Chapter 13: The Properties of Mixtures: Solutions, Colloids

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Hey you!

Ever wondered why sugar seems to just, um, vanish into your coffee?

Or why some medicines have that little instruction shake well before use.

Exactly.

Or even how our own cells manage that incredible chemical balancing act inside.

Mixtures are, well, they're everywhere.

Fundamental.

From the air we're breathing right now to the food we eat.

Yeah.

Okay, so let's unpack this.

Today we're embarking on a deep dive into the fascinating world of solutions and colloids.

And our guide for this is the really insightful work from Silberberg and Amatace's chemistry, the molecular nature of matter and change.

And what's truly captivating here, I think, is that we're not just talking about, you know, dissolving sugar.

Simple stuff.

Right.

We're actually exploring the fundamental forces.

How substances interact, why they mix, or sometimes why they absolutely don't.

And how these principles connect to everything.

Industry, biology.

Exactly, right down to the biochemistry within our own bodies.

It's all connected.

So our mission today is to cut through some of the, let's say, density of these concepts.

Yeah, make it accessible.

Think of this as your shortcut, maybe.

Getting those aha moments about mixtures without getting totally bogged down in the heavy details.

We'll guide you through it, step by step.

Major ideas, key principles, real world examples,

all just through conversation.

No visuals needed.

Pure chemistry chat.

Okay, so let's start at the very beginning.

What is a mixture, fundamentally?

Well, the key things our source highlights are variable composition.

You can have different amounts of components and they keep their original properties.

Like your cappuccino example,

or air, shampoo even.

All mixtures.

But we're going to zoom in on two crucial types today, solutions and colloids.

Right, solutions first.

Counts good.

A solution is what we call homogenous.

Meaning?

Meaning it exists as just one single phase.

Everything's perfectly blended right down at the molecular level.

You've got a solvent.

The main component doing the dissolving.

That's the one most abundant.

And then the solute, the substance being dissolved.

And usually the state of the solvent, liquid, gas, solid dictates the state of the whole solution.

Got it.

And the classic rule, the one everyone remembers from chemistry class.

Ah yes, like dissolves like.

It sounds so simple.

But the mechanics behind it, that's all about those intermolecular forces, isn't it?

Precisely.

That simple phrase hides a lot of complexity about attractions between molecules.

Are we talking ion dipole forces?

Hydrogen bonds?

Maybe weaker dispersion forces?

So give us an example, like salt in water.

Perfect example.

Table salt, NaCl is ionic.

Water is polar.

When salt dissolves, you get these strong ion dipole forces forming.

The water molecules kind of grab onto the ions.

Yeah, they cluster around them.

We call them hydration shells.

Water molecules surround the positive sodium ions, the negative chloride ions, essentially pulling them away from the salt crystal.

So those new water ion attractions have to be strong enough to break the salt crystal apart.

Exactly.

They have to compete with and overcome the forces holding the crystal lattice together.

I like the dance floor analogy sometimes.

If you're doing waltz, you look for other waltzers.

Trying to partner with a break dancer?

Probably not going to work well.

Similar dance styles, similar forces make dissolving easy.

That's a great way to put it.

And this also explains things like dual polarity.

Think about alcohols.

A small one, like methanol.

It dissolves really well in water.

Why?

Because its polar OH group forms strong hydrogen bonds with water molecules.

But if you take a bigger alcohol, like say, one hexanol, with a longer non -polar carbon chain...

Solubility drops way down.

Why is that?

Well, that long non -polar tail, it just doesn't interact well with water.

It can't effectively replace the strong hydrogen bonds between water molecules.

It disrupts water's network.

So it prefers non -polar environments.

Exactly.

It would rather mix with something like hexane using weaker dispersion forces.

It shows how subtle changes in molecular structure dictate where something will dissolve.

It's not just liquids mixing with liquids, right?

We have gas solutions.

Like the air we breathe.

A solution of, what, 18 different gases?

And solid solutions too.

Alloys.

Yeah, like brass zinc atoms mixed in with copper.

Or carbon steel, tiny carbon atoms fitting into the spaces between iron atoms.

Even waxes, like beeswax, can be thought of as solid, solid solutions.

What about gases dissolving in water, like oxygen for fish?

Generally, non -polar gases don't dissolve well in water.

Their solubility is low.

But that tiny amount of dissolved oxygen,

absolutely critical for aquatic life.

Even if it's only, say, three milliliters per 100 milliliters of water at room temp.

Still vital.

Definitely.

