Chapter 12: Solutions and Their Behaviour

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It is the evening of Thursday, August 21st, 1986, down in Lake Neos, Cameroon.

The sun goes down, the air is quiet,

and then suddenly 1 ,700 people and thousands of livestock just, they simply collapse and die overnight.

Yeah, it's horrifying.

And there's no earthquake.

There's no storm, no fire, literally no warning.

Just this quiet, deadly, invisible blanket that rolls out of the crater lake and suffocates everything in like a 26 kilometer radius.

It really remains one of the most chilling natural disasters in modern history.

And the culprit wasn't a pathogen or the volcanic eruption, at least not in the traditional sense.

The killer was carbon dioxide, just 202 or $2.

What happened at Lake Neos was this massive macroscopic catastrophe, but it was driven entirely by the microscopic laws of chemical solutions.

And that is why we're taking you on a deep dive into the hidden behavior of solutions today.

Yeah, welcome to this special session.

Think of this as your custom -tailored one -on -one tutoring session to help you completely conquer Chapter 12, solutions and their behavior from chemistry, human activity, chemical reactivity.

It's a great chapter.

It really is.

So we're going to cover the central chemical concepts in exact sequence today, from the thermodynamics of dissolving right through to colligative properties and colloids.

And you know, the water in our lakes, the oceans, the blood in your veins, the coffee you're probably drinking right now, they are all governed by a very specific set of physical rules.

And when a solution is pushed out of balance, the results can be, well, explosive, like at Lake Neos.

So let's look at the mechanics of Lake Neos.

This lake sits in a volcanic crater, right?

And deep beneath the earth, there's a magma chamber that's just constantly seeping carbon dioxide gas into the water.

Yeah, and under normal conditions at the surface, water cannot hold very much carbon dioxide.

Right, it just bubbles out.

Exactly.

But Lake Neos is 200 meters deep.

So at the bottom of that water column, the conditions are extreme.

You have near freezing temperatures and crushing pressure.

How much pressure are we talking about?

Around 20 bar,

which is 20 times the atmospheric pressure we experience at sea level.

Oh, wow.

So the heavy water on top is basically acting like a pressure cooker, trapping the gas down there.

That's a perfect way to look at it.

The physical weight of the water forces the CO202 molecules to stay dissolved.

So over centuries, the bottom layer of that lake absorbed millions of tons of gas until it became a completely saturated solution.

Saturated.

Right.

It reached what we call dynamic equilibrium.

I want to visualize that dynamic equilibrium for you, the listener, because at the macroscopic level, you look at the lake and it looks perfectly still.

But at the molecular level, it's total chaos.

You have carbon dioxide molecules dissolving into the liquid and other carbon dioxide molecules separating out of the liquid, but they're doing it at the exact same rate.

Yeah, exactly.

The rates of dissolution and exsolution are locked in a dead heat.

So the system is perfectly balanced.

Right.

There is no net change in the concentration.

Yeah.

But that balance is incredibly fragile because it depends entirely on that immense pressure staying constant.

Which means if something shifts that water, the balance is broken.

The lake is basically a giant pressurized sealed bottle of soda.

Yes.

You know, when you unscrew the cap on a bottle of soda, you hear that hiss, you're releasing the pressure and suddenly all those bubbles of gas come rushing out of the liquid.

And at Lake Neos, a trigger event basically caused the lake to open its cap.

What kind of trigger?

Well, it might have been a small underwater landslide or maybe a heavy rainstorm that cooled the surface water and made it sink or even a minor tremor.

We don't know for sure.

But whatever the mechanical trigger was, it pushed the layer of that deep, highly pressurized gas -soaked water upward.

And as that water rises toward the surface, the physical weight of the water column above it is decreasing, so the pressure drops rapidly.

Right.

And at that lower pressure, the water can no longer hold the gas.

In an instant, that saturated water becomes a super saturated solution.

Super saturated.

Yeah.

This is a highly unstable metastable state.

The solution contains a far higher concentration of dissolved sloot than the current pressure and temperature should theoretically allow.

So the gas just has to escape to restore equilibrium.

Exactly.

