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The Chemistry of Solids

  • The common unit cells for solids are primitive cubic, body-centered cubic, and face-centered cubic.
  • The unit cell structure can influence a solid’s density and texture.
  • Lattice energy is used to quantify the strength of the interaction within a crystalline solid.
  • Phase diagrams are used to graphically predict a substance’s physical state at a set of pressure and temperature conditions.

The Chemistry of Solids

Lecture Slides are screen-captured images of important points in the lecture. Students can download and print out these lecture slide images to do practice problems as well as take notes while watching the lecture.

  • Intro 0:00
  • Lesson Overview 0:07
  • Introduction 0:46
    • General Characteristics
    • Particulate-level Drawing
  • The Basic Structure of Solids: Crystal Lattices 1:37
    • The Unit Cell Defined
    • Primitive Cubic
  • Crystal Lattices cont'd 3:58
    • Body-centered Cubic
    • Face-centered Cubic
  • Lattice Enthalpy and Trends 6:27
    • Introduction to Lattice Enthalpy
    • Equation to Calculate Lattice Enthalpy
  • Different Types of Crystalline Solids 9:35
    • Molecular Solids
    • Network Solids
  • Phase Changes Involving Solids 11:03
    • Melting & Thermodynamic Value
    • Freezing & Thermodynamic Value
  • Phase Changes cont'd 12:40
    • Sublimation & Thermodynamic Value
    • Depositions & Thermodynamic Value
  • Phase Diagrams 13:40
    • Introduction to Phase Diagrams
    • Phase Diagram of H₂O: Melting Point
    • Phase Diagram of H₂O: Normal Boiling Point
    • Phase Diagram of H₂O: Sublimation Point
    • Phase Diagram of H₂O: Point C ( Supercritical Point)
  • Phase Diagrams cont'd 16:31
    • Phase Diagram of Dry Ice
  • Summary 18:15
  • Sample Problem 1, Part A: Of the Group I Fluorides, Which Should Have the Highest Lattice Enthalpy? 19:01
  • Sample Problem 1, Part B: Of the Lithium Halides, Which Should Have the Lowest Lattice Enthalpy? 19:54
  • Sample Problem 2: How Many Joules of Energy is Required to Melt 546 mg of Ice at Standard Pressure? 20:55
  • Sample Problem 3: Phase Diagram of Helium 22:42

Transcription: The Chemistry of Solids

Hi, welcome back to

Today's lesson from general chemistry is going to be on the chemistry of solids.0002

We are going to go ahead and start off with a very brief introduction followed by the basic structure of solids.0008

We will get into a few thermodynamic properties, namely the lattice enthalpy,0015

followed by different types of crystalline solids that we can encounter.0022

Then we will something characteristic of solids, liquids, and gases--that is phase changes.0027

Phase changes are going to then lead us into namely phase diagrams.0035

We will finish up the lesson talking about our summary followed by a few sample problems.0039

We have briefly alluded to what we meant by a solid a long0050

time ago when we first did our first one or two lessons.0054

If you recall, solids have the following characteristics; that they are relatively dense.0060

They have a definite shape and a volume; they have very strong attractive forces.0065

If we look at it on a particulate level, we can imagine solids0070

as being atoms or spheres that are closely spaced together like so.0075

The separation that we have in gases and liquids, we basically do not have in solids.0085

That is due to the very strong attractive forces.0090

Now that we have talked about and we reviewed what we0096

know about solids, let's get into how solids look three dimensionally.0101

The basic part of a solid is what we call the unit cell.0107

The unit cell is basically a pattern of a solid that is repeatable in all directions, in all three directions.0113

As you can see from this representative diagram, the unit cell is the basic cube here.0132

You see that this cube is repeatable in all directions.0139

That is what we call the unit cell.0144

A lattice point is basically a fixed location of an atom that composes a solid; a fixed location of an atom.0146

You see that because it is part of a unit cell, this fixed location is reproducible also in all three directions.0162

We are going to learn that there are three basic types of unit cells.0171

The first is what we call primitive cubic.0175

We are going to find that primitive cubic is actually very rare.0178

There is only really a handful of elements, really single-digit number of elements that have this structure.0181

But primitive is your basic cube like so; it is if you will completely empty.0187

Once again primitive cubic is right here.0199

You basically have lattice points at all corners of the unit cell.0205

Each corner actually contributes only one-eighth of the sphere.0220

Because there are eight corners here, we say that we get one net atom per unit cell in primitive cubic structure.0227

