Chapter 24: Complex Ions and Coordination Compounds

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I want you to start by picturing a gemstone.

Just close your eyes for a second unless you're driving, of course, and picture a piece of turquoise.

Ah, that's a great example.

Right.

You know that vivid, opaque blue -green color?

It's totally iconic.

It's been used in jewelry for thousands of years, from ancient Egypt to the American Southwest.

Now, just hold that image in your mind.

It really is a striking color.

And, you know, for a chemist, it poses a very specific question.

Exactly.

The big question is, why is it blue?

Because in our last deep dive, we talked about Chapter 23,

and we talked about color coming from oxidation reduction reactions.

The redox reactions, yes.

Yeah, you know, electrons physically moving from one atom to another, changing the charge, causing all those fireworks and explosions.

Right.

That is the really dramatic side of chemistry.

But this, what we are looking at today, is entirely different.

Because if you look at that turquoise, or a ruby, or an emerald, the metal inside isn't exploding.

It isn't changing its oxidation state at all.

It's just sitting there, completely chemically stable.

Yeah.

Yet it is beaming out this brilliant color.

Precisely.

The color isn't coming from the metal reacting and changing into something else.

It is coming from what is attached to the metal.

It's really about the company the metal keeps.

The entourage.

In a manner of speaking, yes.

The groups bound to the metal center.

And this brings us to the heart of what we are covering today.

We are taking a deep dive into Chapter 24 of your text, which covers complex ions and coordination compounds.

This is effectively a last -minute lecture for anyone staring down a chemistry exam right now.

Or just anyone who wants to know why the world looks the way it does.

Yeah.

We are going to unpack the mystery of these structures.

It's a fascinating feel called coordination chemistry.

It explains how a metal ion can bind to a specific number of neutral molecules, or anions, to form a complex.

And how that structure dictates literally everything from color of a gemstone to the way a cancer drug works.

So our mission today is to go from the history, starting with some 18th century chemical detectives, all the way to the high -tech theories that explain the electrons dancing around inside these things.

We're talking structure, naming,

isomers, and something called crystal field theory.

Which I know sounds intimidating, but it is actually the key to unlocking the color mystery.

And we will wrap up with some very real applications.

But first, let's rewind.

Let's go back to the beginning.

Before we knew about orbitals and electron spins, what did chemists actually think these things were?

Well, we have to travel back to the 18th century.

The first coordination compound was actually discovered completely by accident.

It was Prussian blue.

The pigment?

Yes.

A deep blue pigment found in the early 1700s.

Artists loved it, but chemists, they were absolutely baffled.

They had no idea what the structure was.

It was just this blue mystery powder.

Essentially.

The real scientific puzzle started a bit later, in 1798, with a chemist named Tasseret.

He was working with two very common, very stable things, cobalt chloride, which is CoCl3, and ammonia, NH3.

Okay, so a salt and a gas.

Both are totally happy on their own.

Right.

But Tasseret found that they could combine to form a new yellow crystalline compound.

The formula he derived was CoCl3 dot 6NH3.

That dot in the formula.

I remember seeing that in hydrates.

It usually implies things are just sort of packed together.

That was exactly the prevailing theory at the time.

They called them molecular compounds.

The idea was that two stable compounds just decided to sit next to each other in the crystal.

But the mystery got much deeper because chemists started finding more of them.

Using the same ingredients?

Same ingredients, different ratios, different colors.

In 1851, they found a purple compound.

It was cobalt chloride and ammonia again, but this time the ratio was CoCl3 dot 5NH3.

So six ammonias gives you yellow, five ammonias gives you purple.

And it continues.

They found a green one with four ammonias, and a violet one also with four ammonias.

Wait, two totally different colors for the exact same ratio ingredients?

That must have driven them crazy.

It really did.

It completely defied the logic of the time.

So they decided to do some detective work using the standard tool of 19th century chemistry, silver nitrate, agon O3.

Okay, let's pause here and remind the listener, especially the students, what silver nitrate does.

This is basic solubility rules, right?

Correct.

Silver ions, Ag +, have a very strong affinity for chloride ions, Cl -.

If they find each other in solution, they instantly snap together to form silver chloride, AgCl.

Which is that white solid precipitate.

Exactly.

It's a way of asking the molecule, how loosely are you holding your chlorides?

If the chloride is free -floating, the silver grabs it and makes powder.

Right.

So let's walk through their results.

They started with the yellow compound, the one with six ammonias, CoCl3 .6NH3.

They dissolved it, added the silver, and boom!

They got three moles of silver chloride precipitate.

Okay, let me do the math.

The original formula had three chlorides in it, and they got three moles of precipitate.

That means all three chlorides were loose.

They were totally free -floating in the water.

Spot on.

But then they tested the purple compound, the one with five ammonias.

It also has three chlorides in the written formula.

But when they added the silver… Let me guess.

It wasn't three.

