Browse Courses
Start learning today, and be successful in your academic & professional career. Start Today!
Loading video...
This is a quick preview of the lesson. For full access, please Log In or Sign up.
For more information, please see full course syllabus of General Chemistry
  • Discussion

  • Study Guides

  • Download Lecture Slides

  • Table of Contents

  • Transcription

  • Related Books & Services

Bookmark and Share
Lecture Comments (15)

1 answer

Last reply by: Professor Franklin Ow
Wed May 7, 2014 6:44 PM

Post by Edgar Suarez on May 5, 2014

I have a question, you said that milk was basic, but i thought that it was an acid, because of its lactic acid.

1 answer

Last reply by: Professor Franklin Ow
Sat Apr 26, 2014 5:17 PM

Post by Anthony Mendoza on April 25, 2014

In example 2, how did you know that HPO4 was the conjugate of weak acid H3PO4? In my head, I think I might have chosen that it was the conjugate of -PO4, or conjugate of H2PO4...

1 answer

Last reply by: Professor Franklin Ow
Sat Apr 26, 2014 5:16 PM

Post by Anthony Mendoza on April 25, 2014

Hi Professor Ow
,
Could I get some clarification on how increased electronegativity makes it a stronger acid? HF has a strong electronegativity difference, as in the F- pulls strongly on the H+, so I would imagine it would't want to let it go (and thus dissociate and thus not be that strong).

I'm trying to compare HF to the others on the list, but it's difficult because aren't the rest bases. They already have a negative charge, so obviously aren't going to want to dissociate. But how about comparing HF to HCl. HCl is a strong acid, yet has lower ElectroNegativity (since F- is supposed to be the most EN?). This part still confuses me.

Thanks for your help! Your explanations are making class a lot more manageable!

1 answer

Last reply by: Professor Franklin Ow
Sat Apr 26, 2014 5:15 PM

Post by Melyssa James on April 2, 2014

For your first example, after you determined the pH, i don't see where you calculated the percent ionization??

2 answers

Last reply by: Professor Franklin Ow
Fri Feb 28, 2014 11:40 AM

Post by Jack Miars on February 26, 2014

Can you always make the assumption that x is negligible? If not, what situations cause it to be accounted for?

1 answer

Last reply by: Professor Franklin Ow
Mon Feb 10, 2014 10:58 AM

Post by Miley Xiao on February 9, 2014

what does x at 13:38 stands for? and why are CH3CO2 and H3O the same since he wrote two xs?

0 answers

Post by Cheng Jiang on May 16, 2013

very gucci

0 answers

Post by Max Mayo on April 2, 2013

great job

Related Articles:

Acid-Base Chemistry

  • Bronsted-Lowry acid-base chemistry involves a loss/gain of a proton to/from water.
  • Conjugate pairs only differ by one proton and are inversely related in terms of acidity/basicity.
  • Acidity and basicity can be quantified primarily via pH, Ka ,pKa and percent ionization.
  • The structure of a molecule can have a significant impact on how acidic it can be.
  • ICE tables are easily applied to an acid-base equilibrium situation.

Acid-Base Chemistry

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.

Transcription: Acid-Base Chemistry

Hi, welcome back to Educator.com.0000

Today's lesson from general chemistry is on acid base chemistry.0002

We are going to first start off as usual with our introduction followed by how to quantify acid and base strength.0007

After we learn how to quantify how acidic or how basic something is, we will then go on to0015

writing out aqueous acid base equilibria reactions which is going to become very fundamental for this section.0022

After that, we will learn something we call acidic and basic salts0028

followed by multiple aqueous equilibria which is for acids that have more than one hydrogen.0031

We will then get into some specific topics, namely Lewis acids and bases, followed by molecular structure and acidity.0038

We will wrap up the session with our summary followed by a pair of sample problems.0047

Let's go ahead and begin.0054

We are going to look at two different types of molecules in this chapter--what we call a Bronsted-Lowry acid and a Bronsted-Lowry base.0057

A Bronsted-Lowry acid basically donates a proton to water.0065

A Bronsted-Lowry base is going to accept a proton from water.0076

You see that water is involved in both of these reactions.0088

For a Bronsted-Lowry acid, you can say HA aqueous plus water goes on to form A- aqueous because it loses its proton.0092

It is going to give it to water; H2O then becomes H3O1+ aqueous.0105

This special ion is what we call the hydronium ion.0113

Anytime you see the hydronium cation, it is always indicative of acid.0119

Let's go ahead and look at the reaction for a base.0125

We can have a generic base B aqueous plus H2O liquid goes on to form...0127

The base is going to accept the proton; it becomes BH1+ aqueous.0136

H2O loses a proton; it is going to become OH1- aqueous.0141

This special guy here of course is what we call the hydroxide ion.0145

This is always indicative of aqueous base.0150

Once again hydronium cation is indicative of aqueous acid.0153

Hydroxide anion is indicative of aqueous base.0157

You see that water is involved in both of the reactions.0161

In the acid reaction, water is the base.0166

In the base reaction, water is the acid.0169

We say that water is an amphiprotic molecule meaning it can function both as an acid and as a base.0172

When we take a look at water reacting with itself, we see the dual role of water.0179

For example, one of these water molecules can lose a proton giving hydroxide.0186

Therefore the other water molecule is going to gain a proton forming hydronium.0191

What is the extent of this reaction?0198

It turns out that if we write out the equilibrium expression for this reaction, it is going to be0202

equal to the hydroxide ion concentration at equilibrium times the hydronium ion concentration at equilibrium.0207

It turns out that this value, this product is only 1.0 times 10-14 at 25 degrees Celsius.0216

We give this K a very special name; this is what we call Kw.0225

This is called the auto-dissociation constant for water.0230

It is a very handy one because if you know the hydroxide ion concentration, you can calculate hydronium and vice versa.0236

When we talk about acids and bases, there are going to be what we call strong and weak acids and bases.0247

Basically strong acids are going to donate a proton to water and fully dissociate into a proton and an anion.0253

Because they fully dissociate, we do not use an equilibrium arrow.0262

We only use an arrow pointing in the forward direction.0267

We can take a molecule HX aqueous; this is going to react with water.0273

We only use a single arrow to show full dissociation.0278

That is going to form H3O1+ aqueous and X1- aqueous.0281

It is very imperative that you know what are the strong acids.0287

For our purposes right now... you should always confirm this with your instructor.0292

But there are going to be seven strong acids that you should know.0295

HClO4, HClO3, HNO3, H2SO4, HCl, HBr, and HI.0299

Again these are the seven strong acids that fully dissociate when dissolved in water.0313

The strong bases also are going to become fully protonated.0319

They are going to fully dissociate when dissolved in water.0324

They are going to form a cation and a hydroxide.0328

The typical ones are going to be MOH or MOH2.0332

These are basically group 1 and 2 hydroxides.0340

We can take the typical one, sodium hydroxide.0349

You don't even have to show water because all it does is that0353

it is going to fully break apart into sodium ions and hydroxide anion.0356

Once again these are the strong bases and the strong acids that you should know.0362

Again check with your instructor because every instructor is going to be0368

slightly different on the ones he or she requires you to memorize.0372

We have talked about strong species, what are the weak species.0378

Basically weak species, weak acid and bases do not completely dissociate at all, do not completely dissociate at all.0381

For example, HF aqueous plus water is a weak acid.0389

We are going to use of course an equilibrium arrow just like we do for any type of weak electrolyte.0396

That is going to give us hydronium cation plus F1- aqueous.0403

That is the typical weak acid.0410

The typical weak base is going to going to be ammonia NH3 aqueous plus H2O liquid.0412

Once again equilibrium arrow.0420

That goes on to form NH41+ aqueous and hydroxide aqueous.0421

Anything related to ammonia which is what we call... this is ammonia/amines.0429

These are all going to be our typical weak bases.0439

Again these do not completely dissociate.0449

Let's go ahead and now see how to quantify acidity and basicity.0452

To quantify acid and base strength, we talk about what is known as the pH scale.0457

When we talk about the pH scale, this is on a scale of 0 to 14 where 7 is the neutral area.0464

Any pH less than 7 is acidic.0476

Any pH greater than 7 is going to be basic.0480

The equations for pH is the following.0483

pH is equal to ?log of concentration of hydronium at equilibrium.0485

pOH is equal to ?log of the concentration of hydroxide at equilibrium.0492

There is some common things that you can know that are acidic--citric juice, sodas, stomach acid.0501

Then there is some things that are basic that you should know also.0513

Any household detergents tend to be basic.0516

Milk and human blood also is slightly basic.0519

You have to also make sure that you be able to go both ways.0528

Sometimes you are asked to calculate the hydronium ion concentration at equilibrium.0532

That is going to be anti-log.0537

That is going to be log inverse of ?pH which is 10 raised to the ?pH.0538

The hydroxide ion concentration at equilibrium is equal to log inverse of ?pOH which is equal to 10 raised to the ?pOH.0548

There is going to be a relationship between pH and the pOH.0563

Basically the pH plus the pOH is going to be equal to 14.00 at 25 degrees Celsius.0566

Again you should always consult with your instructor whether or not you have to commit these to memory.0577

This is the pH scale that is typically used for quantifying acid and base strength.0589

Let's now move on to another way of quantifying acid base strength.0596

This is going to be something we are going to encounter quite a bit.0599

This is what we call the acid ionization constant Ka and pKa.0603

For any of the acids where HA reacts with water, you are going to form hydronium and A1-.0608

There is always going to be a certain extent to which this acid is going to dissociate.0624

We are going to call this a Ka.0633

This is just the same equilibrium expression that we have been working with so much already.0636

Remember it is going to be products, hydronium ion, at equilibrium times0640

A1- at equilibrium divided by the concentration of HA at equilibrium.0645

You can see that for stronger acids, the concentration of hydronium at equilibrium is going to be high.0652

Ka is going to be high also.0665

This is what we call the acid ionization constant Ka.0673

Ka is usually written in scientific notation.0678

But there is another way we can go around that, something more convenient system to use.0682