Let's now jump into body centered cubic; body centered cubic is what we abbreviate bcc.0238

Basically for body centered cubic, we take that primitive cubic like so.0245

Let me go ahead and redraw it here.0251

We are going to take that primitive cubic.0253

But in addition to having the spheres at the corners, we also have0257

one sphere that is right in the middle of this cube itself.0268

Not only do we have one-eighth by eight, we also have the entirety of the0274

one sphere in the middle giving us two net atoms per unit cell for bbc.0281

Bcc solids have the characteristic of being brittle.0293

The final unit cell is what we call fcc.0304

For fcc, I am going to go ahead and redraw that primitive cubic again.0309

But for fcc which is called face centered cubic, we are going to0323

have a sphere at the center of each face of this unit cell.0326

There is going to be one right here.0330

There is going to be one right here sticking inward.0333

There is going to be one right here sticking inward.0336

One right here sticking inward; one on the bottom sticking inward.0338

Of course one on the back sphere sticking inward.0342

We have one-eighth by eight from the corners.0347

It turns out that each of the spheres in the center of each face contribute one-half.0351

There is going to be six sides, six faces.0357

We get four net atoms per unit cell.0362

It turns out that what is characteristic of fcc solids is that these solids tend to be the opposite of bcc.0370

They tend to be malleable.0379

Now that we have our brief introduction into crystal lattices, let's go on to a thermodynamic approach.0384

This is what we call lattice enthalpy.0391

Lattice enthalpy is the energy that is released.0394

Again this is released when a solid crystal forms from its constituent0398

elements in the gas phase represented by the following chemical equation.0403

If we take a crystalline solid maybe of the form MX,0407

it is formed from M+ gas and X- gas.0417

The correct equation will be M+ gas plus X- gas goes on to form MX solid.0423

We are going to see that this is what we call the lattice enthalpy0433

which is therefore ΔHlat--is how some textbooks abbreviate it.0436

Basically because mother nature favors low energy, the lower or the more negative0442

the lattice enthalpy, the more favorable the formation of this crystalline solid.0447

Let's examine the equation to calculate its lattice enthalpy.0451

You are going to find that the lattice enthalpy is basically going to be analogous to Coulomb's law.0455

This is going to be analogous to q1q2 over r where q1 and0463

q2 are the charges of the cation and anion respectively; cation/anion charges.0471

r is going to be the internuclear distance between cation and anion adjacent; internuclear distance.0482

We see that as the magnitude of the charges increase...0495

Once again as the magnitude of q1 and q2 goes up,0505

ΔH is going to become more negative which is favorable.0511

That makes sense.0516

As the charges goes up, there is a stronger attractive force between opposite charges.0516

It should be more favorable that the solid is going to form.0522

We also see that as r goes up, ΔH is actually going to become less and less negative.0526

It is actually going to go up.0535

Basically this is saying that the farther apart ions are,0537

the weaker the attractive force between the two nuclei which makes sense also.0541

Basically ΔH negative is going to be given by the following0547

combination of parameters--smaller ions and also ions with greater charge.0556

Smaller ions, greater charge tend to be indicative of a solid that is going to form in the gas phase.0567

We also now are going to focus on crystalline solids.0578

Crystalline solids are those that have these different unit cells such as bcc and fcc.0582

There are different types of crystalline solids; number one is molecular solids.0589

Molecular solids consist of molecules held together by relatively weak0594

intermolecular forces that influence the physical properties, especially melting point.0598

Common examples include ice and dry ice.0603

When you think of ice and dry ice, these don't have really high melting points.0606

Ice melts relatively at a pretty low temperature; dry ice readily sublimes.0611

Molecular solids tend to have relatively low melting and boiling points.0619

Like ice and dry ice, they tend to be brittle.0623

Network solids are essentially the opposite.0627

They consist of molecules that are held together by very strong covalent forces and covalent bonds.0630

They form a framework extending throughout the crystal.0637

Common examples include diamond and graphite.0639

Make sure you guys understand the importance of diamond is that it is the hardest known naturally occurring substance.0644

Network solids tend to be hard and rigid with very very high melting points in the thousands of degrees Celsius.0653

Now that we have talked about the different types of crystalline solids, let's go into phase changes that involve solids.0663