They only got two moles of precipitate.

Whoa.

So one chloride is just missing.

Not missing.

Hiding.

It's in witness protection.

In a sense.

And with the green compound, the one with four ammonias, they only got one mole of precipitate.

Two chlorides were hiding.

This is where Alfred Werner enters the story, right?

1893.

The father of coordination chemistry.

And he was only 26 years old when he proposed this theory.

Werner looked at this data, the colors, the precipitates, and he proposed a truly revolutionary idea.

He said the missing chlorides weren't missing at all.

They were bound directly to the metal center.

So they were holding on too tight for the silver to grab them.

Exactly.

Werner proposed that metals have two types of valence.

First, there's the primary valence.

This corresponds to the oxidation state.

In all these cobalt compounds, cobalt is in the plus three oxidation state.

Okay.

That makes sense.

That balances the overall charge.

But then there is a secondary valence, which we now call the coordination number.

This represents the actual number of groups bound directly to the metal atom.

So this defines the inner circle.

A perfect analogy.

The inner circle is the complex.

It comprises the metal plus whatever is chemically bonded directly to it.

Werner suggested that for cobalt three plus, the coordination number is almost always six.

It demands six friends in the inner circle.

Okay.

Let's apply Werner's inner circle logic back to the mystery compounds.

Start with the yellow one.

Six ammonias.

Since cobalt wants six friends and there are six ammonias available, the ammonias take all the seats.

They form the complex ion, co -NH3 -6, with a three plus charge.

In the text, you'll see this written inside square brackets.

The square brackets indicate the complex.

What about the chlorides?

There's no room for them in the inner circle.

So the three chlorides are relegated to the outer sphere.

They just float around nearby in the solid to balance the charge.

Because they are outside the brackets, when you put it in water, they float away.

The silver can grab them, hence three moles of precipitate.

Man, that is elegant.

Now do the purple one.

It only has five ammonias.

Cobalt still absolutely demands six friends.

It takes the five ammonias, but it still have one empty seat.

So it grabs one of the chlorides and pulls it into the inner circle.

So the complex inside the brackets becomes CoClNH3 -5.

Correct.

And notice the charge changed.

Cobalt is plus three, the five ammonias are neutral, zero, and the chloride inside is minus one.

So the total charge of that complex is plus two.

And since one chloride is locked inside the VIP room, there are only two chlorides left floating outside to balance that plus two charge.

And those two outside chlorides are the only ones the silver can find, hence two moles of precipitate.

Werner, you genius.

That is such a clean explanation.

It perfectly matches the data.

It really was brilliant.

He solved the puzzle by realizing that the metal isn't just a charged ball.

It has a specific geometry and a specific number of attachment points.

So this gives us the vocabulary we need for the rest of the deep dive.

We have the complex, which is the metal plus attachments.

And those attachments have a special name.

Ligands, from the Latin ligar, meaning to bind.

A ligand is any group, whether it's a molecule or an ion that binds to the central metal.

And the number of points where they attach is the coordination number, or CN.

Correct.

And we should definitely mention the shapes that go with these numbers.

If you look at figure 24 to 2 in your text, you'll see this.

Werner predicted that coordination number six corresponds to an octahedral geometry.

Imagine the metal in the center and ligands at the top, bottom, front, back, left, and right.

Oh, like a jack, you know, from the game jacks.

Exactly like a jack.

That is the most common one you'll see.

But we also see coordination number four quite often.

That gives us either a tetrahedral shape or a square planar shape.

And what about two?

Is that a thing?

Two is rare.

You mostly see it with what we call the coinage metals.

Copper one plus, silver one plus, and gold one plus.

That forms a simple linear shape, just a straight line.

Okay, I want to pause here and do a last minute lecture service for the students listening.

There is a work example in the text, example 24 to 1.

It asks us to look at a formula and figure out the numbers.

I want to walk through the logic of this out loud.

That's a great idea.

Let's do it.

The complex in the book is, all right, open bracket, cool sail NO2, then NH3, 4, close bracket, and the whole thing has a plus one charge.

Okay, looking at this alphabet soup, the question asks for the coordination number and the oxidation state of the cobalt.

Let's start with the coordination number.

It is definitely the simpler of the two.

I just count the things inside the square brackets, right?

Yes, specifically you are counting the donor atoms.

So what do we have?

We have one chloride CL, one nitrite NO2, and four ammonias NH3.

So one plus one plus four equals six.

The coordination number is six.

Which implies that octahedral jack shape we talked about.

Now, the oxidation state, this requires a little forensic accounting.

Okay, so we know the total charge of the complex is plus one, because that's the superscript sitting outside the bracket at the end.

Right, so now I need to list the charges of the individual ligands, because you just have to know these from earlier chapters.

Ammonia.

Neutral, charge is zero.

Chloride, that's a halogen, so minus one.

Correct, and nitrite.

That's a polyatomic ion.

Yeah.