It rushes out violently, which displaces the water, creating this massive updraft, which then pulls even more deep water to the surface.

And you get a chain reaction.

Wow.

Yeah.

A 300 ,000 ton cloud of dense carbon dioxide gas just bursts out of the lake, pours over the crater rim, and silently flows down into the valleys.

The tragedy there is just, it's a terrifying demonstration of supersaturation, but we actually see this exact same chemical principle,

you know, a supersaturated solution rapidly returning to equilibrium in completely harmless everyday objects.

Oh, definitely.

Like those reusable hand warmers, you know, the little plastic pouches with the liquid and the metal disk inside.

I love those things.

Right.

You snap the disk and instantly the clear liquid turns into solid white crystals and the pouch gets incredibly hot.

Yeah.

So that heat pack contains a supersaturated solution of sodium acetate.

It was heated up to dissolve a massive amount of the salt and then allowed to cool very slowly.

Okay.

And because it cooled undisturbed, the salt molecules remain trapped in the liquid state.

They desperately want to crystallize, but they lack a nucleation site.

Like a starting point.

Right.

A starting point to begin the process.

So snapping the metal disk creates a tiny acoustic shock wave that provides the starting point.

Exactly.

The disturbance breaks that metastable state.

The excess sodium acetate rapidly crystallizes out of the solution until the liquid that remains drops back down to a normal saturated equilibrium.

But the crystallization process generates heat.

A lot of it, actually.

The pouch jumps up to about 50 degrees Celsius.

It does.

Which brings us to a really fascinating question about the invisible energy changes that happen when things dissolve or crystallize.

Because dissolving isn't just like matter disappearing into a liquid, right?

Not at all.

It is a massive transfer of thermal energy.

We measure this through the molar enthalpy change of solution.

The molar enthalpy change of solution.

Yes.

It represents the total thermal energy absorbed or released when one mole of a substance dissolves in a solvent.

And whether a solution gets hot or cold depends on this microscopic tug of war.

Okay.

Let's break down that tug of war for you.

If we take a solid crystal, let's say potassium fluoride, and we drop it into water,

two distinct things have to happen energetically.

First, you have to rip the solid crystal apart.

A solid crystal is this rigid geometric lattice of positively and negatively charged ions.

And they're bound tightly together by electrostatic forces.

So to dissolve it, you must input enough energy to overcome those bonds and pull the ions apart into isolated gaseous particles.

And that initial step is called the lattice enthalpy.

Yes.

And it is always endothermic.

It demands an input of energy.

I like to think of this as a business transaction, right?

Lattice enthalpy is your startup cost.

I like that analogy.

You have to spend money, or in this case energy, to build the factory and get the business off the ground.

It hurts your wallet initially.

It does.

But the second step is the payoff.

Once those bare ions are floating around, the water molecules rush in to surround them.

Okay.

Because water is polar, right?

It has a partial negative charge on the oxygen atom and a partial positive charge on the hydrogen atoms.

So the oxygen ends cling to the positive ions, and the hydrogen ends cling to the negative ions.

So they form a cage around the ions.

Exactly.

What?

This process is called equation, or hydration.

And forming these new ion dipole attractions releases a massive amount of energy.

It is highly exothermic.

So that's your business revenue.

The customers are paying you.

Yes.

If the revenue from equation is larger than the startup cost of the lattice enthalpy, your business turns a profit.

The overall process releases energy into the surroundings, and the beaker gets hot.

Sodium hydroxide is a classic example of that.

When you dissolve solid sodium hydroxide in water, the equation energy is significantly larger than the lattice energy.

Wow.

The solution heats up so rapidly it can actually boil the water if you aren't careful.

But wait.

What if the startup cost is simply too high?

Like what if pulling the crystal apart takes more energy than the water molecules can give back during equation?

Then the business operates at a loss.

The net process is endothermic.

To make up the energy deficit, the dissolving salt absorbs ambient heat from the water itself.

So the temperature drops.

Right.

The kinetic energy of the water molecules plummets, and the temperature of the solution drops.

That is exactly how an instant cold pack works for sports injuries.