Melting, melting is basically the process in which the intermolecular forces that hold a solid together are overcome and0669

broken, resulting in greater separation between the atoms themselves such that we turn the solid into a liquid.0676

The enthalpy is typically abbreviated as ΔH of fusion under standard conditions.0684

This is going to be an endothermic process because this requires energy.0693

This tends to be positive.0698

Because these are all enthalpies, these are typically in units of kilojoules per mole.0703

Freezing, freezing is basically the exact opposite.0710

It is such that we lower the temperature of our liquid such that atoms can come into contact with each other.0713

There is less motion of the atoms.0719

When they touch each other, they don't have as much time to escape even partially from neighboring atoms.0722

We go into the solidification process.0729

Basically ΔH of freezing is going to be the same magnitude but opposite sign of ΔH of fusion.0732

We see that ΔH of freezing is going to be negative.0744

ΔH of freezing is an exothermic process whereas ΔH of fusion is an endothermic.0750

That is melting and freezing.0758

Let's now go on to the next two; this is sublimation.0760

Sublimation is the process in which a solid coverts directly to a vapor0763

from the solid state at a cold enough temperature and low enough pressure.0769

Because this is going solid to a gas directly, this is going to be an endothermic process0774

because we have to require energy to overcome the very strong attractive forces in the solid.0782

ΔH of sublimation is going to be a negative value.0787

Finally deposition is the process when a vapor converts directly to a thin solid layer.0794

The thermodynamic value is going to be ΔH of deposition.0801

This is going to be equal to negative ΔH of sublimation.0806

We see that deposition is going to be an exothermic process.0811

That is sublimation and deposition.0817

Now that we have talked about the different phase changes that can occur and we have really0821

completed our discussion on all three of the physical state solids, liquids, and gases, let's culminate everything.0826

When we try to culminate everything, what we get is a phase diagram.0833

We can graphically predict when a substance will be a solid, liquid, or gas at any given temperature and pressure combination.0837

When we go ahead and graph this for a substance, when we graph pressure0845

as a function of temperature, we get what is called a phase diagram.0848

Let's take a look at the diagram for water; there is several points of interest.0853

Number one, the melting point; this is called the normal melting point.0857

Normal melting point, anytime you see the word normal, that is always at 1 atm, at standard pressure.0862

We see that at a pressure of 1 atm, we can go ahead and translate it over.0871

When the solid becomes a liquid, that is what we call melting.0877

The melting point is the 0 degrees Celsius that we know for water.0883

The normal boiling point, we just translate along.0888

You see that the liquid to gas transition is 100 degrees Celsius.0893

That is where water boils at standard pressure.0899

The sp, sp is what stands for sublimation point.0904

The sublimation point as you can see for ice is right here.0911

As you can see, this is going to be at a very cold temperature and a very very low pressure0919

which is why we don't see ice readily sublime under normal ambient conditions; sublimation.0922

There is another point of interest here that we want to talk about--that is point C.0932

Point C is the absolute highest temperature where a liquid can exist.0936

This is what we call the super critical point.0944

At the super critical point or the critical point, you cease becoming a liquid and you have nothing but gas.0950

We say that gas overtakes a liquid state.0957

Finally there is a really nice point of interest right here which is0965

what we call the intersection of the solid, liquid, and gas phase.0968

That is called the triple point.0973

Only at this one unique set of pressure and temperature is where we get coexistence of all three physical states.0975

Once again that is what we call the triple point.0984

That is what we call a phase diagram; that is for water.0988

Let's look at another phase diagram, this time with dry ice.0992

Dry ice we know readily sublimes.0995

Anytime you go to see maybe a theatrical production, you sometimes see a fog being generated.0997

This is basically one source of the fog generated is going to be dry ice being blown out to the stage.1003

Dry ice we expect to readily sublime.1009

Let's see if we can detect that in the phase diagram.1012

At 1 atm which is what we are typically at, yes we can go ahead and translate over.1016

Boom, we hit point A.1023

At point A, represents the normal sublimation point for dry ice.1025

You see that when you go down, the temperature is only 197.5 kelvin.1032

That is definitely less than room temperature.1039

Dry ice readily sublimes; that is why it produces a high vapor pressure.1043

Just by looking at a phase diagram, you can easily see that dry ice readily sublimes.1050

Also look how large the sublimation area is; compare that to that of water.1056

The sublimation interface between the solid and gas is quite prominent in the phase diagram compared to that of water.1062

