Chapter 9: The Periodic Table and Some Atomic Properties

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Welcome back to another deep dive.

Today we're kind of shifting gears.

We are entering what I like to call crisis mode, but in the best possible way.

We are doing a last minute lecture.

Crisis mode.

Yeah.

I mean, that makes it sound like an emergency room.

But honestly, for a lot of students, the night before a general chemistry exam feels exactly like a medical emergency.

It really does.

We know who you are out there.

You're sitting there maybe with a double shot of espresso, staring at a textbook that feels like it weighs, you know, 50 pounds, and you have an exam on the periodic table and atomic properties in less than 12 hours.

Right.

Or, hey, maybe you weren't a student at all.

Maybe you're just someone who looks at the world and thinks, OK, I see the surface.

But what is the actual engine driving all of this stuff?

The engine is exactly the right metaphor.

Right.

Because today we are diving exclusively into chapter nine of general chemistry,

principles and modern applications, the 11th edition.

The chapter is titled The Periodic Table and Some Atomic Properties.

And look, I know what you're thinking.

You're thinking, I know the periodic table.

It's the wallpaper in every science classroom I ever sat in.

Ideally, yes.

But it's usually just decoration, right?

It's a colorful chart with boxes and letters.

But our mission today is to take that decoration and turn it into a tool.

We want to show you that the periodic table isn't just like a list of ingredients for the universe.

It is the ultimate map.

It is honestly a cheat sheet.

A cheat sheet is exactly what we need right now.

Yeah.

And if you understand the logic of this map, specifically the quantum mechanical logic we touched on in previous deep dives, you stop memorizing.

You just stop.

Thank goodness.

Exactly.

You don't have to memorize the fluorine is small and the francium is big.

You don't have to memorize ionization energies.

You can predict them.

You can basically derive the behavior of matter just by looking at where an element lives on that grid.

Prediction over memorization.

That is the mantra for this hour.

So here is the roadmap for today.

We are going to walk through this text section by section totally chronologically.

We're going to start with the history because there's some serious drama between Mendeleev and Meyer.

Then we'll look at how we classify elements.

Then we're going to get into the heavy lifting.

We're talking atomic radii, ionization energy,

electron affinity, magnetic properties, and

polarizability.

And the big idea I want you to keep in the back of your mind the whole time is this.

The electron configurations we learned in Chapter 8, the 1S2, 2S2 stuff.

Right.

The code.

That is not just abstract code.

That is the DNA that causes every single physical trend we are going to discuss today.

Chapter 9 is really just Chapter 8 made manifest in the real world.

I love that.

So if you are ready to turn that anxiety into confidence or just expand your understanding of the building blocks of reality, let's get into it.

Section 9 -1, classifying the elements.

Let's set the scene.

It's the mid -19th century, 1869 to be precise.

The wild west of chemistry.

It really was.

At this point, scientists had discovered about 63 elements.

They knew about gold, silver, oxygen, iron.

But there was no system.

Just a pile of facts.

Exactly.

Imagine trying to learn a language where you have a dictionary but absolutely no grammar rules.

That was chemistry in the 1860s.

And usually when we talk about fixing this mess, we talk about one guy, Dmitri Mendeleev.

He gets the stamp.

He gets the statue.

He does.

But the text makes a very specific point to highlight that he wasn't alone.

There's a German chemist named Lothar Meyer working the exact same time.

And Meyer actually cracked the code of periodicity visually in a way that I think is easier to understand than Mendeleev's initial tables.

This is the graph in Figure 9 -1.

It's arguably one of the most important graphs in the history of science.

It really is.

Meyer was looking for a pattern.

He proposed the Periodic Law, which essentially says if you line the elements up by weight, their properties will repeat in a rhythm.

To prove it, he plotted an atomic number on the x -axis.

Well, effectively the order of elements.

Right, the order.

And on the y -axis, he plotted molar volume.

Now we need to define that because it sounds a bit technical.

Yeah, let's break that down.

He took the atomic mass of an element, how heavy it is, and divided it by its density.

Because mass divided by density gives you volume.

Exactly.

It tells you how much space a mole of these atoms takes up.

It's a measure of, let's call it puffiness.

How big is the atom in its solid state?

And when he plotted this, he didn't get a straight line.