Also minus one.

Exactly.

So let's set up a simple algebra equation.

The charge of cobalt plus negative one for chloride, plus negative one for nitrite, plus zero for ammonia equals plus one total.

So cobalt minus two equals plus one.

Solve for cobalt.

Move the two over, cobalt equals plus three.

Perfect.

That is the exact process.

Break it down, sum the parts, and solve for the metal.

I feel smarter already.

But let's zoom in on these ligands for a second.

You said ligand means to bind, but how exactly are they binding?

Is it like a normal covalent bond where we share electrons?

It is a covalent bond, but with a slight twist.

In a standard covalent bond, like the bond between carbon and hydrogen, each atom usually brings one electron to the table to form the pair.

It's potluck dinner.

Everyone brings a dish.

Right.

But in coordination chemistry, the bond is a coordinate covalent bond.

In this specific case, the ligand brings both electrons.

So the ligand is paying for the whole date.

Indeed.

We view this through the lens of Lewis acid -base theory.

The ligand has a lone pair of electrons.

It is an electron pair donor, which makes it a Lewis base.

The metal ion has empty orbitals waiting to accept those electrons, so it acts as a Lewis acid.

Okay, so anything with a lone pair can technically be a ligand.

Theoretically, yes.

Water, H2O, has lone pairs on oxygen.

Ammonia, NH3, has a lone pair on nitrogen.

These are very common ligands.

And these simple ones are called monodentate, right?

Yes.

Mino means one.

Dentate comes from the Latin dens, meaning tooth.

A monodentate ligand has one tooth.

It bites the metal at one spot.

Water, ammonia, chloride, these are all one -bite ligands.

Which implies the existence of multi -toothed monsters.

Oh, absolutely.

We call them polydentate ligands, or much more commonly, chelating agents.

Chilli.

That sounds like chilla.

Which is Greek for a crab's claw.

Imagine a molecule that is long and flexible.

It has a lone pair on one end and another lone pair on the other end.

It can wrap around the metal and bite it in two places at once.

It grips the metal literally like a claw.

That sounds much more secure than just bumping into it with one tooth.

It is remarkably secure.

The classic example you'll see in Figure 24 -3 of the textbook is ethylenediamine.

Which is almost always abbreviated as Just Ender.

Its structure has two amine groups connected by a short carbon chain.

It is a bidentate ligand two teeth.

So it forms this stable little ring structure with the metal.

And are there ligands with even more teeth?

The absolute champion that you need to know is EDTA.

Ethylenediamine tree acetate.

It is hexadentate.

Hexa.

Six teeth.

Six points of attachment.

It wraps around the metal ion completely, encasing it in a chemical cage.

Once a metal is grabbed by EDTA, it is effectively removed from the solution.

It's totally sequestered.

You know, I feel like I've seen EDTA on ingredient labels.

You definitely have.

Go check your shampoo bottle later.

It almost always says Desodium EDTA.

Why is there crab claw chemistry in my shampoo?

To handle hard water, tap water contains calcium and magnesium ions.

Those ions react with the soap molecules to form that gross scum you get on the shower walls.

That's the worst.

Well, the EDTA in the shampoo acts as a scavenger.

It wraps its six teeth around the calcium and magnesium ions as soon as the water hits your hair, locking them up in that cage so they can't mess with the soap.

It keeps the metal ions busy so the soap can actually do its job of cleaning.

That is actually incredibly practical.

It's also used in medicine.

If someone has severe lead poisoning, doctors can administer a form of EDTA intravenously.

It chelates the lead in the bloodstream, wrapping it up securely so the kidneys can excrete it safely without it damaging tissues.

So chelation literally saves lives.

Okay, moving on.

We know what these complexes are made of.

Now, how do we name them?

Because I'm looking at Chapter 24, and I've seen some names that look like they are 40 letters long.

The nomenclature can definitely be daunting.

It looks like a completely foreign language.

But like any language, it has strict grammar rules.

If you follow the algorithm, you can name anything systematically.

All right.

Give us the crash course, the last -minute lecture guide to naming.

What is rule number one?

First, we deal with the ligands.

If the ligand is an anion, meaning it's negative, we change its ending to an O.

So chloride becomes chlorido, sulfate becomes sulfato.

And oxide becomes oxido.

Correct.

Even nitrate becomes nitrito.

Okay, got it.

Rule number two.

What about neutral ligands?

Usually, they just keep their regular names, but there are some critical exceptions you just have to memorize.

And these are the ones that always trip students on exams.

Water is not water.

It is aqua.

Ammonia is not ammonia.

It is Mn.

Quick spelling check on that.

Yes.

A -M -M -I -N -E.

Two Ms.

If you spell it with one M, it refers to the organic functional group.

In complexes, it's always double M.

And the third big exception, carbon monoxide is carbonyl.

Okay.

Aqua, ammion, carbonyl.

I'm locking those in.