And some gases, like CO2, don't just dissolve physically.

There's a chemical reaction too.

Forming carbonic acid.

Precisely.

Which is key for ocean chemistry and, of course, fizzy drinks.

Okay, so we've covered the what mixes and how these forces work.

But the deeper question is, why does dissolving happen at all?

What drives it?

Ah, now we get into the energy changes, the thermodynamics.

The source breaks the solution process into this useful three -step cycle.

Right.

Involves energy costs and gains.

Exactly.

Step one, separating the solute particles, pulling them apart from each other.

That takes energy.

Endothermic.

Step two, separating the solvent particles.

Making space.

Also endothermic.

Takes energy.

Step three, the actual mixing.

Solute and solvent particles come together, attracting each other.

This usually releases energy.

Exothermic.

Right.

So the overall energy change, the heat of solution, or delta H -soul.

It's just the sum of those three steps.

And this balance is crucial.

It determines if the whole process gives off heat or absorbs it.

Absolutely.

When we talk about ionic solids dissolving in water, for instance, that energy released during mixing often relates to something called the heat of hydration, delta H -hydrator.

Which is the energy change when ions get surrounded by water molecules.

Yes, and that hydration process is almost always exothermic.

Why?

Because forming those strong ion dipole forces between ions and water releases a lot of energy, usually more than enough to compensate for breaking some of water's own hydrogen bonds.

And the strength of that hydration depends on the ion's charge density.

What's that again?

Good question.

Charge density is just the ratio of the ion's charge to its size, its volume.

So high charge, small size.

Means high charge density.

And that leads to a much stronger attraction to water molecules, a much more negative, more exothermic heat of hydration.

Can you give an example?

Sure.

Compare magnesium ion, Mg2 plus O, which is small and has a plus two charge, to a sodium ion, Na, plus O, which is larger and only has a plus one charge.

Mg2 plus will have a significantly more exothermic heat of hydration.

And we see this in action, right?

Like with the hot packs and cold packs.

Exactly.

Dissolve something like sodium hydroxide, NaOH in water.

The flask gets hot.

That's a big negative delta H zone.

Great for hot packs.

And the opposite.

Ammonium nitrate, NH4NO3.

Dissolve that and the flask gets noticeably cold.

It absorbs heat from the surroundings.

Large positive delta H zone.

Perfect for those instant cold packs you use for injuries.

But wait, you mentioned ammonium nitrate absorbs heat.

It's endothermic.

Why does it dissolve at all if it costs energy?

Energy isn't the whole story, is it?

Excellent point.

No, energy or enthalpy delta H isn't everything.

We absolutely have to consider entropy.

Entropy.

The measure of disorder.

Or randomness.

Sort of.

More accurately, it's about how energy is distributed, or the amount of freedom of motion particles have.

Gases have higher entropy than liquids.

Liquids higher than solids.

More ways to move.

More ways to arrange things.

And how does that relate to solutions?

Crucially, a solution almost always has higher entropy than the pure solute and solvent separated.

Why?

Because when you mix them, there are simply more possible arrangements.

More ways for the particles and energy to be distributed.

Nature tends towards states with higher entropy, greater dispersal.

Like sugar dissolving spontaneously, but never un -dissolving on its own.

Exactly.

The universe favors that increase in entropy.

So putting it all together, solution formation is this constant battle, this tug of war between enthalpy, the drive for lower energy, and entropy, the drive for more disorder or dispersal.

That's the core idea.

It explains so much.

Why doesn't salt dissolve in hexane, a non -polar solvent?

Because the energy caused to break the salt crystal is huge, and the mixing forces with hexane are weak.

Not enough energy released.

Right.

Enthalpy says no.

But why do octane and hexane, both non -polar, mix so easily, even though the enthalpy change is practically zero?

Because the entropy increases significantly when they mix.

Entropy wins.

Bingo.

And back to ammonium nitrate dissolving, even though it's endothermic.

The entropy increase must be huge.

The solid crystal breaking down, ions spreading out in water.

That massive entropy gain outweighs the energy cost.

Precisely.

It's a beautiful interplay between energy and randomness.

Okay, let's shift gears slightly.

We've talked about forces and simple solutions, but these same forces are absolutely fundamental in biology, aren't they?

They literally shape life.

They are the architects.

Think about proteins.

Huge, complex molecules made of amino acids.

Doing everything from being enzymes to antibodies.

And their function depends entirely on their specific intricate 3D shape.

How is that shape maintained?

By a whole hierarchy of these intermolecular forces.