This is called pKa; pKa is just like the pH.0687

pKa is equal to ?log of Ka.0691

You see that for large Ka values, pKa is actually small.0694

Again this is going to be true for stronger acids.0703

The nice thing about pKa is that pKa, the advantage is that we don't have to use scientific notation when we compute it.0707

It is going to be a nice small number.0716

pKa advantage avoids use of scientific notation.0718

That is the acid ionization constant Ka and pKa.0730

Let's go ahead and do a sample problem now.0734

Calculate the pH of a 1.2 molar solution of acetic acid where Ka is 1.8 times 10-5.0736

Also calculate the percent ionization.0741

Anytime you have a calculation like this, this always involves the use...0746

This is always going to involve the use of ICE tables just like we learned last time.0756

Let's go ahead and write out the equation.0766

Acetic acid CH3CO2H plus H2O liquid goes on to form CH3CO21- aqueous plus hydronium aqueous.0768

When we are reading the statement, calculate the pH of a 1.2 molar solution of acetic acid,0786

that 1.2 molar, remember the problem doesn't mention equilibrium here for 1.2 molar.0790

We assume this is initial and then zero and zero.0796

This is going to be ?x, +x, and +x.0800

This is then 0.2 minus x, x, and x.0804

We are going to set it up then to the expression.0808

Ka is equal to 1.8 times 10-5 which is equal to x squared over 1.2 minus x.0810

You remember from our last session that we talked about the assumption that x is negligible.0819

We can do this for any weak acid.0827

Assume x is negligible for weak acids and bases.0829

This equation then is approximately x squared over 1.2.0840

When we solve for x, that is going to give us the hydronium ion concentration at equilibrium which is going to be 0.0046 molar.0846

When we go ahead and solve for pH, as expected, we should get an acidic pH because this is an acid after all.0857

This is going to be pH is equal to 2.33.0864

The rules for sig figs when you do the logarithm function, the number of sig figs0867

in the concentration is equal to the number of digits after the decimal in pH.0877

As you can here for 0.0046, there is two significant figures.0894

I am going to report my pH value to two digits after the decimal.0899

That is how we quantify acid base strength.0905

Another way we can quantify acid base strength is to do the other one besides acids--that is bases.0908

Instead of using Ka, we use Kb which is the base ionization constant.0917

For Kb, we are going to say B aqueous plus H2O liquid goes on to form BH+1 aqueous and hydroxide aqueous.0922

Basically Kb is going to be equal to the concentration of BH1+ at equilibrium0936

times hydroxide at equilibrium divided by the concentration of B at equilibrium.0944

You can see it is analogous to Ka.0951

Stronger bases are going to have lots of OH- at equilibrium which means Kb is going to be large.0955

Just like pKa, pKb is going to be equal to the negative log of Kb.0969

If Kb is large, this means pKb is small.0976

Once again for stronger bases, pKb is small.0984

That is how we do the base ionization constant.0998

Finally the other ways to quantify acid base strength is what we call percent ionization.1003

For acids, this is going to be equal to the hydronium ion concentration1014

at equilibrium divided by the initial concentration of the acid times 100.1022

Then for bases, this is equal to the concentration of BH1+ at equilibrium divided by the initial concentration of base times 100.1027

Basically strong acids are going to have relatively large percent ionizations near the 99 percent.1039

For weak acids and bases, it is going to be the extreme.1046

The percent ionization is actually going to be single digits, sometimes less than1050

one percent just to show you how drastic the differences can be.1054

Now that we have gone through quantifying acid base strength,1059

let's go ahead and turn our attention to writing out acid base equilibria.1062

As you can see, in order to set up the ICE table, it is very imperative that you be1066

able to write out the correct equilibria, or else the entire problem will be messed up.1070

Let's go ahead and take some practice in writing out equilibria reactions involving acids and bases.1075

For example, HNO2 aqueous, this is going to react with water liquid.1080

Because this is weak, we use an equilibrium arrow.1088

HNO2 is going to lose a proton giving nothing but NO21- aqueous.1093

H2O is going to gain it; that becomes H3O1+ aqueous.1097

Let's go ahead and look at a base, CH3NH2 aqueous plus H2O liquid.1103

This is going to be weak base; we use an equilibrium arrow.1111

The amine is going to gain a proton, CH3NH31+ aqueous.1114

H2O loses the hydrogen giving us hydroxide.1121

A couple of notes anytime you are dealing with an amine.1127

Four amines, you always attach the H to the nitrogen atom.1132

That is note number one.1141

Note number two, you see that when an acid loses a proton, you see that the charge always decreases.1143

HNO2 becomes NO21-; for each hydrogen loss, charge decreases by one.1150

You see for the base of course, for each hydrogen gained, the charge is going to increase by one.1162

That is why we go from CH3NH20 to CH3NH31+.1172

This leads us into what we call conjugate pairs.1181

This is going to become very important conceptually for the rest of our discussion.1184

Consider the following equilibrium.1188

HF aqueous plus H2O liquid goes on to form H3O1+ aqueous and F- aqueous.1192

The only difference between HF and F- is a proton.1207

The only difference between H2O and H3O1+ is a proton.1212

When you add two molecules on opposite sides of the equilibrium that1216

differ by only one hydrogen, these are what we call conjugate pairs.1222

HF and F- are conjugates.1226

Water and hydronium are also a conjugate pair.1229

HF is the Bronsted-Lowry acid, making F- its counterpart the conjugate base.1234

The reason why it is called a conjugate base is because there is a tendency for the reverse reaction to happen1244

where F- aqueous can react with water going on to form HF aqueous and hydroxide.1252

This equation is likely to happen because HF is only a weak acid.1260

What that means is that it doesn't fully dissociate.1264

That means that F- will react with water to reform at least some of the initial HF as we can see from this reaction.1267

In general, the stronger the acid, the weaker its conjugate base.1279

Something like HCl which fully dissociates means that Cl1- will not react with water1284

to reform HCl because HCl has a great tendency to remain dissociated in water.1290

In general, the stronger an acid, the weaker its conjugate base.1297

In general, the stronger a base, the weaker its conjugate acid.1300

What we see is that acid base conjugate strengths are going to be inversely related.1304

Now let's go ahead and examine salt solutions and predict if they are acidic, basic, or neutral.1314

Basically you want to go by the following three rules to help us predict1322

if a salt is acidic, basic, or neutral just going off of its formula.1326

Salts said to produce acidic aqueous solutions contain the following combinations.1330

A conjugate acid of a weak base and a conjugate base of a strong acid--something like NH4Cl aqueous.1335

Cl1- is not going to react with water to give us anything because Cl- is the conjugate of HCl.1349

We are not going to reform HCl.1357

However NH41+ will have a tendency to react with water because it is the conjugate of only a weak base.1359

Here we are going to NH3 aqueous and hydroxide aqueous.1368

Excuse me... NH3 aqueous plus hydronium aqueous; there we go.1375

You see that we form hydronium which makes sense that we expect this solution1381

therefore to be acidic when the salt is dissolved in water.1386

Let's look at basic aqueous solutions.1393

Salts that produce aqueous solutions that are basic contain a group 1 and group 21395

metal cation like Na in combination with the conjugate base of a weak acid like F.1400

Na+ is not going to react with water to form anything.1409

It is going to remain solvated; we don't have to worry about that.1417

Also F-, F- is the conjugate of HF, a weak acid, which means there will be1422

a tendency for the following reaction to occur, formation of hydroxide and HF aqueous.1429

You see that because we form hydroxide, these types of salts are expected to be basic when dissolved in water.1437

Let's now move onto neutral salt solutions containing a group 1 and 2 metal cation and the conjugate base of a strong acid.1445

For example, NaCl, we said that Na1+, this is not going to reform hydroxide when reacting with water.1454

Cl1- is not going to react with water to reform HCl.1464

In either case, both the cation and anion do not react with water to form hydroxide or hydronium.1469

Neither is going to contribute to pH.1476

This is what we call a neutral salt solution.1479

We have so far talked about monoprotic acids, those acids that only contain one hydrogen atom.1485

But what happened for acids like sulfuric acid or phosphoric acid?1491

These are what we call polyprotic acids.1496

To calculate the pH of a 1.2 molar solution of sulfuric acid which is diprotic, we want to remember one thing.1499

It is that the first dissociation is the only one contributing to pH; first dissociation only contributes to pH.1507

Once again the first dissociation only contributes to pH1521

meaning that Ka1 is going to be relatively much greater than Ka21525

which is going to be relatively much greater than Ka3.1530

For H2SO3 aqueous plus water, we are going to write out the dissociation stepwise.1534

That is very important; write out dissociation stepwise; write out steps one at a time.1545

H2SO3 aqueous plus H2O liquid goes on to form HSO31- aqueous and H3O1+ aqueous.1558

This is what we call Ka1.1569

The next one is HSO31- aqueous is going to take its turn.1572

Plus H2O liquid goes on to form H3O1+ aqueous and SO32- aqueous.1577

For this, we are going to see Ka2.1587

We are basically saying Ka1 is much greater than Ka2.1590

To complete calculate the pH, Ka1 is going to be equal to 1.2 times 10-2.1594

You can look that up which is approximately x squared over 1.2 minus x.1604

You can go ahead and use this to solve for the hydronium ion concentration which gets us pH.1610

If you look up Ka2, Ka2 is 6.6 times 10-8.1617

That is more than six orders of magnitude less than Ka1.1622

This shows how drastic a difference the first deprotonation is versus successive deprotonations.1629

Now let's take a look at the salt of a diprotic acid.1638

Calculate the pH of a 1.2 molar solution of Na2SO3.1642

Here Na2SO3, which one of these cation or anion is going to react with water?1647

We know that Na1+ is not going to react with water.1653

There is going to be no reaction; the only thing that is left is sulfite.1659

SO32- aqueous is going to react with water.1664

We just saw that SO32- is the conjugate of a weak acid.1668

Then there is going to be a tendency for this reaction to occur where we reform HSO31- aqueous and generate hydroxide.1674

We don't use Ka for this one; we use Kb.1684

But we don't have Kb; we were only given Ka.1690

But we do know that Kw is equal to Ka times Kb.1693

This is a very important equation which is equal to 1.0 times 10-14.1697

It turns out that Kb for this reaction which we call Kb1, Kb1 represents the first protonation.1704