He didn't get a random scatter.

He got a heartbeat.

He got a wave.

The line goes up to a sharp peak, crashes down, goes along a valley, and shoots up to a sharp peak again.

And here's the kicker.

He looked at which elements were at the peaks.

The absolute highest points on the graph, the puffiest atoms in the universe were lithium, sodium, potassium, rubidium, and cesium.

The alkali metals, group one.

Exactly.

And then he looked at the valleys.

The smallest, tightest atoms were always in the same families, too.

He proved that physical properties cycle.

You add protons, you get different properties, but eventually you loop back around and hit the exact same node again, just an octave higher.

So Meyer finds the rhythm.

He builds the graph.

Why do we celebrate Mendeleev?

Why isn't it the Meyer table hanging in the classrooms?

Because of what they did with the information.

Meyer was a scientist describing the data he had.

Mendeleev was a visionary predicting data he didn't have.

Mendeleev was bold.

He was essentially gambling his entire professional reputation on this system.

The text mentions two specific things Mendeleev did that were just, I mean, audacious.

The first was correcting the experimentalists.

He was arranging his table by atomic mass, but sometimes an element wouldn't fit the pattern of the column it landed in.

Indium and uranium are the examples given in the chapter.

Instead of saying, oh, my table is broken, Mendeleev said, no, your measurements are wrong.

The data is wrong because it doesn't fit my beautiful theory.

That takes some serious confidence.

And he was right.

They remeasured them and the values were actually closer to what Mendeleev needed them to be.

But the second thing, and this is the mic drop moment of 19th century science, was the gaps.

The blank spaces.

Mendeleev arranged his cards and he saw holes.

He saw a gap below silicon and a gap below aluminum.

And he said,

there are elements here.

We just haven't found them yet.

And he named them like Eka silicon.

Right.

Eka meaning beyond or next.

So next silicon.

But he didn't just name it.

In 1871, he profiled it.

He said, this element will have an atomic mass of 72.

It will have a density of 5 .5 grams per cubic centimeter.

It'll be a dirty gray metal.

Its oxide will have a density of 4 .7.

He's giving you the police sketch of a suspect that no one has ever actually seen.

Precisely.

And then 15 years later, in 1886, a German chemist named Clemens Winkler discovers a new metal.

He calls it germanium.

Okay.

Let's look at the scoreboard from the text here.

Mendeleev predicted mass 72.

Germanium is 72 .6.

Incredibly close.

Mendeleev predicted density 5 .5.

Germanium is 5 .47.

Again, extremely close.

He predicted the oxide density would be 4 .7.

It was 4 .703.

That is precision.

That is the moment chemistry stopped being a hobby of just, you know, collecting rocks and weird smells and became a predictive theoretical science.

He proved that there was a fundamental order to the universe.

However, even with all that genius, there was a glitch.

We kept mentioning that they organized this by atomic mass.

Wait.

Right.

And 99 % of the time, that works.

Lighter things generally have fewer protons.

But there were exceptions.

The text points out the argon and potassium problem.

Argon is a noble gas.

Totally unreactive.

Potassium is an alkali metal.

It explodes in water.

Chemically, we know argon must come at the end of a period and potassium must start the next one.

But look at the weights.

Argon has a mass of 39 .9.

Potassium is 39 .1.

Argon is actually heavier.

So if you follow the periodic law, as it was written in 1869, you know, order by mass, you have to put potassium first.

That puts a highly reactive metal in the noble gas column and a noble gas in the metal column.

It breaks the whole game.

Mendeleev assumed the masses were wrong again.

He told them to remeasure argon.

But this time, the experimentalists were right.

Argon really is heavier.

The problem wasn't the data.

The problem was the rule itself.

It took until 1913 to fix this.

Enter Henry Mosley.

A young physicist working in Ernest Rutherford's lab.

He wasn't looking at weight.

He was looking at X -rays.

He set up an experiment where he bombarded metal targets with high -energy electrons.

Like a shooting gallery for subatomic particles.

Exactly.

And when you hit a target with electrons, it emits X -rays.

This is figure nine to two in the text.

Mosley measured the frequency, the pitch, basically, of these X -rays, and he found a beautiful mathematical relationship.

The square root of the frequency was proportional to the nuclear charge.