You punch the bag, it breaks a little inner pouch of water, the water mixes with solid ammonium nitrate, and the solution turns freezing cold.

Because the lattice enthalpy of ammonium nitrate is so high that dissolving it literally sucks the heat right out of your skin.

It's amazing.

And you know, these energy dynamics dictate how solubility responds to outside forces like temperature and pressure.

Like we saw at Lake Nyos.

Exactly.

We already saw how pressure affects gases there.

That relationship is formalized by Henry's Law, which states that the solubility of a gas in a liquid is directly proportional to the partial pressure of that gas above the solution.

So more pressure pushing down on the surface forces more gas molecules into the liquid.

Exactly.

Now, temperature plays an equally critical role, but it often trips people up.

If you heat a liquid, gas has actually become less soluble.

See, I always pushed back on this when I first learned it, because my instinct is that heating water makes things dissolve better.

Like if I want to dissolve a massive scoop of sugar in my tea or salt in a pot of pasta water, I want that water boiling hot.

And your instinct is totally correct for most solids.

For the majority of ionic salts and sugars, solubility does increase as temperature rises.

The added thermal energy helps break apart the solid lattice.

Makes sense.

Although it is worth noting there are bizarre exceptions.

Like calcium sulfate actually becomes less soluble in hot water, which is why it precipitates out and forms hard, crusty scale inside industrial hot water boilers.

That's annoying.

But for gases, you're saying it's the complete opposite.

Hot water holds less gas, which actually makes sense visually if you think about it.

When you heat a pot of water on the stove, long before it actually reaches a rolling boil, you see tiny bubbles forming on the bottom and sides of the pot.

That is the dissolved air coming out of solution because the water is getting too hot to hold it.

Exactly.

And we can understand why by looking at Le Chatelier's principle.

This principle states that when a system at equilibrium is disturbed,

it adjusts to counteract the change.

Okay.

When a dissolved gas exits a solution and becomes a free gas, that specific process is endothermic.

It absorbs heat.

So if I take a beaker of water and put it on a hot plate, I am dumping a ton of excess heat into the system.

Right.

And the system wants to relieve that thermal stress.

According to Le Chatelier's principle, the system will shift in the direction that absorbs that extra heat.

Since the gas escaping into the atmosphere absorbs heat, the rising temperature literally drives the gas out of the water.

That is so elegant.

But tracking all these shifting solubilities requires precise mathematics, right?

It does.

We need a way to measure concentration accurately, whether we're at the freezing depths of a crater lake or inside a boiling industrial pipe.

And the standard unit most people learn first, molarity, which is moles of solute per liter of solution,

it has a fatal flaw.

It really does.

And the flaw is thermal expansion.

Right.

Because the volume of a liquid changes with temperature.

When you heat a solution, the increased kinetic energy causes the molecules to push slightly further apart.

The liquid expands.

The volume increases.

Okay.

So if I have a one liter flask of a one molar solution and I heat it up, the liquid expands, so it's now taking up, say, 1 .05 liters, I haven't added or removed any solute particles.

The mass is exactly the same.

But because my denominator, the volume just got bigger, my molarity calculation says the concentration went down.

Exactly.

The molarity changed simply because the room got warmer.

That's wild.

Yeah.

If you were doing precise thermodynamics in a field environment where temperatures fluctuate wildly,

you just cannot rely on volume based measurements.

You must use mass based units.

Because mass is an intrinsic property of the matter.

Right.

It does not change when the temperature rises.

So we switch to units like molality.

Notice the L instead of the R.

Molality measures the moles of solute per kilogram of solvent.

Yeah.

Or we use mole fraction, which is just the ratio of the moles of one component to the total moles of everything in the mixture.

And for environmental systems, we rely heavily on parts per million, or PPM.

Yes.

PPM is crucial.

We hear PPM all the time in the news, especially regarding atmospheric carbon.

We recently passed 400 PPM of TO2E2 in the atmosphere.

Conceptually, that just means if you grab a random sample of one million air molecules, 400 of them will be carbon dioxide.

Right.

And calculating PPM is an exercise in tracking pure mass.