This is going to be equal to Kw over not Ka1 but Ka2 only.1711

The reason why we use Ka2 is because we don't generate SO32-.1717

It is not formed until both protons have been removed.1726

SO32- not formed until second deprotonation which on the previous slide if you go back is associated with Ka2.1730

Again this is how you deal with a salt solution of a polyprotic acid.1745

Now let's go ahead and take a look at acids from a different perspective.1755

So far we have talked about Bronsted-Lowry acids and bases which involve a transfer of protons.1759

However we can also look at it from a more general approach.1764

This is what we call the Lewis theory of acids and bases.1768

Lewis acids accept a lone pair of electrons to form a new covalent bond.1772

Lewis bases donate a lone pair; Lewis acids accept a lone pair.1778

You want to look for metal cations that tend to be Lewis acids.1785

You want to look for elements that do not have a complete octet,1792

that do not require a complete octet which is boron, aluminum, and beryllium.1796

These do not need a full octet.1802

They have room to accommodate an extra lone pair.1806

Lewis bases however donate a lone pair.1810

You want to look for anions for visual clues.1813

Of course you have to have lone pairs to start with; look for lone pairs in the Lewis structure.1817

How do they relate to Bronsted-Lowry acids and bases?1831

You want to remember that all Bronsted-Lowry acids/bases are also Lewis acids and bases.1834

However the reverse statement is not always true; the converse is not always true.1849

Again these are what we call Lewis acids and Lewis bases.1861

We can take a look at a very interesting example.1865

Let's look at ammonia then, NH3 aqueous.1869

I am going to put its lone pair right there.1872

It is going to react with H2O liquid going on to form NH41+ aqueous and hydroxide aqueous.1874

Let me put the lone pairs on the oxygens just to highlight the point.1885

You see that the nitrogen atom has lost a lone pair making ammonia a Lewis base.1889

You see that the oxygen atom has gained a lone pair making water the Lewis acid.1899

Just like a Bronsted-Lowry acid reacts with a Bronsted-Lowry base,1906

a Lewis acid is also going to react with a Lewis base.1910

Now that we have interpreted acids and bases differently, let's go ahead and look at factors that influence acidity strictly off of molecular structure.1916

Factor number one, the effect of charge, basically acidity is higher with increasing charge.1925

For example, H2SO4 versus HSO41-, this is going to be the stronger acid.1939

Remember we said that for polyprotic acids that successive deprotonations don't really contribute to pH.1950

Successive steps are negligible to pH; look at this.1957

The reason is the following; a hydronium cation is positively charged.1966

It is easier for something with a high charge to give up something positively charged.1971

HSO4 has its -1 charge.1975

It is going to be a relatively harder time to give up that positively charged cation.1977

Once again acidity is higher with increasing charge.1983

Now, within a period/row--CH31- aqueous, NH21- aqueous, OH1- aqueous, and HF aqueous.1988

It turns out that acidity is going to increase left to right.2003

Acidity goes up left to right in a row; the reason is the following.2008

What else increases left to right in a row?--this parallels electronegativity.2018

Remember we say that Bronsted-Lowry acids, which are also Lewis acids, they donate a proton and accept a lone pair.2028

The more electronegative the atom, the more easier it is to accept a lone pair.2039

Higher EN is easier to accept a lone pair.2045

Again acidity increases left to right in a row.2055

Within a group or column is the next one we will look at--HF, HCl, HBr, and HI.2058

It is experimentally determined that acidity increases down a column.2065

This doesn't parallel electronegativity; however this does parallel atomic size of the anion.2073

Parallels atomic size of atom directly bonded to H.2081

Again that is key; it has to be directly bonded to the H.2088

What happens is when HA become A1-, A1- is gaining a lone pair.2095

Let's look at this; it starts off with three lone pairs in HA.2103

It comes with four lone pairs in A-.2107

We are going to be better off adding electron density to a larger volume.2113

Easier to add more electrons to a larger volume.2118

In other words, I- is more stable than F- which makes HI a stronger acid2129

because the dissociation of HI is much more likely to occur and to a greater extent2144

than the dissociation of HF where F- is not as stable because of its smaller size.2151

Oxo acids, for example, HNO3 versus HNO2.2158

In oxo acids, acidity increases with the number of oxygen atoms per hydrogen.2167

To answer the question why, let's go ahead and look at the Lewis structures.2184

HNO3 here; HNO2 is going to be right there.2190

Basically what is happening is we know that HNO3 is a stronger acid.2201

It is one of the seven you have memorized; HNO2 is not.2207

What makes the difference?--the only difference is one oxygen atom.2211

That explains for why HNO3 is more acidic.2214

What happens is the oxygen atoms are highly electronegative.2217

They are going to withdraw electron density away from the bond with hydrogen.2220

That makes this partial positive.2225

That is going to make the rest of the molecule partial negative.2228

This weakens the bond with the hydrogen.2233

It is going to allow for hydrogen to leave the acid much more easily.2236

Here in HNO2, the withdrawing effect is not as great.2242

The electron density is not so lopsided.2250

It is going to be a harder time for hydrogen to fall off here with fewer oxygen atoms.2255

This is what we call an electron withdrawing inductive effect.2260

Once again this is what we call an electron withdrawing inductive effect.2267

That is oxo acids.2273

Carboxylic acids, when we go ahead and look at carboxylic acids, you also go by inductive effect.2277

When you go by the inductive effect, just look for the presence of electronegative groups.2291

For example, we can take this carboxylic acid that has two fluorines versus a carboxylic acid that has no fluorines.2298

They are structurally similar.2307

You see here that with the two fluorines, we are going to get a greater withdrawing effect.2310

Again the stronger the withdrawing effect, the more easily the hydrogen is going to quote and quote fall off.2320

H falls off more easily.2329

Here in acetic acid, there is going to be a minimum withdrawing effect from just the oxygens.2336

Once again for carboxylic acids, you want to go by the inductive effect.2356

You want to look for nearby electronegative atoms.2360

Hydrated metals cations also can be acidic.2364

If we look at Fe(H2O)63+, this is aqueous.2367

This can go ahead and react with water.2375

It can actually function as a Bronsted-Lowry acid.2379

We are going to get Fe Fe(H2O)5OH plus H3O1+ aqueous.2382

How likely is this reaction to occur?2394

This reaction is more likely to occur when this charge is very high.2396

Again for hydrated metal cations, acidity goes up with metal charge.2402

Acidity goes up with metal charge.2412

For example, the Ka of Fe(H2O)63+ is going to be greater2415

than the Ka of Fe(H2O)62+ aqueous just strictly because of charge.2426

That is molecular structure and acidity.2438

Let's go ahead and summarize the section.2441

Bronsted-Lowry acid base chemistry involves a loss or gain of a proton to or from water.2443

Conjugate pairs only differ by one proton and are inversely related in terms of acidity and basicity.2449

We learned many ways of quantifying acid base strength, namely pH, pKa, and Ka and then percent ionization.2456

Finally we saw qualitatively how the structure of a molecule can have a significant impact on how acidic it can be.2466

That is our summary of the lesson.2476

Let's now jump into a pair of sample problems.2478

Calculate the pH of a 1.2 molar solution of NH3 where Kb is 1.8 times 10-5.2481

Just like the previous lecture, the first step for any equilibrium problem is to write out the actual equilibria.2489

NH3 aqueous plus H2O liquid goes on to form NH41+ aqueous and hydroxide aqueous.2496

Let's set up the problem; 1.2 molar is given to us.2510

This is going to be 0 and 0; this is ?x, +x, and +x.2512

This thing goes to 1.2 minus x at equilibrium, x, and x.2519

Kb is 1.8 times 10-5; this is going to be approximately x squared over 1.2.2524

When we go ahead and solve for x, we get the hydroxide ion concentration at equilibrium which is going to be 0.0046 molar.2533

When we solve for pH, we had better get a pH that is basic because this is ammonia after all.2545

We get 11.66; this is sample problem one.2550

Let's now move on to sample problem two--predicting if the following salt solutions are acidic, basic, or neutral.2558

Here potassium bromide, we have a group 1 cation.2563

Br1-, this is the conjugate of HBr which is a strong acid.2569

When we have this combination, we expect this salt solution to be neutral.2576

Sodium, group 1, HPO42-.2581

This is going to be the conjugate of phosphoric acid which is going to be a weak acid.2587

We expect this compound to be basic.2593

Finally lithium cyanide, this is going to be group 1 here.2597

CN is going to be the conjugate of HCN which is considered to be a weak acid.2602