So you have the strong covalent bonds holding the amino acid chain together.

That's the backbone.

Then hydrogen bonds create secondary structures like helices or sheets.

Then the protein folds up.

Polar and ionic side chains tend to be on the outside, interacting with water via ion dipole forces.

And the non -polar ones.

They usually cluster together on the inside, away from the water, interacting through weaker dispersion forces.

It's called the hydrophobic effect.

Like molecular origami held together by all these specific forces and sometimes even stronger links.

Right, like disulfide bridges.

Some amino acids, like cysteine, can form strong covalent bonds between them, locking parts of the protein chain together.

Amazing.

What about other biological examples of this dual polarity, like soaps?

Soaps and detergents are classic examples.

A typical soap molecule has that long, non -polar hydrocarbon tail.

Which likes grease and oil.

Exactly.

Interacts via dispersion forces.

And then it has a polar, often ionic, head group.

Which likes water.

Attracts water through ion dipole forces.

So the soap molecules surround tiny grease droplets, forming micelles, with the tails pointing in and the heads pointing out into the water.

Allowing the grease to be washed away.

Simple but brilliant chemistry.

It is.

And we see the same principle in cell membranes.

Ospilipids.

The main component.

They also have that dual nature.

Typically two non -polar fatty acid tails and a polar ionic head group.

And in water they automatically form.

The lipid bilayer.

Two layers thick.

Tails pointing inwards, away from water, heads pointing outwards, facing the watery environment inside and outside the cell.

Creating that essential barrier.

Absolutely vital for defining the cell.

And membrane proteins fit right into this structure too, with non -polar parts embedded in the bilayer and polar parts exposed to the water.

Even some antibiotics work this way.

Yes.

Like Gramicidin A.

It forms a channel through bacterial membranes, messing up their ion balance, which can kill them.

It exploits that membrane structure.

And finally the big one.

DNA.

The blueprint of life itself.

Its double helix structure relies heavily on these forces.

The backbone is polar, right?

Sugar phosphate.

Yes.

Polar and ionic.

Facing the watery environment.

Inside the helix, you have the nitrogen -containing bases.

Adenine, thymine, guanine, cytosine.

They stack on top of each other, interacting through dispersion forces.

And crucially, they form specific hydrogen bonds between pairs.

A with T, G with C.

Millions of these weak bonds holding the helix together.

Millions.

Collectively strong, but individually weak enough to be broken when the cell needs to replicate DNA or read its instructions.

It's an incredible design.

It really is.

Okay.

So forces dictate structure.

Let's go back to the macroscopic view.

Solubility limits.

What happens when you try to dissolve too much?

You eventually reach a point of equilibrium.

The rate of dissolving equals the rate of recrystallizing.

That point defines a saturated solution.

Holding the maximum amount of solute possible at that temperature.

Correct.

Less than that, it's unsaturated.

More than that, if you manage it carefully.

Super saturated.

Which is unstable, right?

Very unstable.

I love the demonstration with sodium acetate.

You have this clear, super saturated solution.

And you add one tiny seed crystal.

Woosh.

Rapid crystallization.

All the excess solute crashes out, leaving just a saturated solution behind.

It's visually quite dramatic.

Chemistry you can actually see happen.

Now, temperature.

We know it affects solubility.

How?

For most solids dissolving in liquids, like sugar in tea, solubility increases as temperature goes up.

Hotter water dissolves more sugar.

Makes sense.

But for gases?

It's the opposite.

Gases become less soluble in water as the temperature rises.

Which has real world consequences.

Absolutely.

Think thermal pollution.

Power plants discharge warm water into rivers.

That warmer water holds less dissolved oxygen.

Which is bad for fish and other aquatic life.

Especially since their metabolism might increase in warmer water, meaning they need more oxygen when less is available.

It's a double whammy.

Pressure matters too.

Especially for gases.

Henry's law.

Right.

It states that the solubility of a gas in a liquid is directly proportional to the partial pressure of that gas above the liquid.

More pressure, more gas dissolves.

Exactly.

That's why soda is bottled under high CO2 pressure.

When you open it.

Pressure drops.

CO2 comes bubbling out.

Yep.

And it's critical for divers.

As they descend, the increased pressure causes more nitrogen from the air they breathe to dissolve in their blood.

But if they come up too fast?

Pressure drops quickly.

And that dissolved nitrogen can form bubbles in their bloodstream in tissues.

That's decompression sickness.

The bends.

Very dangerous.

Understood.

Okay, let's quickly touch on concentration units.