The number of protons.

The atomic number, yes.

He realized that the soul of an element isn't its weight, which can fluctuate because of isotopes and extra neutrons, but its proton count.

That is the ID number.

When you resort the table by atomic number, argon is number 18, potassium is number 19.

The order flips.

The glitch disappears.

So that gives us the modern periodic law.

Similar properties recur periodically when elements are arranged according to increasing atomic number.

And that leads us to the map we use today.

We have groups, which are the vertical columns, and periods, which are the horizontal rows.

And crucially for us, we have blocks.

The blocks are essential because they link directly to the electron configurations we learned in chapter eight.

Yes.

You have to visualize the table as neighborhoods, groups one and two, that's the S block.

Their valence electrons are filling sorbitals.

Groups 13 through 18 on the far right, that's the P block.

The wide valley in the middle, groups three through 12, that's the D block, or the transition metals.

And the two strips floating at the bottom, the F block, the lanthanides and actinides.

Knowing those blocks is going to save you so much time on the exam.

If you know where an element is physically located, you know its valence configuration instantly.

Exactly.

It's a massive shortcut.

Okay.

Let's move to section nine to two.

Metals and non -metals and their ions.

We can broadly slice the table into three categories.

Metals, non -metals and metalloids.

Metals are the vast majority.

Everything on the left and the center.

They are shiny, malleable, good conductors.

But chemically, the defining feature is that they tend to lose electrons.

Non -metals are clustered on the top right.

They are brittle, if solid, poor conductors.

And chemically, they tend to gain electrons.

And then you have the metalloids, the stair -step elements like silicon and germanium.

They are hybrids.

They look like metals, but behave a bit like non -metals, depending on the situation.

But the real meat of this section, and a huge part of any general chemistry exam, is predicting ions.

Why does sodium become plus one?

Why does oxygen become minus two?

It's all about the role models.

The noble gases.

Group 18.

Noble gases have a full valence shell.

Usually that means Ns2, Np6, that's eight electrons.

The octet.

In the quantum world, this configuration is nirvana.

It is perfectly stable.

So every other element on the table is just trying to achieve that state of grace.

Correct.

Look at a main group metal, like sodium,

Na, atomic number 11.

Its configuration is neon in brackets, then three as one.

It has one electron more than neon.

It is basically holding the bag.

To look like neon, it just has to drop that one electron.

So it loses it.

Na goes to Na plus, plus an electron.

And now Na plus has the exact electron configuration of neon.

It is isoelectronic with neon.

On the other side of the table, look at chlorine, Cl, group 17.

It is neon core, 3s2, 3p5.

It has seven valence electrons.

It is one single electron short of the argon configuration.

So it's desperate.

It is deeply desperate.

It will rip an electron off anything nearby.

It gains one to become Cl minus.

Now it is isoelectronic with argon.

So the rule for main group elements is delightfully simple.

Metals lose electrons to go backward to the previous noble gas.

Non -metals gain electrons to go forward to the next noble gas.

Simple enough.

But, and here comes the trap.

If you're listening to this while driving, pull over.

If you're taking notes, grab a red pen.

We really need to talk about transition metal ions.

This is where students lose points every single semester.

The logic we just use loss to achieve noble gas status doesn't quite work the same way for the D block.

And specifically, the order in which they lose electrons is completely counterintuitive.

Let's walk through an example from the text.

Let's take iron.

Fe atomic number 26.

First, expert, give me the electron configuration for the neutral atom.

Okay, following the Aufbau principle, the building up principle from chapter 8, we fill the 4s before the 3d.

So neutral iron is argon core, 3d6, 4s2.

Okay, so we put the 4s in, then we put the 3d in.

Now, I want to make an iron 2 ion, Fe2 +, I need to remove two electrons.

My gut instinct, and the instinct of basically every student who hasn't studied this specific chapter is to take the electrons from the place we just put them, the 3d orbital.

And your gut is totally wrong.

That is the trap.

The rule is, when forming an ion, transition metals lose the valence s electrons first.

Even though we wrote them earlier in the notation,

physically, those 4s electrons are in the fourth shell.

They are further out from the nucleus.

They are the first to be stripped away.

So for Fe2 +, we don't even touch the d electrons.

No, you script the 4s2 completely.

