Chapter 13: Properties of Solutions

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Have you ever stopped to think about how sugar just disappears into your coffee, but then like oil and water absolutely refuse to mix?

Or that fizz when you pop up in a soft drink.

Exactly.

It seems simple, but it's actually deep chemistry, isn't it?

It really is.

And today, we're doing a deep dive into exactly that.

Properties of solutions.

We're drawing heavily from a cornerstone text, chemistry,

the central science.

Right.

And our goal really is unpack the chemistry behind how things mix or don't mix.

Yeah.

What drives those interactions, how we measure them, how we use them.

We want to show you the chemistry that's kind of hiding right in front of us.

And hopefully by the end of this, you'll have a really solid grasp of solutions.

You'll be able to connect these ideas to things you see every day, whether you're prepping for something specific or just curious about how the world works.

OK, so where should we start?

What exactly is a solution?

Good question.

The basic definition is it's a homogenous mixture.

Homogenous meaning completely uniform.

Exactly.

Every little bit you take looks and is exactly the same as any other bit.

Now, I think most people hear solution and immediately picture a liquid like salt water or something.

That's a super common thought, yeah.

But solutions aren't just liquids.

Not at all.

They can be solids too.

Sterling silver, for instance, that's copper dissolved uniformly in silver.

A solid solution.

And gases.

The air we're breathing right now, it's a solution of nitrogen, oxygen, argon, a few other things, all mixed perfectly.

Huh.

OK, but for today, we're mostly sticking with liquids.

We are, yeah.

Especially aqueous solutions, where water is the main player, the solvent.

OK, so let's define those terms.

Solvent and solute.

Simple enough.

The solvent is the component that's there in the largest amount.

It's the medium, basically.

And the solute?

That's everything else that's dissolved in the solvent.

Right.

So we know what they are.

But why do they form?

What makes two things actually, you know, mixed together?

Is it just random?

Not random, no.

Well, there are two main driving forces.

The first one is, well, it's a natural tendency for things to mix, to spread out.

OK.

We call this entropy.

It's a fundamental concept in science.

Entropy.

It sounds significant.

Can you break that down a bit?

Think of it like this.

The universe tends towards more disorder, more randomness.

If you have two different gases separated by a wall and you remove the wall, they just mix up on their own.

They do, spontaneously.

They spread out into the larger volume, becoming more randomly distributed.

That increases the system's entropy.

For gases, especially, where the forces between molecules are weak, this drive towards higher entropy is really key to why they mix.

So like the universe preferring a messy room over a tidy one.

Uh -huh.

That's one way to think about it.

A bit simplified.

But yeah, the tendency to spread out and mix is fundamental.

So that's entropy.

What's the other force?

The other big piece is intermolecular interactions.

Ah.

The forces between the molecules.

Things like dispersion forces,

dipole interactions, hydrogen bonds, ion -dipole forces, all those little attractions.

And the key idea here is that solutions tend to form when the attraction between the solvent molecule and the solute molecule is strong.

Strong compared to what?

Compared to the forces holding the solute molecules together and the forces holding the solvent molecules together, if the new solute -solvent hug is as good or better than the old solute -solute -solvent hugs.

Then they'll mix.

It's like an energetic trade -off.

Precisely.

It's an energetic tug -of -war.

You have to put energy in to separate the solute particles and to separate the solvent particles.

That's endothermic.

It costs energy.

Right.

But then when the solute and solvent particles come together and interact, energy is released.

That's exothermic.

The overall result depends on the balance.

Exactly.

The overall enthalpy of solution, we call it ASON, is the sum of those energy changes.

Sometimes it releases heat overall, like magnesium sulfate dissolving in water.

Epsom salts.

Using those instant heat packs, right?

That's the one.

That's an exothermic process.

But sometimes it absorbs heat, like ammonium nitrate in instant cold packs.

Ah, okay.

So dissolving can make things hot or cold.

It really shows that energy balance in action.

It does.

And it explains why some things just don't dissolve well.

If it takes way too much energy to pull the solute and solvent apart and you don't get much energy back when they mix.