Imagine you're trying to adjust the chemistry of a massive half million liter municipal swimming pool.

OK, big pool.

You dump in a standard 500 gram container of pool chemicals, say sodium bisulfate, to lower the pH.

A 500 gram bottle seems like nothing compared to half a million liters of water.

It is a tiny addition, but it matters chemically.

And to find the concentration of the sodium ions you just added, you don't worry about the changing volume of the pool under the hot summer sun.

Right.

You calculate the raw mass of the sodium ions in that bottle, and you divide it by the total mass of the half million liters of water.

Because a liter of water is basically a kilogram, so you are dividing a few grams of sodium by 500 ,000 kilograms of water.

Exactly.

You end up with a concentration of roughly one part per million.

You have successfully tracked the ratio of particles strictly by mass, completely eliminating the variable of temperature expansion.

And this obsession with tracking particles brings us to one of the strangest blind spots in nature.

Because once we start precisely counting particles, we discover that water behaves in a very peculiar way.

It really does.

For a specific set of physical properties, the solvent absolutely does not care what is dissolved in it.

It only cares how many particles are present.

Right.

These are known as colligative properties.

They depend entirely on the ratio of the number of solute particles to the number of solvent molecules, totally independent of the solute's chemical identity.

I always think of water acting like a bouncer at the door of a nightclub.

Oh, that's good.

Right.

The bouncer has a little metal clicker in his hand.

He is not checking IDs.

He doesn't care if the person walking in is a VIP, a celebrity, or just a regular guy.

He simply clicks the counter for every single body that walks through the door.

I love that.

And that bouncer's particle count determines some major physical changes.

For instance, Rowald's law shows us that adding solute particles lowers the vapor pressure of the solvent.

Okay.

Because the more solute particles you add, the fewer solver molecules are at the surface to escape into the gas phase.

Right.

And it also governs osmosis, right?

The movement of water across a cell membrane to try and dilute a more concentrated environment.

Yes, absolutely.

Yeah.

But the most visible colligative property we interact with has to be freezing point depression, which is why city trucks spread salt on the roads before a winter storm.

Exactly.

The solute particles physically get in the way of the water molecules trying to organize into a solid ice lattice.

They're blocking them.

Yeah.

The water molecules have to slow down even more, meaning the temperature has to drop even lower to finally overcome the interference of the solute and freeze solid.

Okay.

But let me ask you about our bouncer.

If I toss one molecule of ordinary table sugar into the water versus one molecule of table salt, does the bouncer click once or twice?

Ah.

Well, the bouncer clicks once for the sugar, but twice for the salt.

Wait, really?

Yeah.

This is quantified by the Van Haas factor.

Sugar is a covalent molecule.

When it dissolves, it stays completely intact.

One sugar molecule yields one dissolved particle.

So one body walks into the club.

One click.

Right.

But table salt, sodium chloride, is an electrolyte.

As we discussed with the lattice energy earlier, when salt dissolves, the water pulls it apart into two distinct ions, a sodium ion and a chloride ion.

Oh.

Two bodies walk into the club.

Two clicks.

Exactly.

Because the salt splits into two particles, one mole of salt has double the colligative effect of one mole of sugar.

The water counts twice as many obstructions.

That is exactly why salt is so economical for treating roads.

It punches above its weight class.

It lowers the freezing point of the ice twice as effectively as a non -electrolyte would.

It does.

In fact,

municipalities often use calcium chloride instead of sodium chloride for extreme coal.

Let me guess.

Because calcium chloride consists of one calcium ion and two chloride ions.

You got it.

It splits into three separate particles.

The bouncer clicks three times.

Three times the freezing point depression.

That is brilliant.

Okay.

So, we have explored true solutions where things dissolve perfectly into individual ions or molecules, but there is an entire world of mixtures that exist in an awkward,

like an in -between state.

Yes.

They refuse to fully dissolve, but they don't sink to the bottom either.

You're describing colloidal dispersions or colloids.

This is a state of matter where the particles are significantly larger than single molecules or ions.

Okay.

They're large enough to have a measurable surface area, but they're still small enough, usually between one and a thousand nanometers, that the constant chaotic, kinetic collisions of the water molecules keep them suspended indefinitely.