So Fe2 +, becomes argon core, 3d6.

The 4s is just gone, empty.

And if I want iron 3, Fe3+.

You take the 2s electrons first, and then you take one from the d orbital, so it becomes argon core, 3d5.

Let's do another one really quick, just to cement it, titanium, atomic number 22.

Neutral titanium is argon core, 3d2, 4s2.

If it loses two electrons to become titanium 2 +, we lose the 4s, it becomes argon core, 3d2.

Why is this such a big deal?

I mean, why can't I just say titanium 2 +, is argon core, 4s2?

That's still 20 electrons.

Because that configuration would have different magnetic properties, different bonding capabilities, and a totally different physical size.

It would describe an atom that literally doesn't exist.

If you write the wrong configuration, you get the magnetism question wrong, you get the bonding question wrong, you fail the whole page of the exam.

Always, always remember, transition metals, electrons out first.

Got it.

Moving to section 9 -3, sizes of atoms and ions.

This is probably the most visual part of the chapter.

We are asking a really simple question.

How big is an atom?

Which ironically is a very hard question to answer.

Atoms aren't billiard balls with hard shells, they are quantum clouds.

The electron density just trails off into infinity.

There's no hard boundary that says, welcome to carbon, city limits.

So how do we actually measure it in the lab?

We measure how close they get to each other, figure 9 -3 shows the methods.

We have the covalent radius, which is half the distance between two identical nuclei in a bonded molecule like Cl2.

We have the metallic radius, which is half the distance between nuclei in a solid metal crystal.

We use these distances as our proxies for atomic size.

And when we plot these radii on a graph, this is figure 9 -4, we see two major trends.

Trend number one.

What happens as we go down a group, say lithium to sodium to potassium?

The atoms get bigger, this one is very intuitive.

Lithium's valence electron is in the second shell, n equals 2.

Sodium's in the third, n equals 3.

Potassium is in the fourth, n equals 4.

Each step down adds a whole new layer to the onion.

The principal quantum number, n, determines the size of the orbital.

As n goes up, the probability cloud expands outward.

Easy.

Down equals bigger.

Now, trend number two, the paradox.

What happens as we go across a period, left to right, let's say lithium to neon?

The atoms get smaller.

Now, hold on.

We are adding protons, we are adding neutrons, we are adding electrons.

Lithium has three electrons, neon has 10.

Logic suggests that if I pack more stuff into a bag, the bag gets bigger, or at least stays the same.

Why does the atom shrink?

This is the single most important concept in Chapter 9.

If you understand this, you understand the rest of the deep dive.

It is called Effective Nuclear Charge, or ZF.

Let's really unpack ZF.

Imagine the atom is a tug of war.

In the center you have the nucleus, it's causative, it is pulling the electrons in.

On the outside edge you have the valence electron, it is negative, it is being pulled, but, and this is the crucial part in between them, you have other electrons.

The core electrons, electrons are negative, so they repel each other, the core electrons act as a shield.

They block the positive charge of the nucleus from fully reaching the valence electron.

So if I am a valence electron, I don't feel the full power of the nucleus, I feel a screened or weakened version.

Exactly, you feel the Effective Nuclear Charge.

Mathematically, the text simplifies this as ZF is approximately equal to Z minus S, the number of protons minus the shielding electrons.

So let's apply this math to the trend across the period.

Let's start at lithium.

Atomic number Z is 3.

It has two core electrons, the 1S2.

So the valence electron, the 2S1, feels roughly 3 minus 2, which equals plus 1.

Okay, now move one spot to the right, to beryllium.

Atomic number Z is 4.

It still has only two core electrons, the 1S2.

So the valence electron feels 4 minus 2, which is plus 2.

The magnet just got stronger.

Go all the way to the right side, to fluorine.

Atomic number is 9, still only two core electrons.

The valence electron feels 9 minus 2, which is plus 7.

That is a massive difference.

The valence electron in fluorine is staring at a plus 7 charge, while lithium's valence electron only sees a plus 1.

The nucleus in fluorine pulls that outer shell in tight, it crushes it down.

That is why fluorine is tiny compared to lithium.

But wait, I know what someone is thinking, we are adding valence electrons as we go across too.

Don't they shield each other?