We expect this compound here to also be basic when dissolved in water.2608

That is our lesson on acid base chemistry.2617

I want to thank you for your time.2620

I will see you next time on Educator.com.2621

I. Basic Concepts & Measurement of Chemistry
  Basic Concepts of Chemistry 16:26
   Intro 0:00 
   Lesson Overview 0:07 
   Introduction 0:56 
    What is Chemistry? 0:57 
    What is Matter? 1:16 
   Solids 1:43 
    General Characteristics 1:44 
    Particulate-level Drawing of Solids 2:34 
   Liquids 3:39 
    General Characteristics of Liquids 3:40 
    Particulate-level Drawing of Liquids 3:55 
   Gases 4:23 
    General Characteristics of Gases 4:24 
    Particulate-level Drawing Gases 5:05 
   Classification of Matter 5:27 
    Classification of Matter 5:26 
   Pure Substances 5:54 
    Pure Substances 5:55 
   Mixtures 7:06 
    Definition of Mixtures 7:07 
    Homogeneous Mixtures 7:11 
    Heterogeneous Mixtures 7:52 
   Physical and Chemical Changes/Properties 8:18 
    Physical Changes Retain Chemical Composition 8:19 
    Chemical Changes Alter Chemical Composition 9:32 
   Physical and Chemical Changes/Properties, cont'd 10:55 
    Physical Properties 10:56 
    Chemical Properties 11:42 
   Sample Problem 1: Chemical & Physical Change 12:22 
   Sample Problem 2: Element, Compound, or Mixture? 13:52 
   Sample Problem 3: Classify Each of the Following Properties as chemical or Physical 15:03 
  Tools in Quantitative Chemistry 29:22
   Intro 0:00 
   Lesson Overview 0:07 
   Units of Measurement 1:23 
    The International System of Units (SI): Mass, Length, and Volume 1:39 
   Percent Error 2:17 
    Percent Error 2:18 
    Example: Calculate the Percent Error 2:56 
   Standard Deviation 3:48 
    Standard Deviation Formula 3:49 
   Standard Deviation cont'd 4:42 
    Example: Calculate Your Standard Deviation 4:43 
   Precisions vs. Accuracy 6:25 
    Precision 6:26 
    Accuracy 7:01 
   Significant Figures and Uncertainty 7:50 
    Consider the Following (2) Rulers 7:51 
    Consider the Following Graduated Cylinder 11:30 
   Identifying Significant Figures 12:43 
    The Rules of Sig Figs Overview 12:44 
    The Rules for Sig Figs: All Nonzero Digits Are Significant 13:21 
    The Rules for Sig Figs: A Zero is Significant When It is In-Between Nonzero Digits 13:28 
    The Rules for Sig Figs: A Zero is Significant When at the End of a Decimal Number 14:02 
    The Rules for Sig Figs: A Zero is not significant When Starting a Decimal Number 14:27 
   Using Sig Figs in Calculations 15:03 
    Using Sig Figs for Multiplication and Division 15:04 
    Using Sig Figs for Addition and Subtraction 15:48 
    Using Sig Figs for Mixed Operations 16:11 
   Dimensional Analysis 16:20 
    Dimensional Analysis Overview 16:21 
    General Format for Dimensional Analysis 16:39 
    Example: How Many Miles are in 17 Laps? 17:17 
    Example: How Many Grams are in 1.22 Pounds? 18:40 
   Dimensional Analysis cont'd 19:43 
    Example: How Much is Spent on Diapers in One Week? 19:44 
   Dimensional Analysis cont'd 21:03 
    SI Prefixes 21:04 
   Dimensional Analysis cont'd 22:03 
    500 mg → ? kg 22:04 
    34.1 cm → ? um 24:03 
   Summary 25:11 
   Sample Problem 1: Dimensional Analysis 26:09 
II. Atoms, Molecules, and Ions
  Atoms, Molecules, and Ions 52:18
   Intro 0:00 
   Lesson Overview 0:08 
   Introduction to Atomic Structure 1:03 
    Introduction to Atomic Structure 1:04 
    Plum Pudding Model 1:26 
   Introduction to Atomic Structure Cont'd 2:07 
    John Dalton's Atomic Theory: Number 1 2:22 
    John Dalton's Atomic Theory: Number 2 2:50 
    John Dalton's Atomic Theory: Number 3 3:07 
    John Dalton's Atomic Theory: Number 4 3:30 
    John Dalton's Atomic Theory: Number 5 3:58 
   Introduction to Atomic Structure Cont'd 5:21 
    Ernest Rutherford's Gold Foil Experiment 5:22 
   Introduction to Atomic Structure Cont'd 7:42 
    Implications of the Gold Foil Experiment 7:43 
    Relative Masses and Charges 8:18 
   Isotopes 9:02 
    Isotopes 9:03 
   Introduction to The Periodic Table 12:17 
    The Periodic Table of the Elements 12:18 
   Periodic Table, cont'd 13:56 
    Metals 13:57 
    Nonmetals 14:25 
    Semimetals 14:51 
   Periodic Table, cont'd 15:57 
    Group I: The Alkali Metals 15:58 
    Group II: The Alkali Earth Metals 16:25 
    Group VII: The Halogens 16:40 
    Group VIII: The Noble Gases 17:08 
   Ionic Compounds: Formulas, Names, Props. 17:35 
    Common Polyatomic Ions 17:36 
    Predicting Ionic Charge for Main Group Elements 18:52 
   Ionic Compounds: Formulas, Names, Props. 20:36 
    Naming Ionic Compounds: Rule 1 20:51 
    Naming Ionic Compounds: Rule 2 21:22 
    Naming Ionic Compounds: Rule 3 21:50 
    Naming Ionic Compounds: Rule 4 22:22 
   Ionic Compounds: Formulas, Names, Props. 22:50 
    Naming Ionic Compounds Example: Al₂O₃ 22:51 
    Naming Ionic Compounds Example: FeCl₃ 23:21 
    Naming Ionic Compounds Example: CuI₂ 3H₂O 24:00 
    Naming Ionic Compounds Example: Barium Phosphide 24:40 
    Naming Ionic Compounds Example: Ammonium Phosphate 25:55 
   Molecular Compounds: Formulas and Names 26:42 
    Molecular Compounds: Formulas and Names 26:43 
   The Mole 28:10 
    The Mole is 'A Chemist's Dozen' 28:11 
    It is a Central Unit, Connecting the Following Quantities 30:01 
   The Mole, cont'd 32:07 
    Atomic Masses 32:08 
    Example: How Many Moles are in 25.7 Grams of Sodium? 32:28 
    Example: How Many Atoms are in 1.2 Moles of Carbon? 33:17 
   The Mole, cont'd 34:25 
    Example: What is the Molar Mass of Carbon Dioxide? 34:26 
    Example: How Many Grams are in 1.2 Moles of Carbon Dioxide? 25:46 
   Percentage Composition 36:43 
    Example: How Many Grams of Carbon Contained in 65.1 Grams of Carbon Dioxide? 36:44 
   Empirical and Molecular Formulas 39:19 
    Empirical Formulas 39:20 
    Empirical Formula & Elemental Analysis 40:21 
   Empirical and Molecular Formulas, cont'd 41:24 
    Example: Determine Both the Empirical and Molecular Formulas - Step 1 41:25 
    Example: Determine Both the Empirical and Molecular Formulas - Step 2 43:18 
   Summary 46:22 
   Sample Problem 1: Determine the Empirical Formula of Lithium Fluoride 47:10 
   Sample Problem 2: How Many Atoms of Carbon are Present in 2.67 kg of C₆H₆? 49:21 
III. Chemical Reactions
  Chemical Reactions 43:24
   Intro 0:00 
   Lesson Overview 0:06 
   The Law of Conservation of Mass and Balancing Chemical Reactions 1:49 
    The Law of Conservation of Mass 1:50 
    Balancing Chemical Reactions 2:50 
   Balancing Chemical Reactions Cont'd 3:40 
    Balance: N₂ + H₂ → NH₃ 3:41 
    Balance: CH₄ + O₂ → CO₂ + H₂O 7:20 
   Balancing Chemical Reactions Cont'd 9:49 
    Balance: C₂H₆ + O₂ → CO₂ + H₂O 9:50 
   Intro to Chemical Equilibrium 15:32 
    When an Ionic Compound Full Dissociates 15:33 
    When an Ionic Compound Incompletely Dissociates 16:14 
    Dynamic Equilibrium 17:12 
   Electrolytes and Nonelectrolytes 18:03 
    Electrolytes 18:04 
    Strong Electrolytes and Weak Electrolytes 18:55 
    Nonelectrolytes 19:23 
   Predicting the Product(s) of an Aqueous Reaction 20:02 
    Single-replacement 20:03 
    Example: Li (s) + CuCl₂ (aq) → 2 LiCl (aq) + Cu (s) 21:03 
    Example: Cu (s) + LiCl (aq) → NR 21:23 
    Example: Zn (s) + 2HCl (aq) → ZnCl₂ (aq) + H₂ (g) 22:32 
   Predicting the Product(s) of an Aqueous Reaction 23:37 
    Double-replacement 23:38 
    Net-ionic Equation 25:29 
   Predicting the Product(s) of an Aqueous Reaction 26:12 
    Solubility Rules for Ionic Compounds 26:13 
   Predicting the Product(s) of an Aqueous Reaction 28:10 
    Neutralization Reactions 28:11 
    Example: HCl (aq) + NaOH (aq) → ? 28:37 
    Example: H₂SO₄ (aq) + KOH (aq) → ? 29:25 
   Predicting the Product(s) of an Aqueous Reaction 30:20 
    Certain Aqueous Reactions can Produce Unstable Compounds 30:21 
    Example 1 30:52 
    Example 2 32:16 
    Example 3 32:54 
   Summary 33:54 
   Sample Problem 1 34:55 
    ZnCO₃ (aq) + H₂SO₄ (aq) → ? 35:09 
    NH₄Br (aq) + Pb(C₂H₃O₂)₂ (aq) → ? 36:02 
    KNO₃ (aq) + CuCl₂ (aq) → ? 37:07 
    Li₂SO₄ (aq) + AgNO₃ (aq) → ? 37:52 
   Sample Problem 2 39:09 
    Question 1 39:10 
    Question 2 40:36 
    Question 3 41:47 