Rule number three.

The numbers.

How do we say how many there are?

We use standard Greek prefixes,

mono, di, tri, tetra, penta, hexa.

So if you have five waters, it is pentaqua.

But wait, looking at the rules, I saw B centris in the text.

What's that about?

Ah, that is a special rule for when the ligand name itself is complicated or already contains a prefix.

Take ethylenediamine.

It already has diocese right in the middle of its name.

If we said diethylenediamine to mean two of them, it would sound like one giant molecule with two ethyl groups.

It's confusing.

So we need a totally different set of numbers to avoid the stutter.

Exactly.

We use best for two, trace for three, and tetrachas for four.

And we have put the ligand name in parentheses.

So if you have two ethylenediamines attached, it is bis parentheses ethylenediamine close parentheses.

That actually clarifies a lot.

It prevents ambiguity.

Okay.

Rule number four.

The order.

How do we string this all together?

Legands come first, and they are listed in alphabetical order.

Then the metal comes last.

And immediately after the metal name, you put the oxidation state in Roman numerals in parentheses.

Wait, alphabetical order?

Is that based on the prefix or the base name?

Based on the name itself.

So pentaco counts as A for aqua.

You ignore the penta when alphabetizing.

Okay.

And there is one final trap, right?

Rule five.

What if the whole complex is negative?

Yes, this is crucial.

If the complex ion is an anion, the metal name must end in eight.

So zinc becomes zincate.

Cobalt becomes cobaltate.

Platinum becomes platinate.

And some of them go old school Latin, don't they?

They do.

This is where it sounds really fancy.

Iron becomes ferrate.

Copper becomes cuprate.

Silver becomes argentate.

Lead becomes plumbate.

Okay.

Let's test this algorithm.

Let's reverse engineer a name from example 24 .2 in the book.

The name is deep breath.

Pentaquacloridochromium, Roman numeral three chloride.

That is a massive mouthful.

Let's break it down piece by piece.

Start from the left.

Pentacle.

Five water molecules.

Chlorido.

One chloride ion.

And because it has the O, it's inside the complex.

Chromium three.

The metal is chromium sitting in the middle with a plus three charge.

Perfect.

So the inner circle of the complex inside the brackets is CrClH2O5.

Now we need to check the overall charge of that complex to figure out the very end of the formula.

We know Cr is plus three, Cl is minus one, water is zero.

So plus three minus one equals plus two.

The complex is occasion with a plus two charge.

And the written name just ends with the word chloride.

That's the counter ion sitting outside the bracket.

Since our complex is plus two, how many outside chlorides do we need to balance it?

We need two chlorides because they are minus one each.

So the final complete formula is CrClH2O5Cl2.

See, when you break it down step by step like that, it's really just a logic puzzle.

Exactly.

It is highly systematic.

Once you know the rules, you can't be tricked.

All right.

Speaking of puzzles, let's talk about when the pieces are exactly the same, but the final picture is different.

Isomerism.

Isomers are compounds the exact same chemical formula, but different structures.

In coordination chemistry, the variety is just staggering.

We generally divide them into two major camps.

Structural isomers and stereoisomers.

Structural means the actual bonds are different, right?

Like different things are physically connected to the metal.

Yes.

For example, ionization isomerism.

Imagine you have a complex where a sulfate group is inside the inner circle, bonded directly to the metal, and a bromide ion is floating outside.

The formula would be COPO -SO4 -NH35 -Br.

Okay.

Now imagine they swap places.

Bromide goes inside to bond with the metal.

Sulfate gets kicked outside.

The formula is COPR -NH35 -SO4.

It's the exact same atoms in the beaker, totally different chemical.

One gives a precipitate if you add barium, which is the sulfate test, and the other gives a precipitate if you add silver, the bromide test.

They behave completely differently.

Then there's linkage isomers.

This one is really cool.

Look at figure 24 -4.

It happens with those ambidextrous ligands.

Correct.

Take the nitrite ion NO2-, it has a nitrogen atom and two oxygen atoms.

It can bond to the metal through the lone pair on the nitrogen.

We call that nitro or nitrito -N.

And interestingly, it creates a yellow complex.

Or it can literally flip around and bond through a lone pair on one of the oxygen atoms.

We call that nitrito or nitrito -O.

That creates a red complex.

It's basically like a USB stick.

You can plug it in two ways.

And unlike a USD stick, both ways actually work, but they give you a distinctly different result.

Now let's move to the 3D space.

Stereoisomers.

The connections are exactly the same.

The metal has the exact same neighbors,

but they are arranged differently in three -dimensional space.

This is where we get geometric isomerism.

The classic example is cis -trans isomerism.

You see this most often in square planar and octahedral shapes.

Look at figure 24 -5 for the square planar version.

I remember cis and trans from organic chemistry.

Cis means same side, trans means across.

It works exactly the same way here.

Imagine a square planar complex like PDCl2NH32.