Why doesn't S go up?

That is the key nuance.

Core electrons are great shields.

They are physically located between the nucleus and the edge.

Valence electrons are terrible shields.

They are in the exact same shell, roughly the same distance from the nucleus.

They don't block the view of the nucleus for each other effectively.

So as you go right, you add protons, more pull, but you don't add any significant shielding.

The net result is a much stronger grip.

The book also gives equation 9 .5.

Rr is approximately proportional to N squared divided by Zf.

Radius is a fight between the shell number N and the effective charge, Zf.

Moving down a group, N squared gets huge so the atom grows even though Zf creaks up a little.

Moving across a period, N stays exactly the same, but Zf shoots up so the radius shrinks.

So to summarize the atoms, down equals bigger because of shells, equals smaller because of Zf.

Correct.

Now, what happens when we make ions?

Does an atom change size when it loses an electron?

Drastically.

Think about a cation of a positive ion.

Take sodium.

It loses its one valence electron.

Two things happen.

One, you lost the entire outer shell.

The N equals three shell is just gone.

The atom shrinks to the N equals two boundary.

Two, the proton to electron ratio goes up.

You have 11 protons holding onto only 10 electrons.

They hold them tighter.

Cations are always, always smaller than their parent atoms.

And anions, negative ions.

The exact opposite.

You add an electron, but you didn't add any protons to balance it out.

So now you have more negative charge crowded into the same exact space.

Electron repulsion increases.

The cloud puffs out to relieve the stress.

Anions are always larger than their parent atoms.

This leads perfectly to the classic exam problem, the isoelectronic series.

This is where a professor gives you a list of ions that all have the exact same number of electrons and asks you to rank them by size.

Example from the text.

O2 minus, F minus, NA plus, Mg2 plus.

Okay, let's break it down slowly.

Oxygen starts with eight protons.

It gains two electrons, so it has 10 electrons.

Fluorine has nine protons.

It gains one electron, so 10 electrons.

Sodium has 11 protons.

Loses one, 10 electrons.

Magnesium has 12 protons.

Loses two, 10 electrons.

They all have 10 electrons.

They all have the exact electron configuration of neon.

So the shielding and the number of shells are perfectly identical across the board.

So the only difference is the engine, the protons and the nucleus.

Exactly.

Oxygen has only eight protons pulling on 10 electrons.

That's a weak pull.

The cloud drifts out.

It has a big radius.

Magnesium, on the other hand, has 12 protons pulling on those same 10 electrons.

Very strong pull.

The cloud is sucked in tight.

Small radius.

So the rule for an isoelectronic series is more protons equals smaller radius.

Correct.

Magnesium 2 plus is the smallest.

Oxygen 2 minus is the biggest.

Okay.

We've mapped them.

We've measured them.

Now let's torture them.

Section nine to four.

Ionization energy.

Expert.

Define it for us.

Ionization energy, designated EI, is the energy required to remove an electron from a gaseous atom.

Key words there are remove and gas phase.

You are fighting the nucleus.

You are ripping a negative charge away from a positive center.

That requires physical work.

This process is always endothermic.

You have to put energy in.

The values are always positive.

And the trend for ionization energy connects directly back to size, doesn't it?

It's basically the inverse of size.

Think about it.

If an atom is small, like fluorine, the valence electrons are very close to the nucleus.

They feel that high ZF.

They are held incredibly tightly.

It takes a massive amount of energy to rip one off.

Small atom equals high ionization energy.

And if an atom is big?

Like cesium.

The valence electron is far away.

It's highly shielded by all those core layers.

It's loosely held.

You can pluck it off with very little energy.

Big atom equals low ionization energy.

So, broadly speaking, bottom left, like francium, is easy to ionize.

Top right, like helium, is incredibly hard to ionize.

The trend increases as you go up and to the right.

That represents about 90 % of the data.

But, and you know chemistry loves exceptions.

If you look at the graph in Figure 914, the line isn't perfectly smooth, has these little zags.

There are two specific places where the ionization energy drops when the general trend says it should go up.

And professors love to ask why on exams.

Let's identify them.

Zag number one.

Group two versus group 13.

Specifically, magnesium versus aluminum.

Magnesium is to the left.

Aluminum is to the right.