And it's just not energetically worth it.

Pretty much.

The salt show is too endothermic, too unfavorable.

That's a big reason why, say, ionic salts don't dissolve in nonpolar solvents like oil.

The interactions are just too weak.

And we should be clear, we're talking about a physical process here, right?

Like salt dissolving in water.

You can get the salt back if you evaporate the water.

Absolutely crucial distinction.

This is dissolution.

It's not a chemical reaction where you form totally new substances, like reacting nickel metal with acid to make nickel chloride.

We're focused on just the physical mixing.

Okay, that makes sense.

Okay.

So we know how they form.

What about how much can dissolve?

Let's talk about saturation and solubility.

Right.

So imagine you start adding a solid solute to a solvent.

The solid starts dissolving, particles moving into the liquid.

Okay.

But at the same time, some of those dissolved particles bump back into the solid and stick there again.

That's called crystallization.

So dissolving and crystallizing are happening at the same time.

Exactly.

They're opposing processes.

And eventually they reach a point where the rate of dissolving equals the rate of crystallizing.

A balance.

A dynamic equilibrium.

Things are still happening.

Molecules are moving back and forth, but there's no net change in the amount dissolved.

And that's a saturated solution.

That's it.

A solution that's in equilibrium with undissolved solute.

It's holding the absolute maximum amount of solute it can at that specific temperature.

Like when they harvest sea salt.

Yeah.

They let the water evaporate?

Until the brine becomes saturated and then the salt starts crystallizing out because it can't hold anymore.

The maximum amount that can dissolve is called its solubility, often expressed in grams per hundred milliliter of water or something similar.

So if you've dissolved less than that maximum, the solution is?

Unsaturated.

It can still hold more.

But then there's that weird one.

Yeah.

Supersaturated.

What's going on there?

Ugh.

Supersaturation.

That's when you manage to dissolve more solute than you normally could at that temperature.

It's tricky to do.

Often involves dissolving a lot at a high temperature and then cooling it very carefully.

And it's unstable.

Very unstable.

If you disturb it, maybe by adding just one tiny crystal of the solute?

A seed crystal.

Rouge.

Yeah.

All the excess solute suddenly crashes out of solution, crystallizing really rapidly.

It's quite dramatic to see.

It's like the solution was holding its breath and the seed crystal lets it exhale.

Wow.

Okay, so what actually controls solubility?

What factors make something dissolve well or poorly?

The biggest rule of thumb is probably dissolves like...

Heard that one before.

What does it mean chemically?

It means substances with similar types of intermolecular forces tend to dissolve in each other.

Okay, give me an example.

Well, polyliquids, like say acetone, mix really well with other polar liquids, especially water.

They both have dipole forces.

And water has hydrogen bonding, which acetone can participate in to some extent.

The interactions are favorable.

Makes sense.

So non -polar things.

Like gasoline or cooking oil, which are mostly non -polar molecules held together by weaker dispersion forces, don't mix well with water.

Because water molecules are much happier hydrogen bonding with each other than interacting weakly with the oil molecules.

You got it.

The water -strong attractions kind of exclude the non -polar molecules.

It's not energetically favorable to break up those strong water interactions.

And this even applies within families of molecules, right?

Like alcohols.

Definitely.

Take methanol, one carbon, very small, plus that polar OH group, mixes perfectly with water.

But then look at hexanol, six carbons in a chain, much larger non -polar part.

Much less soluble in water.

Way less soluble.

The long non -polar tail starts to dominate its behavior.

This like dissolves, like has cool real -world consequences like vitamins.

Absolutely.

Think about vitamin C and the B vitamins.

They're loaded with polar groups, OH groups, things that can hydrogen bond with water.

So they're water -soluble, which means...

Your body doesn't really store them long -term, excess just gets flushed out, you need a regular intake.

Okay, and the other vitamins, ADEK.

Those are fat -soluble.

They have long non -polar hydrocarbon structures, they prefer fatty tissues and non -polar environments in the body.

So your body can store those, makes a big difference for nutrition.