  Chemical Reactions II 55:40
   Intro 0:00 
   Lesson Overview 0:10 
   Arrhenius Definition 1:15 
    Arrhenius Acids 1:16 
    Arrhenius Bases 3:20 
   The Bronsted-Lowry Definition 4:48 
    Acids Dissolve In Water and Donate a Proton to Water: Example 1 4:49 
    Acids Dissolve In Water and Donate a Proton to Water: Example 2 6:54 
    Monoprotic Acids & Polyprotic Acids 7:58 
    Strong Acids 11:30 
    Bases Dissolve In Water and Accept a Proton From Water 12:41 
    Strong Bases 16:36 
   The Autoionization of Water 17:42 
    Amphiprotic 17:43 
    Water Reacts With Itself 18:24 
   Oxides of Metals and Nonmetals 20:08 
    Oxides of Metals and Nonmetals Overview 20:09 
    Oxides of Nonmetals: Acidic Oxides 21:23 
    Oxides of Metals: Basic Oxides 24:08 
   Oxidation-Reduction (Redox) Reactions 25:34 
    Redox Reaction Overview 25:35 
    Oxidizing and Reducing Agents 27:02 
    Redox Reaction: Transfer of Electrons 27:54 
   Oxidation-Reduction Reactions Cont'd 29:55 
    Oxidation Number Overview 29:56 
    Oxidation Number of Homonuclear Species 31:17 
    Oxidation Number of Monatomic Ions 32:58 
    Oxidation Number of Fluorine 33:27 
    Oxidation Number of Oxygen 34:00 
    Oxidation Number of Chlorine, Bromine, and Iodine 35:07 
    Oxidation Number of Hydrogen 35:30 
    Net Sum of All Oxidation Numbers In a Compound 36:21 
   Oxidation-Reduction Reactions Cont'd 38:19 
    Let's Practice Assigning Oxidation Number 38:20 
    Now Let's Apply This to a Chemical Reaction 41:07 
   Summary 44:19 
   Sample Problems 45:29 
    Sample Problem 1 45:30 
    Sample Problem 2: Determine the Oxidizing and Reducing Agents 48:48 
    Sample Problem 3: Determine the Oxidizing and Reducing Agents 50:43 
IV. Stoichiometry
  Stoichiometry I 42:10
   Intro 0:00 
   Lesson Overview 0:23 
   Mole to Mole Ratios 1:32 
    Example 1: In 1 Mole of H₂O, How Many Moles Are There of Each Element? 1:53 
    Example 2: In 2.6 Moles of Water, How Many Moles Are There of Each Element? 2:24 
   Mole to Mole Ratios Cont'd 5:13 
    Balanced Chemical Reaction 5:14 
   Mole to Mole Ratios Cont'd 7:25 
    Example 3: How Many Moles of Ammonia Can Form If you Have 3.1 Moles of H₂? 7:26 
    Example 4: How Many Moles of Hydrogen Gas Are Required to React With 6.4 Moles of Nitrogen Gas? 9:08 
   Mass to mass Conversion 11:06 
    Mass to mass Conversion 11:07 
    Example 5: How Many Grams of Ammonia Can Form If You Have 3.1 Grams of H₂? 12:37 
    Example 6: How Many Grams of Hydrogen Gas Are Required to React With 6.4 Grams of Nitrogen Gas? 15:34 
    Example 7: How Man Milligrams of Ammonia Can Form If You Have 1.2 kg of H₂? 17:29 
   Limiting Reactants, Percent Yields 20:42 
    Limiting Reactants, Percent Yields 20:43 
    Example 8: How Many Grams of Ammonia Can Form If You Have 3.1 Grams of H₂ and 3.1 Grams of N₂ 22:25 
    Percent Yield 25:30 
    Example 9: How Many Grams of The Excess Reactant Remains? 26:37 
   Summary 29:34 
   Sample Problem 1: How Many Grams of Carbon Are In 2.2 Kilograms of Carbon Dioxide? 30:47 
   Sample Problem 2: How Many Milligrams of Carbon Dioxide Can Form From 23.1 Kg of CH₄(g)? 33:06 
   Sample Problem 3: Part 1 36:10 
   Sample Problem 3: Part 2 - What Amount Of The Excess Reactant Will Remain? 40:53 
  Stoichiometry II 42:38
   Intro 0:00 
   Lesson Overview 0:10 
   Molarity 1:14 
    Solute and Solvent 1:15 
    Molarity 2:01 
   Molarity Cont'd 2:59 
    Example 1: How Many Grams of KBr are Needed to Make 350 mL of a 0.67 M KBr Solution? 3:00 
    Example 2: How Many Moles of KBr are in 350 mL of a 0.67 M KBr Solution? 5:44 
    Example 3: What Volume of a 0.67 M KBr Solution Contains 250 mg of KBr? 7:46 
   Dilutions 10:01 
    Dilution: M₁V₂=M₁V₂ 10:02 
    Example 5: Explain How to Make 250 mL of a 0.67 M KBr Solution Starting From a 1.2M Stock Solution 12:04 
   Stoichiometry and Double-Displacement Precipitation Reactions 14:41 
    Example 6: How Many grams of PbCl₂ Can Form From 250 mL of 0.32 M NaCl? 15:38 
   Stoichiometry and Double-Displacement Precipitation Reactions 18:05 
    Example 7: How Many grams of PbCl₂ Can Form When 250 mL of 0.32 M NaCl and 150 mL of 0.45 Pb(NO₃)₂ Mix? 18:06 
   Stoichiometry and Neutralization Reactions 21:01 
    Example 8: How Many Grams of NaOh are Required to Neutralize 4.5 Grams of HCl? 21:02 
   Stoichiometry and Neutralization Reactions 23:03 
    Example 9: How Many mL of 0.45 M NaOH are Required to Neutralize 250 mL of 0.89 M HCl? 23:04 
   Stoichiometry and Acid-Base Standardization 25:28 
    Introduction to Titration & Standardization 25:30 
    Acid-Base Titration 26:12 
    The Analyte & Titrant 26:24 
   The Experimental Setup 26:49 
    The Experimental Setup 26:50 
   Stoichiometry and Acid-Base Standardization 28:38 
    Example 9: Determine the Concentration of the Analyte 28:39 
   Summary 32:46 
   Sample Problem 1: Stoichiometry & Neutralization 35:24 
   Sample Problem 2: Stoichiometry 37:50 
V. Thermochemistry
  Energy & Chemical Reactions 55:28
   Intro 0:00 
   Lesson Overview 0:14 
   Introduction 1:22 
    Recall: Chemistry 1:23 
    Energy Can Be Expressed In Different Units 1:57 
   The First Law of Thermodynamics 2:43 
    Internal Energy 2:44 
   The First Law of Thermodynamics Cont'd 6:14 
    Ways to Transfer Internal Energy 6:15 
    Work Energy 8:13 
    Heat Energy 8:34 
    ∆U = q + w 8:44 
   Calculating ∆U, Q, and W 8:58 
    Changes In Both Volume and Temperature of a System 8:59 
   Calculating ∆U, Q, and W Cont'd 11:01 
    The Work Equation 11:02 
    Example 1: Calculate ∆U For The Burning Fuel 11:45 
   Calculating ∆U, Q, and W Cont'd 14:09 
    The Heat Equation 14:10 
   Calculating ∆U, Q, and W Cont'd 16:03 
    Example 2: Calculate The Final Temperature 16:04 
   Constant-Volume Calorimetry 18:05 
    Bomb Calorimeter 18:06 
    The Effect of Constant Volume On The Equation For Internal Energy 22:11 
    Example 3: Calculate ∆U 23:12 
   Constant-Pressure Conditions 26:05 
    Constant-Pressure Conditions 26:06 
   Calculating Enthalpy: Phase Changes 27:29 
    Melting, Vaporization, and Sublimation 27:30 
    Freezing, Condensation and Deposition 28:25 
    Enthalpy Values For Phase Changes 28:40 
    Example 4: How Much Energy In The Form of heat is Required to Melt 1.36 Grams of Ice? 29:40 
   Calculating Enthalpy: Heats of Reaction 31:22 
    Example 5: Calculate The Heat In kJ Associated With The Complete Reaction of 155 g NH₃ 31:23 
   Using Standard Enthalpies of Formation 33:53 
    Standard Enthalpies of Formation 33:54 
   Using Standard Enthalpies of Formation 36:12 
    Example 6: Calculate The Standard Enthalpies of Formation For The Following Reaction 36:13 
   Enthalpy From a Series of Reactions 39:58 
    Hess's Law 39:59 
   Coffee-Cup Calorimetry 42:43 
    Coffee-Cup Calorimetry 42:44 
    Example 7: Calculate ∆H° of Reaction 45:10 
   Summary 47:12 
   Sample Problem 1 48:58 
   Sample Problem 2 51:24 
VI. Quantum Theory of Atoms
  Structure of Atoms 42:33
   Intro 0:00 
   Lesson Overview 0:07 
   Introduction 1:01 
    Rutherford's Gold Foil Experiment 1:02 
   Electromagnetic Radiation 2:31 
    Radiation 2:32 
    Three Parameters: Energy, Frequency, and Wavelength 2:52 
   Electromagnetic Radiation 5:18 
    The Electromagnetic Spectrum 5:19 
   Atomic Spectroscopy and The Bohr Model 7:46 
    Wavelengths of Light 7:47 
   Atomic Spectroscopy Cont'd 9:45 
    The Bohr Model 9:46 
   Atomic Spectroscopy Cont'd 12:21 
    The Balmer Series 12:22 
    Rydberg Equation For Predicting The Wavelengths of Light 13:04 
   The Wave Nature of Matter 15:11 
    The Wave Nature of Matter 15:12 
   The Wave Nature of Matter 19:10 
    New School of Thought 19:11 
    Einstein: Energy 19:49 
    Hertz and Planck: Photoelectric Effect 20:16 
    de Broglie: Wavelength of a Moving Particle 21:14 
   Quantum Mechanics and The Atom 22:15 
    Heisenberg: Uncertainty Principle 22:16 
    Schrodinger: Wavefunctions 23:08 
   Quantum Mechanics and The Atom 24:02 
    Principle Quantum Number 24:03 
    Angular Momentum Quantum Number 25:06 
    Magnetic Quantum Number 26:27 
    Spin Quantum Number 28:42 
   The Shapes of Atomic Orbitals 29:15 
    Radial Wave Function 29:16 