Based on the general trend, aluminum should be smaller and therefore harder to ionize.

But the data says aluminum is actually easier.

It's 578 kilojoules per mole versus 738 for magnesium.

Why?

You have to look at the quantum orbitals.

Magnesium ends in 3S2.

The subshell is full.

That is a stable, happy paired set of electrons.

Aluminum ends in 3S2, 3P1.

It has one lonely electron sitting in the 3P orbital.

The orbital is slightly higher in energy than the S orbital.

It extends a little further out from the nucleus.

And this is the key.

The two S electrons actually provide a little bit of extra shielding for that single P electron.

So that single P electron in aluminum is basically low hanging fruit.

Exactly.

It's significantly easier to snip off that exposed higher energy P electron than it is to break up the stable, full spare, and magnesium.

So the energy cost temporarily drops.

Zag number two.

Group 15 versus group 16.

Phosphorus versus sulfur.

Again, sulfur is to the right.

It should be harder to ionize, but it's easier.

Is this about subshells again?

No, because they are both actively filling the P subshell.

Phosphorus is 3P3, sulfur is 3P4.

This exception is about electron pairing energy.

Look at phosphorus, 3P3.

According to Hunn's rule from chapter 8, those three electrons sit in three separate P orbitals.

Up, up, up.

Everyone has their own room.

Repulsion is minimized.

Very comfortable living situation.

Very.

Now look at sulfur, 3P4.

The fourth electron has no empty room left.

It has to pair up with one of the other electrons.

Electrons are negative.

They repel.

When you force two of them into the exact same tiny orbital space, that repulsion creates instability.

It's like two people trying to share a twin bed.

One of them is very, very ready to leave.

So it's easier to remove that roommate electron because the natural repulsion is already pushing it out the door.

Exactly.

That extra repulsion lowers the ionization energy of sulfur just below phosphorus.

We also need to talk about successive ionization energies.

We talked about removing one electron.

Can we remove two, three?

You can remove as many as you have, assuming you have a big enough laser, but it gets harder every single time.

Removing the first electron is removing a negative charge from a neutral atom.

Removing the second electron is removing a negative from a positive ion.

That attraction is stronger.

Removing the third is removing a negative from a double positive, even harder.

But the text highlights a specific point where the energy cost doesn't just go up incrementally.

It explodes.

The core jump.

Let's look at magnesium again.

To remove the first electron, the 3S, it costs 738 kilojoules.

To remove the second electron, the other 3S, it costs 1 ,451 kilojoules, roughly double.

What about the third?

The text says 7 ,733 kilojoules.

That is a 500 % increase.

That is a brick wall.

Why?

Because the first two were valence electrons.

They were in the n equals three shell.

The third electron you are trying to remove is in the n equals two shell.

It is a core electron.

It is part of the completely stable noble gas configuration of neon.

Breaking a noble gas core requires a massive amount of energy because the ZF jumps wildly when you get that close to the nucleus without any outer shielding.

And this explains the macrochemistry we see every day.

Magnesium forms Mg2 plus ions.

It basically never forms Mg3 plus ions.

Not because it's theoretically impossible, but because the universe demands a payment of 7 ,700 kilojoules.

And normal chemical reactions just don't have that kind of cash lying around.

Exactly right.

The energy economics absolutely forbid it.

Let's flip the script.

Section nine to five.

Electron affinity.

Ionization was about taking electrons away.

Electron affinity, EEA, is about giving them.

The formal definition is the energy change that occurs when an electron is added to a gaseous atom.

You'll listen carefully to this part.

Sign convention.

In thermodynamics, if a system releases energy, meaning it becomes more stable,

the sign is negative.

Exothermic.

If a system absorbs energy, meaning you have to force the process to happen,

the sign is positive.

Endothermic.

So, if an atom really wants an electron, the affinity value will be a large negative number.

Correct.

So, who wants electrons the most?

The halogens.

Group 17.

Fluorine, chlorine.

They are one step away from the octet.

If you offer them an electron, they grab it and release a huge amount of energy and relief.

Chlorine actually has the most negative electron affinity on the entire table, minus 349 kilojoules per mole.

Who doesn't want electrons?

The noble gases.

Group 18.

They are totally full.

S2P6.

If you're trying to give helium an electron, it says, I have no room in level one.