Platinum is in the middle, and all four ligands are flat like they're sitting on a table.

You have two chlorides.

If the two chlorides are sitting at adjacent corners, right side by side, 90 degrees apart, that is cis.

And if they are across from each other, 180 degrees apart, opposite corners, that is trans.

Exactly.

And does this seemingly tiny change in angle actually matter?

Enormously.

We will discuss this more in the application section, but that specific difference, cis versus trans, in that platinum molecule is why one of them cures cancer, and the other one is totally useless.

Wow, okay.

Geometry literally saves lives.

Now, geometric isomers happen in those octahedral jacks too, Figure 24 -6.

They do.

In an octahedron where the coordination number is 6, you can have cis and trans if you have two ligands of one type.

But if you have three ligands of one type, and three of another, like an MA3B3 formula, you get a special geometric type called facial and meridional, usually abbreviated as fac and mer.

Fac and mer.

Sounds like two guys you'd meet at a hardware store.

Imagine the octahedron again.

If the three identical ligands occupy one triangular face, like the three corners of a pyramid on one side that is facial, they are all bunched together on one side of the molecule.

And meridional.

If they are arranged in an arc right around the middle, like a meridian or the equator on a globe that is meridional, they form a flat T shape through the center.

I can visualize that.

Now, we have to talk about the mind -bending one.

Optical isomerism, Figure 24 -8.

This is where we need the listener to just put down their pen for a second and look at their hands.

Your left hand and your right hand are mirror images of each other.

They look the same.

But if you try to put your left glove on your right hand, it doesn't fit.

They are non -superimposable.

We call that handedness or chirality.

Exactly.

Some coordination complexes are chiral.

A classic example is cobalt, with three of those ethylenediamine rings wrapped around it.

Because of the way those bidentate rings twist, they angle sort of like the blades of a desk fan or a propeller.

The entire molecule has a handedness.

So you can actually have a left -handed cobalt complex and a right -handed one.

Yes.

They are called enantiomers.

They are chemically identical in almost every conceivable way.

Same boiling point, same color, except for how they interact with plane polarized light.

One will rotate the light to the right, which is dextrorotatory, and the other rotates it to the left, leverotatory.

And Alfred Werner actually used this exact phenomenon to prove his octahedral theory, didn't he?

He did.

This was really his mic drop moment.

Critics in the late 1800s were arguing that the optical activity they were seeing was just coming from the carbon atoms in the ligands themselves.

Because organic chemists already knew carbon could be chiral.

They didn't believe the metal geometry itself was causing the twist.

So what did he do?

Werner went into the lab and synthesized a complex with absolutely zero carbon, only inorganic ligands.

It was a massive complex with cobalt, ammonia, and hydroxide bridges.

And guess what?

It twisted the light.

It was optically active.

He separated the left and right -handed versions.

It proved, undeniably, that the chirality was a property of the 3D octahedral geometry itself.

It was the final nail in the coffin for the old molecular compound theories.

Werner really was playing 4D chess while everyone else was playing checkers.

He truly was.

Okay, so we have the structure figured out.

We have the names.

We have the isomers.

But we still haven't answered the big question from the very intro.

Why are they colored?

Why is the turquoise blue?

Werner's theory elegantly explains the shape, but it doesn't explain the light show.

To answer that, we have to leave the simple Lewis acid -base model behind and enter the world of crystal field theory, or CFT.

This is section 24 to 5.

Crystal field theory.

Okay, this sounds advanced.

Break it down for us.

The old theories just couldn't explain color or magnetic properties.

CFT approaches the problem by looking purely at electrostatics.

It simplifies things greatly.

It just treats the ligands as negative point charges approaching the metal, like little negative magnets.

Okay, so we have the metal sitting in the middle with its positive nucleus and its outer D electrons.

And here come the ligands bringing their negative charge.

Right.

And remember, a fundamental rule.

Electrons repel electrons.

Now, I need you to visualize the 5D orbitals of the metal.

Look at figure 2411 if you have the book.

They are labeled DXCD, DXE, DZ2, and DX2 minus E2.

I'm visualizing five weird sort of balloon animal shapes.

Good.

In an isolated metal atom floating in space, all five of these orbitals have the exact same energy.

They are degenerate.

But now, imagine six ligands approaching to form an octahedron.

They come in straight along the X, Y, and Z axis.

Like an alien invasion fleet marching directly down the axis.

Exactly.

Now, look closely at those orbital shapes.

Two of them, the D squared and the DX squared minus E squared, have lobes that point directly along those axes.

They are sitting right in the direct path of the oncoming negative ligands.

So that's a head -on collision of electrons.

Yes.

It represents a region of extreme repulsion.

That direct repulsion drives the energy of those two specific orbitals way up.

But the other three orbitals, the DX, DS, DX, their lobes point between the axes.

The ligands slide right past them without that direct head -on conflict.

So they are less repelled.