Huge difference.

Okay, so that's salt, lute, solvent interactions.

What else affects solubility?

Pressure is a big one, but mainly for gases.

Right.

Pressure doesn't really change how much salt dissolves in water, does it?

Not significantly, no.

But for gases, big effect.

The higher the pressure of a gas above a liquid, the more of that gas will dissolve in the liquid.

There's a law for that, isn't there?

There is.

Henry's law.

It basically says gas solubility, SG, is directly proportional to the partial pressure of the gas above the solution.

SG, aeol's KPG, where K is a constant.

And the everyday example is?

Carbonated drinks, they bottle sodas under high CO2 pressure.

So lots of CO2 dissolves.

Then you open it.

The pressure drops suddenly to atmospheric pressure, the CO2 solubility plummets, and fizz bubbles escape.

Simple but elegant.

And this has more serious applications too.

Like diving.

Oh, definitely.

Deep sea diving.

Under high pressure, deep underwater, more nitrogen gas from the air mixture dissolves in the diver's blood.

If the diver comes up too fast, the pressure decreases rapidly, and that dissolved nitrogen can form bubbles inside their blood vessels and tissues.

Sounds bad.

It is.

That's the bends.

Or decompression sickness.

Really dangerous.

Is that why they sometimes use helium mixes?

Exactly.

Helium is much less soluble in blood and tissues than nitrogen, even under pressure.

So switching to a helium -oxygen mix reduces the risk of the bends significantly.

It's applied chemistry saving lives.

Wow.

Okay, pressure for gases.

What about temperature?

Temperature effects are interesting because they differ for solids and gases.

For most solid solutes dissolving in water, solubility increases as temperature goes up.

Think about dissolving sugar in iced tea versus hot tea.

Much easier in hot tea.

Right.

More dissolves faster.

There are a few exceptions, but that's the general trend.

But for gases?

It's the opposite.

The solubility of gases in water decreases as temperature increases.

Ah, that's why warm soda goes flat faster.

Precisely.

The gas molecules have more energy.

They escape the liquid phase more easily.

Same reason you see tiny bubbles forming if you heat -tap water -dissolved airs coming out of solution.

Or why ice cubes made from boiled water are clearer, you've removed most of the dissolved gases beforehand.

That makes sense.

Heating helps gases escape, but helps solids break apart and dissolve.

Generally speaking, yes.

For ionic solids, the extra heat energy helps overcome the lattice energy holding the ions together and helps the hydrating process.

Okay, we've covered how solutions form, what affects how much dissolves.

Now how do we actually talk about how much is dissolved?

We need numbers.

We do.

We use qualitative terms like dilute or concentrated, but chemistry needs precision, so we use quantitative concentration units.

What are the main ones?

Well, there's mass percentage.

Pretty straightforward.

A 10 % salt solution by mass means 10 grams of salt in every 100 grams of the total solution.

Easy enough.

For really, really dilute solutions like environmental pollutants, we use parts per million ppm or even parts per billion ppb.

Like arsenic in drinking water.

The limits are tiny.

Exactly.

The WHO guideline is around 0 .01 bpm, or 10 ppb, that's like 10 drops of arsenic in an entire swimming pool.

Ppm and ppb are crucial for those trace amounts.

Okay.

What about units involving moles?

Right, there's mole fraction x, which is just the moles of one component divided by the total moles of all components in the solution, useful in certain physical chemistry contexts.

And the really common one in the lab.

Molarity, capital M.

That's moles of solute per liter of solution.

Super convenient for stoichiometry and solution prep.

But it has a downside.

It's temperature dependent.

Because the volume of the collusion changes slightly with temperature, the molarity can change too.

Ah.

So if temperature stability is important.

Then we often use molality, lowercase m, that's moles of solute per kilogram of solvent.

Since mass doesn't change with temperature, molality is temperature independent.

Very useful for studying properties over a range of temperatures.

Got it.

Molarity for volume -based work, molality when temperature changes matter.

Okay, this leads us into something really fascinating.

Colligative properties.