We need ways to express how much solute is in a solution.

We have several common ways.

Molarity moles of solute per liter of solution.

That's probably the most common in general chemistry labs.

But there's also molality.

Yes.

Molality moles of solute per kilogram of solvent.

Notice the difference.

Solvent mass, not solution volume.

And things like mass percent, volume percent, mole fraction.

All useful in different contexts.

Molality is often preferred for very precise work, especially involving temperature changes.

Why is that?

Why is molality temperature independent, but molarity isn't?

Because molarity depends on the volume of the solution.

And liquid volumes change slightly with temperature.

They expand when heated, contract when cooled.

But mass doesn't change with temperature.

Exactly.

So molality, being based on mass of solvent, stays constant regardless of temperature.

Clever.

Okay, now for the big finale of solution properties.

Colligative properties.

You said these depend only on the number of solute particles.

That's the key thing.

Not what the solute is, but how many particles of it are dissolved.

There are four main ones.

First,

vapor pressure lowering.

Adding a solute lowers the solvent's vapor pressure.

Why?

It comes back to entropy, actually.

The solution already has higher entropy than the pure solvent.

So fewer solvent molecules need to escape into the gas phase to reach equilibrium.

Lower vapor pressure.

Raoult's law helps quantify this.

And lower vapor pressure leads directly to?

A higher boiling point.

That's boiling point elevation.

It takes more energy, a higher temperature, to get the vapor pressure up to atmospheric pressure so it can boil.

Yeah, conversely.

Lower freezing point.

Freezing point depression.

The solute particles disrupt the solvent's ability to form an orderly solid crystal structure, so you need to go to a lower temperature to freeze it.

So adding a sol basically stretches out the temperature range where the solvent stays liquid.

That's a great way to think about it.

And the fourth one is osmotic pressure.

This involves a semi -permeable membrane.

Yes.

Imagine separating pure solvent from a solution with a membrane that only lets solvent molecules pass through.

What happens?

Solvent molecules will naturally move from the pure solvent side into the solution side, trying to dilute it, driven by that tendency towards higher entropy or equal concentration.

Creating a pressure difference.

Exactly.

The pressure you need to apply to the solution side to stop that flow of solvent is called the osmotic pressure.

And it can be surprisingly large.

Useful for finding molar masses of big molecules, right?

Like proteins.

Incredibly useful, especially in biochemistry.

Measure the osmotic pressure, and you can calculate the molar concentration, and thus the molar mass if you know the mass you dissolved.

Now what about solutes like salt, NaCl, that break apart into ions?

Ah, good point.

Strong electrolytes.

NaCl doesn't just give one particle per formula unit, it gives two Na plus and Cl.

So it should have double the effect on colligative properties compared to something like sugar, which doesn't dissociate.

Ideally, yes.

We use the Van't Hoff factor I to account for this expected number of particles.

For NaCl is ideally 2.

For CaCl2 it would be ideally 3.

1CO2 plus SC2CO.

But you said ideally, does it always work out perfectly?

Not quite.

Especially in more concentrated solutions or with highly charged ions, ions don't always act completely independently.

Opposite charges attract, forming temporary clusters or ionic atmospheres.

This slightly reduces their effective number, so the observed Van't Hoff factor is often a bit lower than the ideal integer value.

Fascinating.

And the real -world applications of these colligative properties are everywhere.

Oh, definitely.

Anti -freeze in your car radiator, usually ethylene glycol.

It lowers the freezing point of the water in winter and raises the boiling point in summer.

Uses both depression and elevation.

And de -icing roads with salt.

Classic freezing point depression.

Salt dissolves in the thin layer of ice or moisture, lowering its freezing point below the ambient temperature causing it to melt.

CaCl2 is often used because it dissociates into three ions, giving a bigger effect per mole.

And biologically.

Osmosis is huge, isn't it?

Fundamental.

It controls the water balance in our cells.

IV drips have to be isotonic, have the same osmotic pressure as body fluids, usually using .159 NaCl solution.

What happens if they're not?

If the IV fluid is hypotonic, lower solute concentration, water flows into red blood cells, causing them to swell and potentially burst.

If it's hypotonic, higher concentration, water flows out of the cells, causing them to shrivel up.

Neither is good.

So osmosis regulates cell shape.

What else?

It's how plants draw water up from the soil through their roots.

It's involved in how our kidneys regulate body fluid volume by managing sodium iron concentration, even preserving food, like salting meat or making pickles in brine.

That uses osmosis to dehydrate microbes.