Correct.

Their energy actually goes down relative to the average.

So the five orbitals that used to be totally equal are now split into two distinct groups.

A high energy upstairs group of two, which we call egg, and a lower energy downstairs group of three, which we call T2G.

Figure 2412 shows this beautifully.

And that gap between the low ones and the high ones, that is a very important name.

It is called the crystal field splitting energy.

We denote it with a delta symbol and a little o for octahedral delta o.

This energy gap is the key to literally everything we are about to discuss.

OK.

So we have a split.

Downstairs bunks and upstairs bunks.

This explains magnetism first, right?

This is the high spin versus low spin dilemma.

Exactly.

Imagine you are a D electron trying to find a place to live.

You want to be in the lowest energy state possible.

So you fill the bottom three bunks first, one in each.

But what if you are the fourth electron?

You suddenly have a choice.

Choice A.

I can room with another electron in one of the bottom bunks.

But electrons are negatively charged, so they naturally hate each other.

Forcing them together in one orbital costs energy.

We call that the pairing energy.

Choice B.

I can pay the energy cost to jump the gap entirely and take a top bunk all by myself.

And that cost is the splitting energy, delta o.

So it's just a simple thermodynamic cost benefit analysis.

If the splitting energy is small, meaning the gap is narrow, the electron looks at it and says, hey, it's cheaper to move upstairs than to deal with a roommate.

So it jumps up.

This means the electrons spread out as much as possible.

Maximum unpaired electrons.

We call this high spin.

This happens when you have weak field likens.

Things like iodide, bromide, chloride, fluoride.

They don't push very hard on the orbitals, so the split they cause is small.

But what if the gap is huge?

If you have strong field likens like cyanide, CN-, or carbon monoxide,

CO, they push aggressively.

They cause a massive split.

The gap is huge.

The fourth electron looks at that giant jump and says, no thanks, I'll deal with a roommate in the basement.

So it pairs up in the bottom level instead.

Exactly.

This forces pairing and minimizes the number of unpaired electrons.

We call this low spin.

And chemists have actually ranked these likens in a master list called the spectrochemical series, right?

Yes, you'll see it in the text.

It ranges from the weak field halogens on one end to medium field like water and ammonia in the middle up to the strong field cyanide and carbonyl on the far end.

So if I have a metal with chloride likens, it'll probably have a small gap.

High spin, lots of unpaired electrons, which means it's highly magnetic.

If I chemically swap those chlorides for cyanides, the gap widens the electrons pair up.

It becomes low spin and it might lose its magnetism entirely.

Precisely.

And we can actually physically weigh the magnetism in a lab to figure out the structure of a complete unknown.

This is the experiment shown in figure 2415, right?

The Goya balance.

Yes, the Goya balance.

You take a tube of your sample and weigh it with an electromagnet turned off.

Then you turn the magnet on.

If the substance is paramagnetic, meaning it has those unpaired electrons, it is physically attracted to the magnetic field.

The magnet literally pulls it down, making it appear to weigh more.

And if it's diamagnetic, meaning all the electrons are paired up.

It is actually slightly repelled by the magnetic field.

It appears to weigh less.

By measuring this tiny weight change, we can calculate exactly how many unpaired electrons are inside the molecule.

Let's apply this.

Look at example 2425.

It mentions the complex NiCN4 2 -.

Okay.

Nickel 2 plus is a D8 ion, meaning it has eight D electrons.

The problem states that lab tests show this complex is diamagnetic.

So all eight electrons must be paired up.

Now we have four ligands, so it could be tetrahedral or square planar.

If it were tetrahedral, the crystal field splitting is inherently very small.

Electrons would spread out.

It would be high spin, giving us unpaired electrons.

But the scale said it's diamagnetic, so it can't be tetrahedral.

Exactly.

It must be square planar.

Square planar geometry creates a unique, extremely large splitting gap that forces those eight electrons to fully pair up in the lower energy levels.

So the physical pull of a magnet confirmed the microscopic 3D shape.

That is just incredible.

It's beautiful deductive chemistry.

Right now, let's get to the color.

Yes.

Finally, bringing it home.

How does this energy gap make the gemstone blue?

Well, color is just light.

And white light contains all the colors of the rainbow, all different energies.

When white light hits a complex ion, the complex absorbs some of that light.

Specifically, an electron in the lower T2G level can absorb a photon, but only if that photon has the exact same energy as the delta O gap.

So the photon hits an electron in the downstairs bunk, gives it a burst of energy, and boosts it to the upstairs bunk.

We call that a daily transition.

The photon is consumed in the process.

Now think about the size of the gap.

If the gap is small because we have weak field ligands, it only takes a low energy photon to make the jump.

Low energy light is on the red side of the spectrum.

Okay, so if the complex absorbs the red light out of the white light, what's left over to hit my eye?

The complementary colors.

Green and blue.

So the compound looks green or blue to us.

