Chapter 8: Chemical Reactions

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Welcome to the deep dive.

We're the place where we take a whole stack of information and, well, pull out the key insights just for you.

That's the goal.

Today, we are diving deep into the pretty essential world of gases.

We're drawing from the textbook chemistry, an introduction to general organic and biological chemistry.

A solid foundation.

Right.

And our mission today is to really get a grip on how gases behave, the basic rules, but maybe more importantly, see how vital this stuff is in, you know, health and medicine.

Oh, absolutely.

It's fundamental chemistry.

But the cool part is how directly these, well, seemingly abstract laws actually connect to our everyday lives.

Like what?

Like breathing, for starters,

medical treatments, safety things.

Yeah.

Yeah.

We'll uncover some surprising connections.

I think it really shows why understanding this is so key.

Okay, let's make it real then.

To kick us off, picture this.

Whitney, young soccer player.

Right.

After practice, she suddenly finds it really hard to breathe.

Chest feels tight.

Scary situation.

Yeah.

Her dad rushes her to the ER.

Respiratory therapist checks her out and boom, asthma attack.

And this right here is where our dive into gases really starts because understanding what happened to Whitney and how they treated her, it all comes down to the principles we're about to unpack.

Exactly.

The treatment directly involves manipulating how gases behave in her airways.

So let's start with the basics.

The air around us, we're basically swimming in this sea of gases, the atmosphere.

Right.

Living at the bottom of it.

And it's mostly nitrogen, what, 78 %?

Yeah, about that.

Then the really crucial part, 21 % oxygen.

The stuff we need.

The stuff we definitely need.

And then there's this tiny 1 % leftover argon, CO2, water vapor, bits and pieces.

Okay, so that's what air is.

But what makes a gas, well, a gas?

How's it different from a liquid or a solid block?

Good question.

It comes down to something called the kinetic molecular theory of gases.

It's kind of the rule book for gas behavior.

Okay, break it down for us.

Kinetic molecular theory.

Right.

So first, think of gas particles,

atoms, molecules, whatever they are, as tiny things zipping around randomly, really fast, high velocity.

Bouncing everywhere.

Pretty much constantly moving, which is why a gas always, always fills up whatever container you put it in.

It just expands.

Okay, point one, random fast movement.

What else?

Point two,

the forces between these particles, usually super, super small,

almost negligible, especially compared to liquids or solids where particles stick together more.

Ah, so they don't really attract each other much.

Not much at all, which leads to point three.

The actual space the particles themselves take up is tiny compared to the total space the gas fills.

Imagine a few bees buzzing around in a huge cathedral.

Okay, lots of empty space.

Exactly.

And that's why you can compress gases so easily.

You're just squeezing out that empty space.

Try doing that with water.

Right, doesn't work so well.

So they move fast, don't stick together, take up little space.

What else?

Point four,

they're in constant motion, moving in straight lines until they smack into something, another particle or the container wall.

Okay.

And those collisions with the walls, that's what we feel is pressure.

More collisions or harder collisions mean higher pressure.

Ah, okay, the force of all those little impacts adds up.

You got it.

And finally, point five, this is really key.

The average kinetic energy,

the energy of motion of these particles is directly proportional to the temperature.

But, and this is important, the Kelvin temperature.

Pelvin, not Celsius or Fahrenheit.

Always Kelvin for gas logs because Kelvin is an absolute scale.

Zero Kelvin means zero energy.

So heat things up, temperature rises, particles move faster, hit the walls harder and more often.

Then the pressure goes up, makes sense.

Exactly.

So putting it all together, this theory explains, well, why perfume spreads across a room?

Those fast moving particles just diffuse out.

Yep.

Or why a tire feels harder on a hot day, the air inside heats up, particles move faster, pressure increases, or like you said earlier, the danger of heating a sealed can.

Right.

That pressure buildup could be huge.

Okay.

So when we talk about or measure a gas, what are the key things we need to know?

There are four fundamental properties we focus on.

Pressure, which we just talked about, the force on the walls.

We may have written atmospheres,

or millimeters of mercury, M -A -H -G, sometimes C -I like in your tires.

Okay.

Pressure P.

What's next?

Volume V.

That's just the space the gas takes up.

And for a gas, that's always the volume of its container.

Usually liters, L or milliliters, temperature, temperature T.

Related to that kinetic energy, how fast the particles are moving.

And remember,

always Kelvin for the calculations.

Higher temperature, faster particles.

Okay.

P, V, T, one more.

The amount of gas.

N.

Basically, how many gas particles are in there?

We usually measure this in moles.

It's just a standard scientific counting unit for a huge number of particles.

Add more gas, like pumping air into a tire.