    Probability Distribution Function 32:08 
   The Shapes of Atomic Orbitals 34:02 
    3-Dimensional Space of Wavefunctions 34:03 
   Summary 35:57 
   Sample Problem 1 37:07 
   Sample Problem 2 40:23 
VII. Electron Configurations and Periodicity
  Periodic Trends 38:50
   Intro 0:00 
   Lesson Overview 0:09 
   Introduction 0:36 
   Electron Configuration of Atoms 1:33 
    Electron Configuration & Atom's Electrons 1:34 
    Electron Configuration Format 1:56 
   Electron Configuration of Atoms Cont'd 3:01 
    Aufbau Principle 3:02 
   Electron Configuration of Atoms Cont'd 6:53 
    Electron Configuration Format 1: Li, O, and Cl 6:56 
    Electron Configuration Format 2: Li, O, and Cl 9:11 
   Electron Configuration of Atoms Cont'd 12:48 
    Orbital Box Diagrams 12:49 
    Pauli Exclusion Principle 13:11 
    Hund's Rule 13:36 
   Electron Configuration of Atoms Cont'd 17:35 
    Exceptions to The Aufbau Principle: Cr 17:36 
    Exceptions to The Aufbau Principle: Cu 18:15 
   Electron Configuration of Atoms Cont'd 20:22 
    Electron Configuration of Monatomic Ions: Al 20:23 
    Electron Configuration of Monatomic Ions: Al³⁺ 20:46 
    Electron Configuration of Monatomic Ions: Cl 21:57 
    Electron Configuration of Monatomic Ions: Cl¹⁻ 22:09 
   Electron Configuration Cont'd 24:31 
    Paramagnetism 24:32 
    Diamagnetism 25:00 
   Atomic Radii 26:08 
    Atomic Radii 26:09 
    In a Column of the Periodic Table 26:25 
    In a Row of the Periodic Table 26:46 
   Ionic Radii 27:30 
    Ionic Radii 27:31 
    Anions 27:42 
    Cations 27:57 
    Isoelectronic Species 28:12 
   Ionization Energy 29:00 
    Ionization Energy 29:01 
   Electron Affinity 31:37 
    Electron Affinity 31:37 
   Summary 33:43 
   Sample Problem 1: Ground State Configuration and Orbital Box Diagram 34:21 
    Fe 34:48 
    P 35:32 
   Sample Problem 2 36:38 
    Which Has The Larger Ionization Energy: Na or Li? 36:39 
    Which Has The Larger Atomic Size: O or N ? 37:23 
    Which Has The Larger Atomic Size: O²⁻ or N³⁻ ? 38:00 
VIII. Molecular Geometry & Bonding Theory
  Bonding & Molecular Structure 52:39
   Intro 0:00 
   Lesson Overview 0:08 
   Introduction 1:10 
   Types of Chemical Bonds 1:53 
    Ionic Bond 1:54 
    Molecular Bond 2:42 
   Electronegativity and Bond Polarity 3:26 
    Electronegativity (EN) 3:27 
    Periodic Trend 4:36 
   Electronegativity and Bond Polarity Cont'd 6:04 
    Bond Polarity: Polar Covalent Bond 6:05 
    Bond Polarity: Nonpolar Covalent Bond 8:53 
   Lewis Electron Dot Structure of Atoms 9:48 
    Lewis Electron Dot Structure of Atoms 9:49 
   Lewis Structures of Polyatomic Species 12:51 
    Single Bonds 12:52 
    Double Bonds 13:28 
    Nonbonding Electrons 13:59 
   Lewis Structures of Polyatomic Species Cont'd 14:45 
    Drawing Lewis Structures: Step 1 14:48 
    Drawing Lewis Structures: Step 2 15:16 
    Drawing Lewis Structures: Step 3 15:52 
    Drawing Lewis Structures: Step 4 17:31 
    Drawing Lewis Structures: Step 5 19:08 
    Drawing Lewis Structure Example: Carbonate 19:33 
   Resonance and Formal Charges (FC) 24:06 
    Resonance Structures 24:07 
    Formal Charge 25:20 
   Resonance and Formal Charges Cont'd 27:46 
    More On Formal Charge 27:47 
   Resonance and Formal Charges Cont'd 28:21 
    Good Resonance Structures 28:22 
   VSEPR Theory 31:08 
    VSEPR Theory Continue 31:09 
   VSEPR Theory Cont'd 32:53 
    VSEPR Geometries 32:54 
    Steric Number 33:04 
    Basic Geometry 33:50 
    Molecular Geometry 35:50 
   Molecular Polarity 37:51 
    Steps In Determining Molecular Polarity 37:52 
    Example 1: Polar 38:47 
    Example 2: Nonpolar 39:10 
    Example 3: Polar 39:36 
    Example 4: Polar 40:08 
   Bond Properties: Order, Length, and Energy 40:38 
    Bond Order 40:39 
    Bond Length 41:21 
    Bond Energy 41:55 
   Summary 43:09 
   Sample Problem 1 43:42 
    XeO₃ 44:03 
    I₃⁻ 47:02 
    SF₅ 49:16 
  Advanced Bonding Theories 1:11:41
   Intro 0:00 
   Lesson Overview 0:09 
   Introduction 0:38 
   Valence Bond Theory 3:07 
    Valence Bond Theory 3:08 
    spᶟ Hybridized Carbon Atom 4:19 
   Valence Bond Theory Cont'd 6:24 
    spᶟ Hybridized 6:25 
    Hybrid Orbitals For Water 7:26 
   Valence Bond Theory Cont'd (spᶟ) 11:53 
    Example 1: NH₃ 11:54 
   Valence Bond Theory Cont'd (sp²) 14:48 
    sp² Hybridization 14:49 
    Example 2: BF₃ 16:44 
   Valence Bond Theory Cont'd (sp) 22:44 
    sp Hybridization 22:46 
    Example 3: HCN 23:38 
   Valence Bond Theory Cont'd (sp³d and sp³d²) 27:36 
    Valence Bond Theory: sp³d and sp³d² 27:37 
   Molecular Orbital Theory 29:10 
    Valence Bond Theory Doesn't Always Account For a Molecule's Magnetic Behavior 29:11 
   Molecular Orbital Theory Cont'd 30:37 
    Molecular Orbital Theory 30:38 
    Wavefunctions 31:04 
    How s-orbitals Can Interact 32:23 
    Bonding Nature of p-orbitals: Head-on 35:34 
    Bonding Nature of p-orbitals: Parallel 39:04 
    Interaction Between s and p-orbital 40:45 
    Molecular Orbital Diagram For Homonuclear Diatomics: H₂ 42:21 
    Molecular Orbital Diagram For Homonuclear Diatomics: He₂ 45:23 
    Molecular Orbital Diagram For Homonuclear Diatomic: Li₂ 46:39 
    Molecular Orbital Diagram For Homonuclear Diatomic: Li₂⁺ 47:42 
    Molecular Orbital Diagram For Homonuclear Diatomic: B₂ 48:57 
    Molecular Orbital Diagram For Homonuclear Diatomic: N₂ 54:04 
    Molecular Orbital Diagram: Molecular Oxygen 55:57 
    Molecular Orbital Diagram For Heteronuclear Diatomics: Hydrochloric Acid 62:16 
   Sample Problem 1: Determine the Atomic Hybridization 67:20 
    XeO₃ 67:21 
    SF₆ 67:49 
    I₃⁻ 68:20 
   Sample Problem 2 69:04 
IX. Gases, Solids, & Liquids
  Gases 35:06
   Intro 0:00 
   Lesson Overview 0:07 
   The Kinetic Molecular Theory of Gases 1:23 
    The Kinetic Molecular Theory of Gases 1:24 
   Parameters To Characterize Gases 3:35 
    Parameters To Characterize Gases: Pressure 3:37 
    Interpreting Pressure On a Particulate Level 4:43 
   Parameters Cont'd 6:08 
    Units For Expressing Pressure: Psi, Pascal 6:19 
    Units For Expressing Pressure: mm Hg 6:42 
    Units For Expressing Pressure: atm 6:58 
    Units For Expressing Pressure: torr 7:24 
   Parameters Cont'd 8:09 
    Parameters To Characterize Gases: Volume 8:10 
    Common Units of Volume 9:00 
   Parameters Cont'd 9:11 
    Parameters To Characterize Gases: Temperature 9:12 
    Particulate Level 9:36 
    Parameters To Characterize Gases: Moles 10:24 
   The Simple Gas Laws 10:43 
    Gas Laws Are Only Valid For… 10:44 
    Charles' Law 11:24 
   The Simple Gas Laws 13:13 
    Boyle's Law 13:14 
   The Simple Gas Laws 15:28 
    Gay-Lussac's Law 15:29 
   The Simple Gas Laws 17:11 
    Avogadro's Law 17:12 
   The Ideal Gas Law 18:43 
    The Ideal Gas Law: PV = nRT 18:44 
   Applications of the Ideal Gas Law 20:12 
    Standard Temperature and Pressure for Gases 20:13 
   Applications of the Ideal Gas Law 21:43 
    Ideal Gas Law & Gas Density 21:44 
   Gas Pressures and Partial Pressures 23:18 
    Dalton's Law of Partial Pressures 23:19 
   Gas Stoichiometry 24:15 
    Stoichiometry Problems Involving Gases 24:16 
    Using The Ideal Gas Law to Get to Moles 25:16 
    Using Molar Volume to Get to Moles 25:39 
   Gas Stoichiometry Cont'd 26:03 
    Example 1: How Many Liters of O₂ at STP are Needed to Form 10.5 g of Water Vapor? 26:04 
   Summary 28:33 
   Sample Problem 1: Calculate the Molar Mass of the Gas 29:28 
   Sample Problem 2: What Mass of Ag₂O is Required to Form 3888 mL of O₂ Gas When Measured at 734 mm Hg and 25°C? 31:59 
  Intermolecular Forces & Liquids 33:47
   Intro 0:00 
   Lesson Overview 0:10 
   Introduction 0:46 
    Intermolecular Forces (IMF) 0:47 
   Intermolecular Forces of Polar Molecules 1:32 
    Ion-dipole Forces 1:33 
    Example: Salt Dissolved in Water 1:50 
    Coulomb's Law & the Force of Attraction Between Ions and/or Dipoles 3:06 
   IMF of Polar Molecules cont'd 4:36 
    Enthalpy of Solvation or Enthalpy of Hydration 4:37 
   IMF of Polar Molecules cont'd 6:01 
    Dipole-dipole Forces 6:02 
   IMF of Polar Molecules cont'd 7:22 
    Hydrogen Bonding 7:23 
    Example: Hydrogen Bonding of Water 8:06 
   IMF of Nonpolar Molecules 9:37 