They never settle out by gravity.

We are surrounded by colloids.

Milk is a colloid of butterfat suspended in water.

Fog is a colloid of liquid water droplets suspended in the air.

Even jello is a colloid.

They're everywhere.

And the easiest way to identify a colloid is through the Tyndall effect.

The Tyndall effect.

Right.

Because colloidal particles are relatively large, they're on a similar scale to the wavelength of visible light.

When light hits them, the electromagnetic waves scatter.

So if you shine a laser pointer through a glass of pure salt water, the beam is totally invisible from the side.

The ions are just too small to interact with the light.

Exactly.

But if you shine that same laser through a dilute glass of milk or through gelatin, you see the solid line of the laser beam clearly illuminated in the liquid.

Oh, because the particles are scattering the light.

Yes.

It is the exact same physics as a car's headlights illuminating the path through a dense fog or sunlight streaming through the dust in a forest canopy.

That's so cool.

And colloids are generally broken into two categories based on how they interact with water, right?

Hydrophilic, meaning water loving,

and hydrophobic, meaning water fearing.

Oil and grease are famously hydrophobic.

They repel water.

Which begs a very practical question.

If the grease on my hands is terrified of water, how do I ever wash it off in the sink?

Well, you force a colloidal dispersion using soap.

Soaps and detergents are fascinating anthophilic molecules.

They have a split personality.

Oh, so?

They consist of a highly polar charged head that is hydrophilic.

It loves water, but attached to that head is a long non -polar hydrocarbon tail that is hydrophobic.

It hates water, but it loves lipids and oils.

So the soap molecule acts as a chemical bridge.

Exactly.

When you scrub a greasy pan with soapy water, the hydrophobic tails flee the water by embedding themselves deep into the grease.

The hydrophilic heads stay on the outside facing the water.

So the soap molecules completely surround the grease droplets, forming a microscopic spherical cage known as a micelle.

A micelle.

So the grease is locked inside and the entire outside of the cage is covered in water loving polar heads.

The grease droplet has essentially been disguised as a water soluble particle.

It becomes suspended as a colloid and the flowing water just washes the micelle right down the drain.

That is incredible.

We have literally traveled from the devastating depths of Lake Nyos to the grease traps of your kitchen sink, all guided by the fundamental behavior of solutions.

It's all connected.

We unpacked the tug of war of lattice energy and equation, the shifting balance of Le Chat

and the blind particle counting bouncer dictating colligative properties.

But you know, the beauty of these chemical laws is how they intersect to shape the future of our planet.

Consider this final thought for you to take away today.

Let's hear it.

We know the ocean is a massive sink for atmospheric carbon dioxide.

To counteract the resulting carbonic acid, the ocean is slowly dissolving vast deposits of solid calcium carbonate from the seafloor, things like shells and coral.

Right.

It's acting as a buffer to neutralize the acid.

Exactly.

But look at the stoichiometry of that buffering reaction.

One solid unit of calcium carbonate reacts to form three dissolved ions, one calcium ion and two bicarbonate ions.

Wow.

This incredibly slow planet wide process is drastically increasing the total number of dissolved particles in the sea.

Three particles entering the water for every one solid unit dissolved.

Right.

If we recall the colligative properties like osmotic pressure depends strictly on the total particle count, we face a staggering long term consequence.

Oh man.

This massive century long influx of ions is fundamentally altering the osmotic pressure of the entire ocean.

The bouncer's clicker is working overtime on a global scale.

Precisely.

How will this shift in the ocean's fundamental balance stress the delicate semi -permeable membranes of marine life that evolved in a less concentrated sea?

What happens when the ultimate solution is pushed past its limits?

That is heavy.

It's a profound demonstration that chemistry doesn't just happen in a beaker, it governs the entire world around us.

And we want to give a huge thank you to you, the listener, for taking this journey with us today.

Yes.

Thank you for studying with us.

Keep questioning the chemistry hidden in plain sight, officially signing off on behalf of the last minute lecture team.

Stay curious.

See you on the next deep dive.