Figure 2416 shows this complementary color wheel.

But what if the gap is large?

If you have strong field ligands, the gap is big.

It takes a high energy photon to boost the electron across that huge divide.

High energy light is violet or blue, so it absorbs the blue.

If you subtract blue from white light, what is left?

Yellow or orange.

So it looks yellow.

Let's look at example 24 to 6.

It compares chromium 3 plus with water ligands, which is violet, versus chromium 3 plus with ammonia ligands, which is yellow.

Why?

Walk us through it using the series.

Okay, let me try.

Ammonia is further to the right on the spectrochemical series than water.

It's a stronger light.

So ammonia pushes harder and creates a bigger gap.

A bigger gap means it requires higher energy light to make the jump.

Higher energy is blue or violet.

So the ammonia complex absorbs the violet light.

And if it absorbs violet, the color wheel says it transmits yellow.

Spot on.

While the water complex has a smaller gap, absorbs lower energy yellow -green light, and transmits violet.

The color you see is always the complement of the color absorbed.

It is entirely about the magnitude of that energy gap.

That is so cool.

It literally connects the invisible geometry of the molecule to the physics of light.

The gemstone is basically a microscopic machine that subtracts specific wavelengths of light based purely on how hard the ligands are pushing on the metal.

That is a very poetic and completely accurate way to put it.

Before we move to the real world applications, I want to touch on one more theoretical concept from section 24 to 8.

Equilibria.

We talk about how stable these complexes are, but they aren't totally permanent, right?

Ligands can swap in and out.

Yes.

Complex formation is always an equilibrium process in solution.

We measure it with formation constants KOF.

If you dissolve a bare metal ion in water, it instantly forms an aqua complex.

If you then add ammonia, the ammonia molecules replace the water molecules one by one in stepwise fashion.

And we mentioned earlier that the crab claws, the chelates like EDTA or NN, hold on much tighter than monodentate ligands.

Why is that?

The text calls it the chelate effect.

This is a beautiful example of thermodynamics in action.

Compare nickel binding to six individual ammonia molecules versus nickel binding to three ethylenediamins, those bidentate and rings.

In both cases, the metal is coordinating to six nitrogen donor atoms.

The actual chemical bond strengths are very similar, but the un -complex is 10 billion times more stable.

10 billion.

That's a massive difference for the same type of bonds.

Why?

Entropy.

Disorder.

Look at the balance reaction equation.

When three un -molecules bind to the nickel, they displace six water molecules.

So you put in three particles plus the metal, and you get out the complex plus six free water particles released into the solution.

So the total number of free floating particles increases.

It nearly doubles.

The system becomes much messier.

Entropy increases.

And the universe loves disorder.

That positive entropy change strongly drives the reaction forward.

It is much harder to un -claw a chelate simply because to reverse the reaction, you'd have to gather up all those free waters and fight that entropy term.

So the universe inherently prefers the chaos of the claw.

Exactly.

Now what about kinetics?

Sections 24 -9 and 24 -10.

Is thermodynamic stability the same thing as being, you know, stuck?

No, and this is a crucial distinction that trips up a lot of students.

Stability having a large KF is purely thermodynamic.

It just means the equilibrium lies far to the right.

Lobility, on the other hand, is kinetic.

It describes how fast the ligands actually pop on and off.

Give me an example of something that is stable but fast.

Take the copper ammonia complex.

It's very thermodynamically stable.

But if you put it in water, the ammonia ligands and water ligands are swapping places billions of times a second.

It is highly labile.

Contrast that with inert.

Inert means slow exchange.

Chromium 3 plus and cobalt 3 plus are famous for being kinetically inert.

If you dissolve chromium chloride in water, it starts out green.

It is thermodynamically unstable.

It really wants to turn violet by replacing the chlorides with water.

But it takes days or even weeks for that color change to happen at room temperature.

So it's thermodynamically unstable but kinetically inert.

It's stuck.

And this property was actually a massive stroke of luck for Alfred Werner.

He studied cobalt and chromium precisely because they stayed put long enough for him to isolate the isomers and study them.

If they had been labile, they would have scrambled into a mess before he could analyze anything.

Lucky for Werner, but he capitalized on it.

He did.

All right.

We've covered the heavy theory.

We've done the math.

Now for section 2411, let's talk about why this matters to the person listening.

Applications.

The most historically impactful application is undoubtedly cisplatin.

This is the famous cancer drug, right?

Yes.

And the story of its discovery is fascinating.

It was found in the 1960s by a biophysicist named Barnett Rosenberg.

He wasn't even trying to cure cancer.

He was studying the effect of electric fields on E.

coli bacteria.

To do this, he used platinum electrodes in the cellular broth.

What happened?

The bacteria stopped dividing.

They just grew into these massive long filaments.

He soon realized it wasn't the electricity stopping cell division.

It was a chemical compound forming from the platinum electrode, slowly reacting with ammonium and chloride in the solution.