    Dipole-induced Dipole Attraction 9:38 
   IMF of Nonpolar Molecules cont'd 11:34 
    Induced Dipole Attraction, London Dispersion Forces, or Vand der Waals Forces 11:35 
    Polarizability 13:46 
   IMF of Nonpolar Molecules cont'd 14:26 
    Intermolecular Forces (IMF) and Polarizability 14:31 
   Properties of Liquids 16:48 
    Standard Molar Enthalpy of Vaporization 16:49 
    Trends in Boiling Points of Representative Liquids: H₂O vs. H₂S 17:43 
   Properties of Liquids cont'd 18:36 
    Aliphatic Hydrocarbons 18:37 
    Branched Hydrocarbons 20:52 
   Properties of Liquids cont'd 22:10 
    Vapor Pressure 22:11 
    The Clausius-Clapeyron Equation 24:30 
   Properties of Liquids cont'd 25:52 
    Boiling Point 25:53 
   Properties of Liquids cont'd 27:07 
    Surface Tension 27:08 
    Viscosity 28:06 
   Summary 29:04 
   Sample Problem 1: Determine Which of the Following Liquids Will Have the Lower Vapor Pressure 30:21 
   Sample Problem 2: Determine Which of the Following Liquids Will Have the Largest Standard Molar Enthalpy of Vaporization 31:37 
  The Chemistry of Solids 25:13
   Intro 0:00 
   Lesson Overview 0:07 
   Introduction 0:46 
    General Characteristics 0:47 
    Particulate-level Drawing 1:09 
   The Basic Structure of Solids: Crystal Lattices 1:37 
    The Unit Cell Defined 1:38 
    Primitive Cubic 2:50 
   Crystal Lattices cont'd 3:58 
    Body-centered Cubic 3:59 
    Face-centered Cubic 5:02 
   Lattice Enthalpy and Trends 6:27 
    Introduction to Lattice Enthalpy 6:28 
    Equation to Calculate Lattice Enthalpy 7:21 
   Different Types of Crystalline Solids 9:35 
    Molecular Solids 9:36 
    Network Solids 10:25 
   Phase Changes Involving Solids 11:03 
    Melting & Thermodynamic Value 11:04 
    Freezing & Thermodynamic Value 11:49 
   Phase Changes cont'd 12:40 
    Sublimation & Thermodynamic Value 12:41 
    Depositions & Thermodynamic Value 13:13 
   Phase Diagrams 13:40 
    Introduction to Phase Diagrams 13:41 
    Phase Diagram of H₂O: Melting Point 14:12 
    Phase Diagram of H₂O: Normal Boiling Point 14:50 
    Phase Diagram of H₂O: Sublimation Point 15:02 
    Phase Diagram of H₂O: Point C ( Supercritical Point) 15:32 
   Phase Diagrams cont'd 16:31 
    Phase Diagram of Dry Ice 16:32 
   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 
X. Solutions, Rates of Reaction, & Equilibrium
  Solutions & Their Behavior 38:06
   Intro 0:00 
   Lesson Overview 0:10 
   Units of Concentration 1:40 
    Molarity 1:41 
    Molality 3:30 
    Weight Percent 4:26 
    ppm 5:16 
   Like Dissolves Like 6:28 
    Like Dissolves Like 6:29 
   Factors Affecting Solubility 9:35 
    The Effect of Pressure: Henry's Law 9:36 
    The Effect of Temperature on Gas Solubility 12:16 
    The Effect of Temperature on Solid Solubility 14:28 
   Colligative Properties 16:48 
    Colligative Properties 16:49 
    Changes in Vapor Pressure: Raoult's Law 17:19 
   Colligative Properties cont'd 19:53 
    Boiling Point Elevation and Freezing Point Depression 19:54 
   Colligative Properties cont'd 26:13 
    Definition of Osmosis 26:14 
    Osmotic Pressure Example 27:11 
   Summary 31:11 
   Sample Problem 1: Calculating Vapor Pressure 32:53 
   Sample Problem 2: Calculating Molality 36:29 
  Chemical Kinetics 37:45
   Intro 0:00 
   Lesson Overview 0:06 
   Introduction 1:09 
    Chemical Kinetics and the Rate of a Reaction 1:10 
    Factors Influencing Rate 1:19 
   Introduction cont'd 2:27 
    How a Reaction Progresses Through Time 2:28 
    Rate of Change Equation 6:02 
   Rate Laws 7:06 
    Definition of Rate Laws 7:07 
    General Form of Rate Laws 7:37 
   Rate Laws cont'd 11:07 
    Rate Orders With Respect to Reactant and Concentration 11:08 
   Methods of Initial Rates 13:38 
    Methods of Initial Rates 13:39 
   Integrated Rate Laws 17:57 
    Integrated Rate Laws 17:58 
    Graphically Determine the Rate Constant k 18:52 
   Reaction Mechanisms 21:05 
    Step 1: Reversible 21:18 
    Step 2: Rate-limiting Step 21:44 
    Rate Law for the Reaction 23:28 
   Reaction Rates and Temperatures 26:16 
    Reaction Rates and Temperatures 26:17 
    The Arrhenius Equation 29:06 
   Catalysis 30:31 
    Catalyst 30:32 
   Summary 32:02 
   Sample Problem 1: Calculate the Rate Constant and the Time Required for the Reaction to be Completed 32:54 
   Sample Problem 2: Calculate the Energy of Activation and the Order of the Reaction 35:24 
  Principles of Chemical Equilibrium 34:09
   Intro 0:00 
   Lesson Overview 0:08 
   Introduction 1:02 
   The Equilibrium Constant 3:08 
    The Equilibrium Constant 3:09 
   The Equilibrium Constant cont'd 5:50 
    The Equilibrium Concentration and Constant for Solutions 5:51 
    The Equilibrium Partial Pressure and Constant for Gases 7:01 
    Relationship of Kc and Kp 7:30 
   Heterogeneous Equilibria 8:23 
    Heterogeneous Equilibria 8:24 
   Manipulating K 9:57 
    First Way of Manipulating K 9:58 
    Second Way of Manipulating K 11:48 
   Manipulating K cont'd 12:31 
    Third Way of Manipulating K 12:32 
   The Reaction Quotient Q 14:42 
    The Reaction Quotient Q 14:43 
    Q > K 16:16 
    Q < K 16:30 
    Q = K 16:43 
   Le Chatlier's Principle 17:32 
    Restoring Equilibrium When It is Disturbed 17:33 
    Disturbing a Chemical System at Equilibrium 18:35 
   Problem-Solving with ICE Tables 19:05 
    Determining a Reaction's Equilibrium Constant With ICE Table 19:06 
   Problem-Solving with ICE Tables cont'd 21:03 
    Example 1: Calculate O₂(g) at Equilibrium 21:04 
   Problem-Solving with ICE Tables cont'd 22:53 
    Example 2: Calculate the Equilibrium Constant 22:54 
   Summary 25:24 
   Sample Problem 1: Calculate the Equilibrium Constant 27:59 
   Sample Problem 2: Calculate The Equilibrium Concentration 30:30 
XI. Acids & Bases Chemistry
  Acid-Base Chemistry 43:44
   Intro 0:00 
   Lesson Overview 0:06 
   Introduction 0:55 
    Bronsted-Lowry Acid & Bronsted -Lowry Base 0:56 
    Water is an Amphiprotic Molecule 2:40 
    Water Reacting With Itself 2:58 
   Introduction cont'd 4:04 
    Strong Acids 4:05 
    Strong Bases 5:18 
   Introduction cont'd 6:16 
    Weak Acids and Bases 6:17 
   Quantifying Acid-Base Strength 7:35 
    The pH Scale 7:36 
   Quantifying Acid-Base Strength cont'd 9:55 
    The Acid-ionization Constant Ka and pKa 9:56 
   Quantifying Acid-Base Strength cont'd 12:13 
    Example: Calculate the pH of a 1.2M Solution of Acetic Acid 12:14 
   Quantifying Acid-Base Strength 15:06 
    Calculating the pH of Weak Base Solutions 15:07 
   Writing Out Acid-Base Equilibria 17:45 
    Writing Out Acid-Base Equilibria 17:46 
   Writing Out Acid-Base Equilibria cont'd 19:47 
    Consider the Following Equilibrium 19:48 
    Conjugate Base and Conjugate Acid 21:18 
   Salts Solutions 22:00 
    Salts That Produce Acidic Aqueous Solutions 22:01 
    Salts That Produce Basic Aqueous Solutions 23:15 
    Neutral Salt Solutions 24:05 
   Diprotic and Polyprotic Acids 24:44 
    Example: Calculate the pH of a 1.2 M Solution of H₂SO₃ 24:43 
   Diprotic and Polyprotic Acids cont'd 27:18 
    Calculate the pH of a 1.2 M Solution of Na₂SO₃ 27:19 
   Lewis Acids and Bases 29:13 
    Lewis Acids 29:14 
    Lewis Bases 30:10 
    Example: Lewis Acids and Bases 31:04 
   Molecular Structure and Acidity 32:03 
    The Effect of Charge 32:04 
    Within a Period/Row 33:07 
   Molecular Structure and Acidity cont'd 34:17 
    Within a Group/Column 34:18 
    Oxoacids 35:58 
   Molecular Structure and Acidity cont'd 37:54 
    Carboxylic Acids 37:55 
    Hydrated Metal Cations 39:23 
   Summary 40:39 
   Sample Problem 1: Calculate the pH of a 1.2 M Solution of NH₃ 41:20 
   Sample Problem 2: Predict If The Following Slat Solutions are Acidic, Basic, or Neutral 42:37 
  Applications of Aqueous Equilibria 55:26
   Intro 0:00 
   Lesson Overview 0:07 
   Calculating pH of an Acid-Base Mixture 0:53 
    Equilibria Involving Direct Reaction With Water 0:54 
    When a Bronsted-Lowry Acid and Base React 1:12 
    After Neutralization Occurs 2:05 
   Calculating pH of an Acid-Base Mixture cont'd 2:51 
    Example: Calculating pH of an Acid-Base Mixture, Step 1 - Neutralization 2:52 
    Example: Calculating pH of an Acid-Base Mixture, Step 2 - React With H₂O 5:24 