Exactly, draws water out of them, killing them or inhibiting their growth.

Okay, one last category to cover.

Colloids.

You call these the in -between mixtures.

Yeah, they're kind of halfway between a true solution and a coarse suspension where particles settle out.

Colloidal particles are larger than simple molecules, but small enough to stay dispersed.

What size range are we talking?

Roughly one to a thousand nanometers in diameter, big enough to scatter light, but small enough to avoid settling due to gravity.

Examples.

Lots.

Whipped cream is gas dispersed in liquid.

Mayonnaise is liquid in liquid.

Milk, fog, smoke, jello, even blood plasma contains colloidal components.

Many biological fluids are aqueous souls, solid particles dispersed in water.

And you said they scatter light.

Yes, that's a key characteristic called the Tyndall effect.

Because the particles are relatively large, they intercept and scatter light beams.

So that's why you can see a sunbeam through dusty air or headlight beams and fog.

Precisely.

True solutions don't do that because the solute particles are too small.

What else is unique about colloids?

You often see Brownian motion under a microscope.

The colloidal particles are constantly being bombarded by the smaller invisible molecules of the dispersing medium, causing them to jiggle around randomly.

Which was actually evidence for molecules existing.

It was crucial evidence back in the day.

And why don't they just clump together and settle out?

Often it's because the surfaces of the colloidal particles absorb ions from the disposing medium, giving them all the same type of charge, either all positive or all negative.

So they repel each other.

Exactly.

That electrostatic repulsion keeps them dispersed and stable.

But you can make them clump up, coagulate them.

Yes.

Heating can give particles enough energy to overcome the repulsion.

Or, more commonly, you can add an electrolyte with oppositely charged ions.

These neutralize the surface charges on the colloidal particles.

Allowing them to stick together.

Right.

They aggregate into larger clumps and eventually settle out.

This is exactly the principle used in water treatments.

Ah, so adding chemicals like aluminum sulfate.

Or iron -thoride.

These provide highly charged positive ions, AL3 +, F3 +, that are very effective at neutralizing the negative charges typically found on colloidal clay particles and microbes suspended in raw water.

They clump together, form larger flock, and then settle out or are filtered.

So water purification relies heavily on understanding colloids.

Absolutely.

It's a multi -step process, screening out large objects, coagulation and flocculation, as we just discussed, sedimentation, filtration, and finally disinfection with chlorine, ozone, or UV light to kill any remaining pathogens.

And sometimes water softening, too.

Yes.

Removing hard water ions like calcium, Ca2 +, and magnesium, Mg2 +, often by ion exchange, swapping them for sodium Na plus ions.

And there's one more purification method mentioned, reverse osmosis.

RO.

Yeah.

This uses that osmotic pressure principle, but in reverse.

Yeah.

You apply external pressure to the salty or impure water, a pressure greater than its natural osmotic pressure.

This forces water molecules through a semi -permeable membrane, leaving the dissolved ions and other impurities behind.

Powerful stuff.

Used for home filters.

And large -scale desalination plants, turning seawater into fresh drinking water.

It's energy -intensive, but very effective.

RO.

Wow.

Okay.

So from the tiniest molecular forces dictating how things dissolve, to the structure of DNA, to keeping ourselves alive, and even purifying the water we drink,

solutions in colloids are just fundamental.

They really are.

We've seen that like dissolves like is driven by intermolecular forces.

We've seen the crucial interplay of energy and entropy that decides if something dissolves.

And we've seen how colligative properties depend just on the number of particles, impacting everything from antifreeze to cell life.

CB.

What really stands out to you from this whole deep dive is there one connection that seems particularly elegant.

RO.

Hmm.

That's tough.

I think the way those relatively weak intermolecular forces, like hydrogen bonds and dispersion forces, orchestrate the precise folding and function of massive proteins in DNA, that's just incredibly elegant.

Life hinges on these subtle interactions.

CB.

Yeah, that's a good one.

So next time you stir sugar into your tea, or maybe you see those headlight beams cutting through fog.

RO.

Remember the chemistry.

CB.

Remember you're witnessing these profound,

sometimes hidden properties of mixtures in action.

Maybe take a moment to think what other solutions or colloids are shaping your world right now in ways you hadn't considered.

RO.

Makes you look at things differently.

CB.

We certainly hope this deep dive has given you a fresh perspective and some really valuable insights into solutions and colloids.

From all of us here at the deep dive and the last minute lecture team, thank you so much for listening.

Thanks everyone.

We'll explore another fascinating topic with you again soon.