Ah, yes.

These are properties of solutions that depend only on the number of solute particles, not on what those particles are.

Wait, doesn't matter if it's sugar or salt or something else.

Yeah.

Just how many particles?

Isn't that wild?

It depends only on the concentration of solute particles relative to solvent particles.

Not their identity, not their size, not their charge, mostly.

Just the quantity.

That seems counterintuitive.

What's the classic example?

Antifreeze in your car radiator.

You add ethylene glycol to the water.

Why?

To stop it freezing in winter and boiling over in summer.

Exactly.

Adding the solute, ethylene glycol, lowers the freezing point of the solvent water and raises its boiling point.

And the amount it changes depends on how much antifreeze you add to its concentration.

It's a colligative property in action.

Okay, so what are the main colligative properties?

There are four, right?

Four key ones.

First, vapor pressure lowering.

Meaning the solution doesn't evaporate as easily as the pure solvent.

That's the idea.

Adding a non -volatile solute one that doesn't evaporate easily itself lowers the vapor pressure of the solvent.

The solute particles essentially get in the way, making it harder for solvent molecules to escape into the JAS phase.

Is there a way to quantify that?

Yes.

Routh's law describes it for ideal solutions.

It says the vapor pressure of the solvent above the solution is equal to the mole fraction of the solvent, X -olvent, times the vapor pressure of the pure solvent, P -degree solvent.

Solution is X -olvent, P -degree solvent.

And this is useful.

Oh yeah.

Think about distillation.

You separate liquids based on differences in their boiling points, which are related to their vapor pressures.

This principle is how we refine crude oil into gasoline, kerosene, etc.

The components with higher vapor pressures vaporize more easily.

Okay, so vapor pressure lowering is number one.

What's next?

Boiling point elevation.

This follows directly from vapor pressure lowering.

Well, boiling happens when the vapor pressure equals the external atmospheric pressure.

If the solute has lowered the vapor pressure, you need to heat the solution to a higher temperature to get its vapor pressure up to atmospheric pressure.

So the boiling point goes up.

Makes sense.

Yeah.

Is there a formula?

There is.

Bae bb is i kbm.

Bae tb is the change in doling point.

Kb is a constant specific to the solvent, and m is the molality of the solution.

Wait, what's that i?

i is the van hoff factor.

This is crucial.

It represents the number of particles the solute breaks into when it dissolves.

For something like sugar, glucose, which doesn't break apart in water, i, OMS -1.

But for sodium chloride, NaCl, it dissolves into two ions, Na plus and Cl.

So ideally i2 ,2.

For calcium chloride, Kqs2, it breaks into Ca2 plus and 2 Cl ions, so ideally i3.

So the effect depends on the total number of dissolved particles, ions included.

Exactly.

It's the total concentration of all solute particles that matters.

Okay, boiling point goes up.

What about freezing point?

Third property.

Freezing point depression.

Adding a solute lowers the freezing point of the solvent.

Like salting icy roads.

Perfect example.

Salt dissolves in the thin layer of water on the ice, lowers its freezing point, and causes the ice to melt, even if the air temperature is slightly below 0 degrees C.

32 degrees zero.

Same reason antifreeze works in winter.

And the formula looks similar.

Very similar.

HaF, iKfm.

Note the negative sign indicating a lowering.

Kf is the freezing point depression constant for the solvent.

Again, i and molality m are key.

Makes sense.

Lower freezing point, higher boiling point, what's the fourth one?

The fourth one is osmotic pressure, symbol i.

This one involves semi -permeable membranes.

Like cell membranes in biology.

Exactly.

These membranes are selective, they let small solvent molecules like water pass through, but block larger solute molecules, or ions.

Okay, so what happens?

Imagine you have pure water on one side of such a membrane and a salt solution on the other.

Water molecules can move back and forth, but the salt ions can't cross to the pure water side.

There's a net tendency for water molecules to move from the pure water side, where water concentration is high, into the solution side, where water concentration is effectively lower because of the solute.

This net movement of solvent across a semi -permeable membrane is called osmosis.