You increase the number of particles hitting the walls, so the pressure goes up.

Got it.

Precisely.

Those are the big four.

P, V, T, and N.

And you know, this connects directly back to and releases on your arm.

That's phygmomanometer is measuring the pressure your blood exerts on your artery walls, usually in the M -E -H -G.

Cystolic over diastolic, like 120 over 80.

Exactly.

That's a direct application of pressure measurement, just like measuring gas pressure.

And knowing those numbers is vital.

Too high, risk of stroke, heart attack, too low, maybe not enough oxygen getting around.

It's pure physics and chemistry applied to physiology.

Amazing how fundamental it is.

Okay, so we have these properties.

Now let's get into the specific relationships between them.

The gas laws.

Where should we start?

Maybe with breathing again.

Perfect place to start.

That brings us to Boyle's law.

This one describes the relationship between pressure and volume.

Okay.

Boyle found that if you keep the temperature and the amount of gas the same, pressure and volume are inversely related.

Inverse, meaning opposite.

Exactly.

Squeeze the volume down, make the container smaller, and the pressure inside goes up.

Give the gas more room, increase the volume, and the pressure goes down.

Why does that happen again?

Think about it.

Same number of particles, same speed, because the temperature is constant.

If you cram them into a smaller space, they're just going to hit the walls more often.

More collisions per second means higher pressure.

Makes sense.

Like pushing the plunger on a sealed syringe, volume down, pressure up.

Perfect example.

And this is exactly how breathing works.

It's Boyle's law running your respiratory system.

How so?

Walk us through inhaling.

Inspiration breathing in.

Your diaphragm muscle contracts, pulls down.

Your rib cage muscles pull your ribs up and out.

So the chest cavity gets bigger.

Right.

The volume inside your chest, including your lungs, increases.

Now Boyle's law,

volume goes up.

So what happens to the pressure inside your lungs?

It goes down.

It goes down.

It drops just slightly below the atmospheric pressure outside your body.

Ah, and air flows from high pressure to low pressure.

Precisely.

That tiny difference is enough to suck air into your lungs.

It flows down the pressure gradient.

Wow.

Okay.

So exhaling must be the opposite.

Expiration.

Yep.

Your diaphragm relaxes, moves up.

Ribs move down and then the volume of your chest cavity decreases.

So Boyle's law again,

volume down, pressure.

Pressure inside the lungs goes slightly above atmospheric pressure and the air gets pushed out.

Simple as that.

Well, mechanically simple, biologically complex.

So back to Whitney and her asthma attack.

Her airways were constricted tight.

Right.

That inflammation and constriction made it hard for her chest to expand properly and crucially hard for air to flow easily.

She couldn't increase her lung volume enough.

Which means she couldn't decrease the pressure enough to draw in a full breath.

She's fighting against Boyle's law.

Exactly.

She couldn't create that necessary pressure gradient.

So the respiratory therapist gave her a bronchodilator medication.

To open the airways?

Yep.

Relax those constricted muscles, widen the passages.

That allowed her lungs to expand more easily,

decrease the internal pressure and let the air flow in again.

A direct medical intervention based on understanding P and V.

That also explains the scuba diver thing you mentioned.

Exhaling while coming up.

Absolutely critical.

As a diver ascends, the water pressure around them decreases dramatically.

If they hold their breath, the air already in their lungs, according to Boyle's law, will try to expand as the external pressure drops.

Oh, wow.

So the volume has to increase.

And if it can't escape by exhaling, that expansion can literally rupture lung tissue.

It's called pulmonary barotrauma.

So rule number one of diving.

Never hold your breath on ascent.

Always breathe normally.

Boyle's law is literally a lifesaver there.

Okay, what's next?

We've linked P and V.

What about temperature?

Let's picture a hot air balloon.

Good visual.

That brings us to Charles's law.

This one connects volume and temperature.

Okay.

Charles found that if you keep the pressure and amount of gas constant, the volume of gas is directly related to its Kelvin temperature.

Directly related, meaning they go up or down.

Heat the gas up, increase T, its volume increases, cool it down, decrease T, its volume decreases.

What?

Back to the particles.

Yep.

Heat them up, they move faster, hit the walls harder and more often.

To keep the pressure constant, which is the condition for Charles's law, the volume has to expand to give them more room.

Okay.

So in the hot air balloon, they heat the air inside.

The air expands, its volume increases.

Since it's taking up more space, but has the same amount of air, it becomes less dense than the cooler air outside.

And lighter things float.

Less dense things float.

So the balloon lifts off.

Simple application of Charles's law.

Cool.

Any medical examples for Charles's law?

Well, it's relevant anytime gases are used where temperature might change.