   Buffers 7:45 
    Introduction to Buffers 7:46 
    When Acid is Added to a Buffer 8:50 
    When Base is Added to a Buffer 9:54 
   Buffers cont'd 10:41 
    Calculating the pH 10:42 
    Calculating the pH When 0.010 mol NaOH is Added to 1.0 L of the Buffer 14:03 
   Buffers cont'd 14:10 
    Calculating the pH When 0.010 mol NaOH is Added to 1.0 L of the Buffer: Step 1 -Neutralization 14:11 
    Calculating the pH When 0.010 mol NaOH is Added to 1.0 L of the Buffer: Step 2- ICE Table 15:22 
   Buffer Preparation and Capacity 16:38 
    Example: Calculating the pH of a Buffer Solution 16:42 
    Effective Buffer 18:40 
   Acid-Base Titrations 19:33 
    Acid-Base Titrations: Basic Setup 19:34 
   Acid-Base Titrations cont'd 22:12 
    Example: Calculate the pH at the Equivalence Point When 0.250 L of 0.0350 M HClO is Titrated With 1.00 M KOH 22:13 
   Acid-Base Titrations cont'd 25:38 
    Titration Curve 25:39 
   Solubility Equilibria 33:07 
    Solubility of Salts 33:08 
    Solubility Product Constant: Ksp 34:14 
   Solubility Equilibria cont'd 34:58 
    Q < Ksp 34:59 
    Q > Ksp 35:34 
   Solubility Equilibria cont'd 36:03 
    Common-ion Effect 36:04 
    Example: Calculate the Solubility of PbCl₂ in 0.55 M NaCl 36:30 
   Solubility Equilibria cont'd 39:02 
    When a Solid Salt Contains the Conjugate of a Weak Acid 39:03 
    Temperature and Solubility 40:41 
   Complexation Equilibria 41:10 
    Complex Ion 41:11 
    Complex Ion Formation Constant: Kf 42:26 
   Summary 43:35 
   Sample Problem 1: Question 44:23 
   Sample Problem 1: Part a) Calculate the pH at the Beginning of the Titration 45:48 
   Sample Problem 1: Part b) Calculate the pH at the Midpoint or Half-way Point 48:04 
   Sample Problem 1: Part c) Calculate the pH at the Equivalence Point 48:32 
   Sample Problem 1: Part d) Calculate the pH After 27.50 mL of the Acid was Added 53:00 
XII. Thermodynamics & Electrochemistry
  Entropy & Free Energy 36:13
   Intro 0:00 
   Lesson Overview 0:08 
   Introduction 0:53 
   Introduction to Entropy 1:37 
    Introduction to Entropy 1:38 
   Entropy and Heat Flow 6:31 
    Recall Thermodynamics 6:32 
    Entropy is a State Function 6:54 
    ∆S and Heat Flow 7:28 
   Entropy and Heat Flow cont'd 8:18 
    Entropy and Heat Flow: Equations 8:19 
    Endothermic Processes: ∆S > 0 8:44 
   The Second Law of Thermodynamics 10:04 
    Total ∆S = ∆S of System + ∆S of Surrounding 10:05 
    Nature Favors Processes Where The Amount of Entropy Increases 10:22 
   The Third Law of Thermodynamics 11:55 
    The Third Law of Thermodynamics & Zero Entropy 11:56 
   Problem-Solving involving Entropy 12:36 
    Endothermic Process and ∆S 12:37 
    Exothermic Process and ∆S 13:19 
   Problem-Solving cont'd 13:46 
    Change in Physical States: From Solid to Liquid to Gas 13:47 
    Change in Physical States: All Gases 15:02 
   Problem-Solving cont'd 15:56 
    Calculating the ∆S for the System, Surrounding, and Total 15:57 
    Example: Calculating the Total ∆S 16:17 
   Problem-Solving cont'd 18:36 
    Problems Involving Standard Molar Entropies of Formation 18:37 
   Introduction to Gibb's Free Energy 20:09 
    Definition of Free Energy ∆G 20:10 
    Spontaneous Process and ∆G 20:19 
   Gibb's Free Energy cont'd 22:28 
    Standard Molar Free Energies of Formation 22:29 
    The Free Energies of Formation are Zero for All Compounds in the Standard State 22:42 
   Gibb's Free Energy cont'd 23:31 
    ∆G° of the System = ∆H° of the System - T∆S° of the System 23:32 
    Predicting Spontaneous Reaction Based on the Sign of ∆G° of the System 24:24 
   Gibb's Free Energy cont'd 26:32 
    Effect of reactant and Product Concentration on the Sign of Free Energy 26:33 
    ∆G° of Reaction = -RT ln K 27:18 
   Summary 28:12 
   Sample Problem 1: Calculate ∆S° of Reaction 28:48 
   Sample Problem 2: Calculate the Temperature at Which the Reaction Becomes Spontaneous 31:18 
   Sample Problem 3: Calculate Kp 33:47 
  Electrochemistry 41:16
   Intro 0:00 
   Lesson Overview 0:08 
   Introduction 0:53 
   Redox Reactions 1:42 
    Oxidation-Reduction Reaction Overview 1:43 
   Redox Reactions cont'd 2:37 
    Which Reactant is Being Oxidized and Which is Being Reduced? 2:38 
   Redox Reactions cont'd 6:34 
    Balance Redox Reaction In Neutral Solutions 6:35 
   Redox Reactions cont'd 10:37 
    Balance Redox Reaction In Acidic and Basic Solutions: Step 1 10:38 
    Balance Redox Reaction In Acidic and Basic Solutions: Step 2 - Balance Each Half-Reaction 11:22 
   Redox Reactions cont'd 12:19 
    Balance Redox Reaction In Acidic and Basic Solutions: Step 2 - Balance Hydrogen 12:20 
   Redox Reactions cont'd 14:30 
    Balance Redox Reaction In Acidic and Basic Solutions: Step 3 14:34 
    Balance Redox Reaction In Acidic and Basic Solutions: Step 4 15:38 
   Voltaic Cells 17:01 
    Voltaic Cell or Galvanic Cell 17:02 
    Cell Notation 22:03 
   Electrochemical Potentials 25:22 
    Electrochemical Potentials 25:23 
   Electrochemical Potentials cont'd 26:07 
    Table of Standard Reduction Potentials 26:08 
   The Nernst Equation 30:41 
    The Nernst Equation 30:42 
    It Can Be Shown That At Equilibrium E =0.00 32:15 
   Gibb's Free Energy and Electrochemistry 32:46 
    Gibbs Free Energy is Relatively Small if the Potential is Relatively High 32:47 
    When E° is Very Large 33:39 
   Charge, Current and Time 33:56 
    A Battery Has Three Main Parameters 33:57 
    A Simple Equation Relates All of These Parameters 34:09 
   Summary 34:50 
   Sample Problem 1: Redox Reaction 35:26 
   Sample Problem 2: Battery 38:00 
XIII. Transition Elements & Coordination Compounds
  The Chemistry of The Transition Metals 39:03
   Intro 0:00 
   Lesson Overview 0:11 
   Coordination Compounds 1:20 
    Coordination Compounds 1:21 
   Nomenclature of Coordination Compounds 2:48 
    Rule 1 3:01 
    Rule 2 3:12 
    Rule 3 4:07 
   Nomenclature cont'd 4:58 
    Rule 4 4:59 
    Rule 5 5:13 
    Rule 6 5:35 
    Rule 7 6:19 
    Rule 8 6:46 
   Nomenclature cont'd 7:39 
    Rule 9 7:40 
    Rule 10 7:45 
    Rule 11 8:00 
    Nomenclature of Coordination Compounds: NH₄[PtCl₃NH₃] 8:11 
    Nomenclature of Coordination Compounds: [Cr(NH₃)₄(OH)₂]Br 9:31 
   Structures of Coordination Compounds 10:54 
    Coordination Number or Steric Number 10:55 
    Commonly Observed Coordination Numbers and Geometries: 4 11:14 
    Commonly Observed Coordination Numbers and Geometries: 6 12:00 
   Isomers of Coordination Compounds 13:13 
    Isomers of Coordination Compounds 13:14 
    Geometrical Isomers of CN = 6 Include: ML₄L₂' 13:30 
    Geometrical Isomers of CN = 6 Include: ML₃L₃' 15:07 
   Isomers cont'd 17:00 
    Structural Isomers Overview 17:01 
    Structural Isomers: Ionization 18:06 
    Structural Isomers: Hydrate 19:25 
    Structural Isomers: Linkage 20:11 
    Structural Isomers: Coordination Isomers 21:05 
   Electronic Structure 22:25 
    Crystal Field Theory 22:26 
    Octahedral and Tetrahedral Field 22:54 
   Electronic Structure cont'd 25:43 
    Vanadium (II) Ion in an Octahedral Field 25:44 
    Chromium(III) Ion in an Octahedral Field 26:37 
   Electronic Structure cont'd 28:47 
    Strong-Field Ligands and Weak-Field Ligands 28:48 
   Implications of Electronic Structure 30:08 
    Compare the Magnetic Properties of: [Fe(OH₂)₆]²⁺ vs. [Fe(CN)₆]⁴⁻ 30:09 
    Discussion on Color 31:57 
   Summary 34:41 
   Sample Problem 1: Name the Following Compound [Fe(OH)(OH₂)₅]Cl₂ 35:08 
   Sample Problem 1: Name the Following Compound [Co(NH₃)₃(OH₂)₃]₂(SO₄)₃ 36:24 
   Sample Problem 2: Change in Magnetic Properties 37:30 
XIV. Nuclear Chemistry
  Nuclear Chemistry 16:39
   Intro 0:00 
   Lesson Overview 0:06 
   Introduction 0:40 
    Introduction to Nuclear Reactions 0:41 
   Types of Radioactive Decay 2:10 
    Alpha Decay 2:11 
    Beta Decay 3:27 
    Gamma Decay 4:40 
    Other Types of Particles of Varying Energy 5:40 
   Nuclear Equations 6:47 
    Nuclear Equations 6:48 
   Nuclear Decay 9:28 
    Nuclear Decay and the First-Order Kinetics 9:29 
   Summary 11:31 
   Sample Problem 1: Complete the Following Nuclear Equations 12:13 
   Sample Problem 2: How Old is the Rock? 14:21