It's like the water is trying to dilute the solution.

That's a great way to think about it.

It's trying to equalize the solute concentration on both sides.

Now osmotic pressure is the external pressure you'd have to apply to the solution side to just stop this net flow of water, to counteract osmosis.

And this is important biologically.

Hugely important.

Think about red blood cells.

Their membranes are semi -permeable.

If you put them in pure water, a hypotonic solution, water rushes in trying to dilute the cell contents, and the cells swell and can burst.

That's hemolysis.

Ouch.

And if you put them in really salty water?

A hypertonic solution.

Water rushes out of the cells into the saltier surroundings, and the cells shrivel up.

That's crenation.

So IV fluids.

They don't have to be isotonic with blood.

Same effective solute concentration, so there's no net movement of water to damage the cells.

It's absolutely critical.

And this explains other things too, like pickles.

Yup.

Cucumbers in salty brine lose water via osmosis and become pickles.

Eating super salty food can cause water retention, edema, because your tissues hold on to water trying to dilute the excess salt.

Salting meat or canning fruit preserves them, partly by drawing water out of bacteria via osmosis, dehydrating and killing them.

Wow.

So versatile.

Is there a formula for osmotic pressure too?

There is, and it looks surprisingly like the ideal gas law at MMRG's IMRT, where I is the Van Haaf factor again, M is the molarity of the solution, R is the ideal gas constant, and T is the absolute temperature.

Interesting.

And you mentioned this is used to find molar mass.

Yes.

Especially for large molecules like proteins or polymers.

Osmotic pressure changes are relatively large even for dilute solutions, making them easier to measure accurately than, say, tiny freezing point changes.

You measure A, you know, T and R, you can calculate M, molarity, and if you know the mass concentration, grams per liter, you can find the molar mass, grams termoli, very powerful tool.

One quick thing about that Van Haaf factor I.

You said ideally I2 for NaCl, does it always work out perfectly?

Not quite perfectly in reality, especially in more concentrated solutions.

Sometimes the measured I value is a bit less than the ideal integer value.

Why is that?

It's due to something called ion pairing.

In solution, a positive ion, like Na +, and a negative ion, like Cl, might briefly stick together due to electrostatic attraction.

They act like a single unit for a moment, effectively reducing the total number of independent particles.

Ah, so the effective concentration of particles is slightly lower than you'd expect.

Exactly.

This effect is stronger for more concentrated solutions.

Ions are closer together.

And for ions with higher charges, stronger attraction.

But for dilute solutions of simple salts, like NaCl, the ideal value is usually a pretty good approximation.

Okay, that clarifies things.

We've covered true solutions thoroughly.

But you mentioned something else earlier, a sort of in -between category.

Right.

We need to talk about colloids or colloidal dispersions.

These are fascinating mixtures.

How are they different from true solutions?

In a true solution, the solute particles are individual molecules or ions.

Super tiny.

In a colloid, the dispersed particles are much larger, maybe large molecules like proteins or aggregates of smaller molecules.

They're typically between, say, 5 and 1 ,000 nanometers in diameter.

So bigger than molecules, but small enough not to settle out.

Exactly.

They're too small to settle out due to gravity like particles in a suspension, like muddy water, but large enough to be distinct from the solvent.

They stay dispersed indefinitely.

And they're common.

Incredibly common.

Milk is a colloid of fat globules dispersed in water.

Fog is a colloid of tiny water droplets in air.

Smoke, solid particles in air.

Whipped cream, air bubbles in cream.

Paint, cytoplasm in cells.

Colloids are everywhere.

Even jello.

Okay.

How can we tell if something is a colloid versus a true solution?

They both look uniform.

They often do, but there's a telltale sign.

The Tyndall effect.

The what now?

Tyndall effect.

Because colloidal particles are large enough to scatter light, if you shine a beam of light through a colloid, you can actually see the beam's path.

Ah.

Like headlights in fog.

Or sunlight streaming through dusty air.

Perfect examples.

True solutions don't do that.

The particles are too small to scatter light effectively.