Think about sterilizing equipment.

If you heat sealed containers with gases inside, you need to account for volume changes.

Or in surgery, like laparoscopy, where they inflate the abdomen.

You mentioned helium for that.

Sometimes, yeah.

If they fill the space with a certain volume of gas at room temperature, and then maybe the patient's body heat or operating room lights warm that gas up.

Its volume will increase.

Right.

They need to understand that relationship.

Say you have 5 .4 liters of helium at 15 degrees C.

If it warms up to 42 C, which isn't unrealistic, Charles's law tells us it will expand to almost 6 liters.

You need to account for that.

So it's about predicting behavior under changing conditions.

Got it.

Now, what about the link between temperature and pressure?

You mentioned the dangerous scenario.

An oxygen tank in a fire.

Right.

That perfectly illustrates Gay -Lussac's law.

This law states that pressure is directly related to Kelvin temperature if the volume and amount of gas are constant.

So like Charles's law, they go up and down together.

Higher T, higher P.

Exactly.

If you heat a gas in a fixed container, so volume can't change, those faster moving particles just slam into the walls much harder and more frequently.

Pressure skyrockets.

It does.

Think about that oxygen tank.

Maybe it's rated for, say, 120 atmospheres at room temp, 25 degrees C.

If a fire heats it up to, let's say, 400 degrees C.

Which is really hot.

Very hot.

According to Gay -Lussac's law, the pressure inside could jump to

maybe 270 atmospheres or even more.

Yeah.

Wow.

And if the tank can only handle, say, 108 adiometer.

Ooh.

Catastrophic failure.

It ruptures.

That's why there are incredibly strict rules about storing compressed gas cylinders away from heat sources in hospitals, labs, homes where people use oxygen.

It's a direct safety application of Gay -Lussac's law.

Understanding that direct relationship is critical for safety.

Okay.

So we have Boyle, P &V, Charles, Gay -Lussac, P &T.

Seems like they're all connected.

They are.

And that leads us to the combined gas law.

It basically rolls Boyle's, Charles' and Gay -Lussac's laws into one single equation.

Handy.

Very.

It looks like this.

P1V1T1 includes P2V2T2.

It lets you calculate what happens when multiple conditions change at once as long as the amount of gas stays the same.

So if pressure and temperature change, you can figure out the new volume.

Exactly.

Let's go back to that diver's bubble.

Imagine a tiny 25 millilo bubble release from their tank deep down where the pressure is high, maybe four atmospheres, and the water is cold, say, 11 degrees C.

Okay.

High P, low T.

Now that bubble rises to the surface.

What happens?

Well, the pressure drops way down, maybe to one atmosphere.

And the surface water is warmer, say, 18 degrees C.

Right.

So P decreased, T increased.

Both will affect the volume.

Using the combined gas law, you plug in the initial P, V, and T, remember to convert T to Kelvin, and the final P and T.

And it tells you the final volume.

Yep.

In this case, that little 25 millilo bubble would expand to over 100 milliloL by the time it hits the surface.

The combined gas law lets us predict that.

Crucial for understanding decompression effects.

That's a powerful tool for situations where everything's changing.

Okay, one property left.

The amount of gas.

And what if that changes, like blowing up a balloon?

Now we're talking about Avogadro's law, named after Amadeo Avogadro.

This one's pretty intuitive, too.

It says that the volume of a gas is directly related to the number of moles of gas if the pressure and temperature stay constant.

So more gas, more volume.

Simple as that.

Double the moles of gas, you double the volume.

Triple the moles, triple the volume.

That's why the balloon gets bigger as you blow more air, more moles of gas into it.

Or pumping up a bike tire, you're adding more moles of air, increasing the volume inside until the pressure gets too high.

Exactly.

And Avogadro's law leads to a really useful benchmark, STP.

You mentioned that standard temperature and pressure.

Right.

It's a reference point scientists agreed on.

Exactly zero degrees Celsius, which is 273 Kelvin,

and one atmosphere of pressure.

Okay.

Standard conditions.

Why is that useful?

Because Avogadro figured out something amazing.

At STP,

one mole of any ideal gas occupies the exact same volume.

Any gas?

Does it matter if it's helium or oxygen or carbon dioxide?

It doesn't matter.

One mole of any of them at STP takes up 22 .4 liters.

That's called the molar volume.

22 .4 liters.

How big is that?

Think about three standard basketballs.

Roughly that much space.

Wow.

So that number, 22 .4 limol at STP, lets you easily convert between moles of gas and volume.

Absolutely.

If you know you have, say, two moles of nitrogen gas at STP, you know it occupies 44 .8 liters.

Or if you measure 11 .2 liters of oxygen at STP, you know you have half a mole.

