloric
acid is -167.2 kJ per mole; for sodium hydroxide it’s -469.15 kJ per mole; for liquid water it’s -285.8;
and for sodium chloride it’s -407.27. I’m not gonna do the mole calculations on-screen, but trust me when I say that we used 0.100
mole of HCL and the amount of NaOH. Because everything in the equation balances
out, it’s just a 1:1:1:1 ratio, we can assume that they all have the same
amount of each product as well. If we plug these into Hess’s Law and do the
calculation, we found that the change in heat or enthalpy
of the reaction is -5.67 kJ. The system is releasing or losing energy,
so the number is negative, but again it’s really the magnitude that we
wanna know. So there you go, the calorimetry formula gave
an absolute enthalpy change of 6.2 kJ, while Hess’s Law gives a change of 5.67 kJ. So, why the difference? Well, the greatest factor is probably that we used the specific heat capacity of pure water instead of the salt water that we actually
created. We also didn’t include the heat capacity of
our calorimeter itself. The calorimeter walls and the thermometer
were heated too, resulting in some of the produced heat not
being accounted for. The insulation of the calorimeter is obviously
a bit light, which allowed some heat to escape entirely
and that’s another major factor. Even so, I’d say we did pretty well, the important thing is that it showed us what we needed to see, even though it was just a little plastic bottle
in a koozie. For a quick simple method, the calorimeter
got us pretty close to the calculated value. If we were calculating the amount of a particular
fuel we would need to travel to Mars, or inventing a cold pack that won’t give you
frostbite, we’d wanna use a more sophisticated system and work more carefully, but this was pretty cool for our purposes. Thanks for watching this episode of Crash
Course. If you paid attention, you learned that we
don’t necessarily have to use standard enthalpies of formation to solve Hess’s Law, you learned what a calorimeter is, that calorimetry is another way to investigate
heat changes in chemical reactions, and that specific heat capacity tells us how much heat energy affects the molecules
in a substance without changing its temperature. And finally, you learned some potential sources
of error related to calorimetry. The episode of Crash Course Chemistry was
written by Edi Gonzalez. The script was edited by Blake de Pastino and our chemistry consultant was Dr. Heiko Langner. It was filmed, edited, and directed by Nicholas
Jenkins. Our script supervisor was Caitlin Hofmeister,
and our sound designer is Michael Aranda. And, of course, our graphics team is
Thought Cafe.