The Tyndall effect is also why the sky is blue air molecules, and tiny dust particles scatter sunlight, with blue light scattered more effectively.

And why sunsets are red, the light travels through more atmosphere, scattering away the blue, leaving the red.

It's all colloid and light -scattering physics.

That's really cool.

So how do these larger particles stay suspended, especially in water?

Good question.

Colloids in water can be broadly categorized as hydrophilic, water -loving, or hydrophobic water -fearing.

Hydrophilic.

Like proteins.

Yeah.

Many large biological molecules, like proteins and enzymes, are hydrophilic.

They fold up in a way that puts lots of polar or charged groups on their surface.

These groups interact strongly with water molecules through hydrogen bonding or ion -dipole forces, keeping the molecule happily suspended.

Okay, but what about hydrophobic things, like oil droplets in water?

They don't want to be near water.

How do they form a stable colloid, like in milk?

They need help.

Hydrophobic colloids need to be stabilized.

One way is for the particles to absorb ions from the surrounding solution onto their surface.

So they all get coated with, say, negative ions.

Right.

Then all the particles have the same charge and repel each other, preventing them from clumping together and settling out.

Clever.

Is there another way?

Yes.

Using a mellocifying agent.

These are molecules that have both a hydrophobic part, a nonpolar tail, and a hydrophilic part, a polar or charged head.

Think of soap or detergent molecules.

How they work.

The hydrophobic tails stick into the oil droplet, while the hydrophilic heads stick out into the water.

They form a layer around the droplet, effectively giving it a water -friendly surface and allowing it to stay dispersed.

This is how milk stays stable, proteins act as emulsifiers, and how soap helps wash away grease.

That connects back to something you mentioned earlier.

Sickle cell anemia.

How does that relate to colloids and solubility?

It's a tragic, powerful example.

In sickle cell anemia, there's a tiny change in the hemoglobin protein molecule.

Just one amino acid is swapped.

A polar one gets replaced by a nonpolar one.

And that's small change.

Makes the hemoglobin molecule significantly less hydrophilic, less soluble in the aqueous environment of the red blood cell, especially when oxygen levels are low.

So they stick together.

They aggregate, forming long, rigid fibers inside the red blood cell.

This distorts the cell into that characteristic sickle shape.

Then those sickled cells cause all the problems.

Exactly.

They block small blood vessels, causing pain, organ damage.

It's a devastating illustration of how fundamentally molecular properties like solubility, driven by intermolecular forces, impact health.

Just like bile salts acting as emulsifiers are essential for digesting fats, it's all connected.

Incredible.

One last thing about colloids that keeps those particles bouncing around.

You said they don't settle.

Ah, that's Brownian motion.

Mamed after.

Robert Brown, a botanist who first observed it with pollen greens in water.

It's the random, erratic, jiggling motion of colloidal particles.

What causes it?

It's not the particles themselves moving purposefully.

It's because they are constantly being bombarded from all sides by the much smaller, much faster moving solvent molecules, like water.

The collisions are uneven at any given instant, pushing the larger particle around randomly.

So the solvent molecules keep them agitated and suspended.

Precisely.

This constant bombardment effectively counteracts gravity and keeps the colloidal particles dispersed throughout the medium.

Wow.

What a tour.

We've gone from why sugar dissolves, through entropy and energy changes, to why antifreeze works, how IV drips have to be just right, why the sky is blue, and even the molecular basis of a disease like sickle cell anemia.

It really shows how these concepts, solution formation, solubility, concentration, colligative properties, colloids, are woven into everything around us and inside us.

Understanding these principles from texts like chemistry, the central science just opens up a new way of seeing the world.

It really does.

It turns everyday observations into understandable chemical processes.

It feels empowering, actually.

Absolutely.

So maybe a final thought for our listeners to ponder.

Considering this delicate balance of forces that govern solubility and these colligative effects,

what other really critical biological processes might depend fundamentally on these very same chemical principles?

That's a great question to chew on.

Thank you for joining us on this deep dive into the properties of solutions.