It's a super helpful conversion factor.

Very neat.

Okay, last scenario then.

What if you don't have just one gas, but a mixture, like air?

Excellent question.

That brings us to our final law for today,

Dalton's law of partial pressures.

Partial pressures.

Yeah.

Dalton figured out that in a mixture of gases, each gas behaves as if it were the only gas present.

It exerts its own pressure independent of the others.

So each gas contributes a piece of the total pressure.

Exactly.

And the total pressure of the mixture is simply the sum of all those individual pressures, the partial pressures.

So the air pressure we feel right now.

Isn't just one thing.

It's the partial pressure of nitrogen PLUS, the partial pressure of oxygen PLUS, the partial pressure of argon, CO2, water vapor, all added together.

So if air is 78 % nitrogen, does nitrogen contribute 78 % of the pressure?

Pretty much, yeah.

At sea level, total pressure is about 760 mmHg.

Nitrogen's partial pressure is around 594 mmHg.

Oxygen's is about 160 mmHg, and the others are tiny.

Add them up, you get 760.

Okay, that makes sense.

How is this important medically?

You mentioned hyperbaric chambers.

Right.

This is a fantastic application of Dalton's law.

In a hyperbaric chamber, they increase the total pressure, often to two or three times normal atmospheric pressure.

Why do that?

Well, think about the air inside.

It's usually close to 100 % oxygen, or at least a very high percentage.

If you increase the total pressure, you dramatically increase the partial pressure of oxygen, according to Dalton's law.

Ah, so much higher pressure of just O2 pushing on the body.

Exactly.

And that high partial pressure forces much more oxygen than usual to dissolve into the blood plasma and body tissues.

What does that extra dissolved oxygen do?

It's amazing for certain conditions.

For severe burns or infections with certain bacteria, anaerobic ones that hate oxygen, this oxygen flood helps fight the infection and promotes healing for carbon monoxide poisoning.

Where CO binds to your hemoglobin instead of oxygen.

Right.

The massively increased partial pressure of oxygen helps to displace that CO from the hemoglobin, essentially pushing it off so oxygen can bind again.

It can be life -saving.

Wow.

But you also mentioned a danger related to this, the bends.

Yes.

The flip side of messing with partial pressures and dissolved gases.

This affects scuba divers primarily, but also relates to leaving a hyperbaric chamber.

How does it work?

When a diver is deep underwater,

the high pressure forces more nitrogen from the air they're breathing to dissolve into their blood and tissues, just like the oxygen in the hyperbaric chamber.

But this time it's nitrogen.

Okay, so they have extra nitrogen dissolved in them.

Right.

Now, if they ascend too quickly, the external pressure drops fast.

That dissolved nitrogen suddenly wants to come out of solution back into gas form.

Like opening a soda bottle, bubbles appear.

Exactly like that.

But these bubbles form inside the diver's blood vessels and tissues.

Very painful and can be very dangerous, causing blockages, tissue damage.

That's decompression sickness or the bends.

Ouch.

So how do they prevent it?

Slow control descent.

Coming up gradually, often with safety stops at certain depths, allows that excess dissolved nitrogen time to diffuse slowly out of the tissues, into the blood, travel to the lungs, and be exhaled safely.

It keeps the nitrogen from forming bubbles.

So managing dissolved gases based on partial pressures is critical for diver safety.

Absolutely fundamental.

It's all about understanding Dalton's law and how gases behave under pressure.

This has been fascinating.

So really, what's the big takeaway for everyone listening?

I think it's realizing that these gas laws aren't just abstract equations in a textbook.

They are constantly, silently operating all around us and inside us.

Yeah, from the simple act of breathing controlled by Boyle's law.

To the medical gases kept safe by understanding Gaylew Sachs, the treatments enabled by Dalton's law in hyperbaric chambers, even the technology of a hot air balloon using Charles's law.

It truly connects the microscopic behavior of particles to macroscopic things we see and experience, especially in health.

Definitely.

Understanding these basics gives you a much deeper appreciation for how our bodies work, how medicine intervenes, and even how to stay safe in certain environments.

It's practical science.

So maybe next time you take a breath, or inflate a tire, or see a weather report mention barometric pressure,

take a second.

Think about those invisible gas particles buzzing around following these laws we talked about.

Yeah, how are these principles shaping that moment?

And maybe consider what else in your daily life could be explained by these fundamental rules of gas behavior, something to mull over.

Last minute lecture team.

On that note, a really warm thank you from all of us here at the last minute lecture team for joining us on this deep dive into the world of gases.

We hope you've gained a clearer, maybe more practical perspective on this bit of chemistry.

Thanks for listening.