You have to put this in level two.

Level two is a much higher energy state.

You have to pay energy to put it there.

So, noble gas affinities are positive.

What about group two?

The alkaline earths.

Beryllium, magnesium.

They have filled subshells, S2.

To add an electron, you have to open up the P subshell.

Again, that's higher energy.

So their affinities are also positive, or near zero.

They just don't want it.

Now, here is the tricky one from the chapter, oxygen.

Oxygen forms O2 - ions all the time.

It loves electrons, right?

Well, yes and no.

Adding the first electron to oxygen, making O-, is exothermic.

It releases energy, minus 141 kilojoules.

That's good.

But adding the second electron, going from O - to O2-, You are trying to push a negative electron onto an ion that is already negative.

It's like pushing the north poles of two strong magnets together.

Massive repulsion.

To add that second electron, you have to do significant work.

The second electron affinity of oxygen is strictly positive.

Plus 744 kilojoules.

It is heavily endothermic.

This blows people's minds.

It costs a ton of energy to make an oxide ion.

Then why is the earth covered in rocks made of oxydons?

Why does magnesium oxide even exist?

Because of what happens after the ion is formed.

Oxygen isn't just soaring alone in a void.

It is next to magnesium.

You have a highly charged Mg2 +, and a highly charged O2 -.

They slam together to form a solid crystal lattice.

That snap forming the ionic bonds releases a colossal amount of energy.

It's called lattice energy.

That energy releases so huge, thousands of kilojoules, that it completely pays off the initial debt of forming the oxide ion and leaves a massive profit.

It's an investment.

You pay 700 kilojoules to make the ion, but you get back 3000 kilojoules when the crystal forms.

Exactly.

Chemistry is all about the net bottom line.

Moving on to section nine to six, magnetic properties.

This is a shorter section, but it gives us a really direct window into the quantum world.

It all comes down to electron spin.

An electron spinning creates a tiny magnetic field.

It's a micro magnet.

We have two main categories of materials here, diamagnetic and paramagnetic.

Let's define diamagnetic first.

In a diamagnetic atom, all the electrons are perfectly paired up.

For every up spin, there is a down spin in the same orbital.

The magnetic fields cancel each other out completely.

The result.

The atom has no net magnetic field.

In fact, it is weakly repelled by an external magnet.

On paramagnetic.

You have unpaired electrons.

Bachelorette electrons.

Their magnetic fields don't have a partner to cancel them out, so they add up.

If you put a paramagnetic substance near a magnet, it is attracted.

The more unpaired electrons you have, the stronger the attraction.

So on the exam, they will ask, is this specific atom paramagnetic or diamagnetic?

How do we solve it?

You have to draw the orbital diagram.

You have to physically see the boxes.

Example.

Nitrogen.

Configuration is helium core.

Two is two.

Two P three.

Look at the two P three electrons.

Hans rule says separate orbitals first.

Up, up, up.

Three unpaired electrons.

Verdict.

Paramagnetic.

Let's do another.

Zinc.

Configuration is argon core.

Three D ten.

Four is two.

The D subshell has ten electrons, which is five pairs.

Full.

The subshell has two electrons, one pair.

Full.

Zero unpaired electrons.

Verdict.

Diamagnetic.

Right.

Now let's go back to our transition metal rule.

Manganese two ion.

MN two plus.

If you fill for the trap and remove the D electrons, you'd have argon core three D three four S two.

You'd see three unpaired in the D and zero in the S.

You'd write down three unpaired electrons.

And you would be wrong.

You'd lose the points.

The correct way is to remove the four S electrons first.

Result.

Argon core three D five.

Five electrons in the D subshell.

Up, up, up, up.

Five unpaired electrons.

MN two plus is highly paramagnetic.

And the beautiful thing is we can verify this in the lab.

We can actually weigh a sample in a magnetic field.

The physical data confirms the S first rule.

It's experimental proof that our quantum model is correct.

That is so satisfying.

Final property.

Section nine to seven, polarizability.

This sounds like, you know, shields up, Captain.

It does sound very sci -fi, but I want you to replace the word polarizability with a much simpler word, squishiness.

Squishiness, I like it.

Very technical.

We draw atoms as perfect hard seers in textbooks.