Sublimation, a solid to gas interface curve very prominent in dry ice.1073

That is the phase diagram of a different compound versus water.1092

Now let's go ahead and summarize our brief session on the chemistry of solids.1096

The common unit cells for solids are primitive cubic, body centered cubic, and face centered cubic.1102

The unit cell structure we see can influence a solid's density and texture.1107

We saw that lattice energy quantifies the extent or tendency for a molecular1112

or network solid to form from its ions in the gas phase.1121

Finally we see that phase diagrams plot pressure versus temperature.1125

They allow us to predict how a compound is expected to exist,1130

as a solid, a liquid, or as a gas at any set of these conditions.1136

Let's go ahead and look at sample problem one.1142

Of the group 1 fluorides, which should have the highest lattice enthalpy?1144

When we mean highest lattice enthalpy, we really mean most negative or the most exothermic.1149

Of the group 1 fluorides, LiF, NaF, KF, etc; the only difference is atomic size.1160

Lithium is going to be the smallest cation here.1172

The smallest cation results in a smaller r value which results in a more negative ΔH.1178

The lithium fluoride is definitely your answer for that one.1191

Of the lithium halides, which should have the lowest lattice enthalpy here?1195

Lowest lattice enthalpy is going to be...1199

Sorry about that, this should be the most negative.1206

Let me go ahead and rephrase that; most negative lattice enthalpy.1209

This should also be the most negative; very sorry about that.1213

Lithium halides, LiF, LiCl, LiBr, and LiI.1220

Once again the only difference is the atomic size.1227

Fluorine is going to have the smallest r value.1231

The smallest r value is going to give you the more negative ΔH.1236

This is a nice sample problem which illustrates the periodic trends and its effect on the value of ΔH.1244

Sample problem two, how many joules of energy is required to melt 546 milligrams of ice at standard pressure?1256

You are given a ΔH of fusion for ice to be 6.01 kilojoules per mole.1263

Basically we are going to use kJ per mole as a conversion factor.1268

Basically we want to go from milligrams of ice to moles of ice and1279

then from moles of ice to what the question is asking for--to joules.1287

As you can see, we can use kilojoules per mole as a nice convenient conversion factor for the moles to cancel.1292

Let's go ahead and do this.1301

546 milligrams of ice times 10-3 grams divided by 1 milligram times1303

1 mole of the ice which is water divided by its molar mass 18.016 grams.1315

Then times 6.01 kilojoules divided by 1 mole.1323

Finally we have to go from kJ to regular joules.1329

We are going to take all of that.1333

We are going to multiply it by 103 joules over 1 kilojoule.1334

That is going to give us our answer in units of joules.1338

There is nothing really special about this problem.1342

It is basically just using the dimensional analysis with ΔH of the phase change as a conversion factor.1344

Dimensional analysis is something we have done so many times.1353

Again this is just nothing new.1357

We are just introducing a new type of conversion factor.1358

Finally onto sample problem three which involves a phase diagram.1362

Below is the phase diagram for helium; we know helium is a noble gas.1367

At standard temperature and pressure, we expect a gas to exist.1376

We can also spot that; it makes sense of that.1381

What is the normal bp for helium?--let's locate 1 atm or 1 bar.1384

The boiling point is when a liquid goes into the gas.1391

When we translate over, here is our transition between liquid to gas right there.1393

When we translate down, our normal boiling point is right there which is 4.22 in this case kelvin.1401

The next question is can solid helium sublime?1411

What we want to look for is the interface between a solid and a gas.1416

You see that it is right there going from a solid to a gas.1422

Absolutely liquid helium can sublime.1426

However it is going to be not readily occurring because we as you can see from1429

this curve, it only occurs at super low pressures, at only very low pressures.1435

What is the maximum pressure at which solid helium can exist?1447

Let's go ahead and look at the solid curve.1452

The solid curve here is strictly ending right there.1457

The highest pressure at which solid helium can exist is only at 100 bar.1463

Anything after that, we are going to become a liquid.1471

Finally at what set of conditions do all three states coexist?1476

Again that is what we call the triple point.1481

At the triple point which is right there, we are going to be1484

at a pressure maybe approximately 0.05 bar and a temperature of 2.17 kelvin.1488

That is a nice problem that illustrates a phase diagram.1501

I want to thank you all for your attention.1507

I will see you next time on