He isolated it.

Cisdiamonide chloroplatinum 2.

Cisplatin.

So how does a square planar platinum complex work against cancer?

It enters the cell.

The chloride ligands are somewhat labile, so they fall off inside the cell.

The platinum atom then binds directly to the DNA.

Specifically, it coordinates to the nitrogen atoms on adjacent guanine bases.

It forms a cross -link.

It essentially handcuffs the DNA to itself.

Yes.

It puts a huge 45 -degree kink in the DNA double helix.

When the cancer cell tries to replicate its DNA to divide, the replication machinery hits that kink and gets completely stuck.

The cell realizes its DNA is irreparably damaged and triggers apoptosis, programmed cell suicide.

It has saved countless lives, especially in treating testicular and ovarian cancers.

But here's the kicker, and it ties perfectly back to our isomer section transplatin.

The geometric isomer.

The exact same molecule, but with the chlorides across from each other.

Transplatin is completely inactive as a cancer drug.

Why?

It's the same formula.

Same metal.

Same ligands.

The geometry is wrong.

The physical spacing between the leaving chlorides and the trans isomer doesn't match the spacing of the DNA bases.

It can't form that critical cross -link kink.

Also, kinetically, transplatin is more reactive and gets deactivated by other proteins in the blood before it even reaches the DNA.

Hey, it is mind -blowing.

The difference between a blockbuster, life -saving cure and a completely useless, potentially toxic molecule is just a 90 -degree angle in the bonds.

It really highlights the vital importance of stereochemistry.

What about photography?

I know digital is totally taken over, but old -school black -and -white film relied heavily on this stuff.

The fixing step in developing film is pure coordination chemistry.

When you take a picture, you have unreacted silver bromide, AgBr, left on the film.

You need to wash it away to stop the film from further reacting to light and turning completely black.

But AgBr is highly insoluble.

It won't just wash off with water.

So use a chemical fixer.

Sodium thiosulfate.

The thiosulfate ion, S2O32-, acts as a very strong ligand.

It binds to the silver to form the complex AgS2O32 with a 3 -minus charge.

This complex is highly soluble in water.

So the solid silver bromide literally dissolves into the liquid and washes away, leaving your permanent image behind.

And finally, biology.

Yeah.

We are basically powered by coordination complexes.

Absolutely.

We mentioned porphyrins earlier in the text.

These are massive, naturally -occurring tetradentate ligands.

We call them macrocycles.

They form a giant flat ring with four nitrogens pointing inward.

And if you drop a metal in the middle.

If you put a magnesium ion in the middle of a porphyrin ring, you get chlorophyll.

The molecule that turns sunlight into sugar for plants.

And notice its color.

It absorbs red light, reflecting green.

Now, if you take that exact same porphyrin ring and swap the magnesium for an iron 2 -plus ion, you get heme, the core active site of hemoglobin.

Which carries oxygen from our lungs to every cell in our blood.

And again, notice the color change.

Magnesium complex.

Green.

Iron complex.

Red.

It's the specific metal and the ligands tuning that crystal field energy gap, literally sustaining life as we know it.

And there is a dark side to this too, right?

Carbon monoxide poisoning.

This goes right back to the spectrochemical series.

Remember we said CO is a strong field ligand?

The strongest on the list.

Because of that, it binds to the iron in your hemoglobin much, much more tightly than oxygen does.

If you inhale CO gas, it knocks the oxygen off your hemoglobin and grabs the iron in a death grip.

And it refuses to let go.

It effectively shuts down your entire oxygen transport system.

Exactly.

It simply outcompetes the oxygen for that single coordination site on the metal.

So, quite literally, understanding the principles of the spectrochemical series is a matter of life and breath.

Wow.

A sobering but incredibly powerful thought to end on.

It's all about who grabs the metal the tightest.

Indeed it is.

Well, we have covered a massive amount of ground today.

We've gone from Werner's deductive reasoning and the mystery of the hiding chlorides.

To the crab claws of chelation and the entropy effect.

To the invisible splitting of deorbitals and crystal field theory that paints our entire world in color.

And finally, to the delicate, precise geometry that makes cisplatin a lifesaver.

My final thought for the listener is this.

Next time you see a vivid color, whether it's a firework, a piece of turquoise jewelry, or just the green leaves outside your window, remember you aren't just seeing color.

You are seeing a quantum measurement.

You are seeing the exact energy gap between two microscopic orbitals perfectly tuned by the entourage of atoms dancing around a metal center.

The world is just a giant spectroscopy experiment.

And we are the lucky observers.

Thank you so much for guiding us through the complex world of, well, complexes.

It was my absolute pleasure.

And to you, the learner, thanks for sticking with us for this last -minute lecture deep dive.

Good luck with the chemistry exam.

Or just enjoy the colors of the world a little bit more today.

This is the last -minute lecture team signing off.