But they are really loose clouds of probability.

If I put an atom in an electric field, say, I bring a highly positive k -dimedition near it.

The electron cloud reacts.

The negative electrons get attracted to the positive charge.

The whole cloud shifts.

It distorts.

It bulges out on one side.

That distortion is called polarization.

Polarizability is just a measure of how easy it is to distort that cloud.

And what determines the squishiness?

Size and control.

Think about fluorine.

Very small.

Very high ZF.

The nucleus has an absolute death grip on those valence electrons.

That cloud is like a hard rubber ball.

It does not distort easily.

Low polarizability.

Think about iodine at the bottom of the halogen.

Huge.

The valence electrons are way out in the boonies.

The nucleus barely has a hold on them because of all the shielding.

It's like a giant marshmallow.

You bring a positive charge near it and the whole electron cloud just slashes over.

High polarizability.

So big equals squishy equals highly polarizable.

Small equals hard equals low polarizability.

And anions.

Anions are extra puffy because of that electron repulsion we talked about.

So anions are generally very polarizable compared to neutral atoms or arcations.

This concept might seem small now, but it is going to come back with a vengeance in chapter 11 when we talk about intermolecular forces, specifically London dispersion forces.

The reason gasoline is a liquid and not a gas at room temperature is entirely due to this squishiness.

We have covered the map.

We have covered the trends.

The chapter ends with an integrative example.

I love these because they force us to combine everything we just learned.

The text asks us to look at a specific series, lithium beryllium plus, boron two plus, carbon three plus.

Notice the pattern there.

Lithium has three protons, three electrons.

Its valence is two m's one.

Beryllium plus has four protons, three electrons.

Boron two plus has five protons, three electrons.

Valence is two m's one.

They're all hydrogen -like in their valence shell.

They all just have one single electron in the outermost shell.

And the book plots the square root of the ionization energy versus the atomic number Z.

And it gets a perfectly straight line.

This is basically a victory lap for Niels Bohr.

Remember chapter eight.

The Bohr model predicted that for a one -electron system, the energy is proportional to Z squared.

So logically, if you take the square root of the energy, it should be directly proportional to Z.

The straight line on that graph proves that the Bohr model, as simple as it is, actually captures the fundamental physics of how charge and energy relate.

It even works for multi -electron atoms if you replace Z with ZF.

The graph shows that our stories, the screening, the effective charge, the shells, they aren't just, you know, convenient stories.

They are mathematical descriptions of reality.

Which is the entire goal of science.

Oh, okay, we're at the finish line.

We've gone from Mendeley's blank cards to this squishiness of iodine.

Let's summarize.

What is the main takeaway for the student walking into the exam hall tomorrow morning?

The main takeaway is trust the map.

Do not just memorize that chlorine has a high electron affinity.

Ask yourself why.

Chlorine is top right.

Top right means high ZF.

High ZF means the nucleus pulls really hard.

Strong pull equals small radius.

Strong pull equals high ionization energy.

Strong pull equals high electron affinity.

Strong pull equals low polarizability.

It's all one drafted web.

It is.

The underlying engine driving everything on the table is just Coulomb's law.

Positive attracts negative.

Everything else we talked about is just details about distance and shielding.

I wanna leave you with one final thought.

We talked a lot about Mendeley if predicting elements that didn't even exist yet.

The crazy thing is we are still doing this.

If you look at the very bottom of the table, row seven, elements 113 through 118, Nihonia, Muscovium, Tennessean, Ogenessen, these were only officially named recently.

Right, and we are actively hunting for a 119 and 120 right now.

These elements do not exist in nature.

We have to build them in massive particle accelerators, literally atom by atom, and they only last for fractions of a millisecond before decaying.

But because of the periodic law we just learned, we know their chemistry before they're even born.

We know element 119 will be an alkali metal.

It will sit right below Francium.

It will be the most reactive, largest, most explosive metal ever created.

We are mapping the coastlines of continents that we haven't even visited yet.

That is the true power of the periodic table.

It's the treasure map of the universe.

Thank you for listening to this last minute lecture.

If you have an exam tomorrow, take a breath, draw your orbital diagrams, watch out for the transition metal trap.

You've got this.

Good luck out there.

Go crush it.

We'll see you next time on the Deep Dive.