General Chemistry Atoms, Molecules, and Ions

Hi, welcome back to Educator.com.0000

Today's lesson in general chemistry is on atoms, molecules, and ions.0003

We are going to go ahead and look at the lesson overview right now.0009

We are going to get a brief introduction into atomic structure which is more of a historical background.0012

Following this, we will get into what we mean by isotopes.0019

We will get introduced to the periodic table in this lesson followed by naming compounds.0022

Basically ionic compounds, those that are made from metals and nonmetals.0028

Polyatomic ions, hydrates, and also what we mean by molecular compounds.0036

We are then going to get associated with the concept of the mole0042

and how that can be used to determine what is called a percentage composition0047

which is going to ultimately lead to the finale of this lesson0053

which is on determining what is known as the empirical and molecular formulas.0057

By the beginning of the twentieth century, physicists had already known the following.0065

Number one, that opposite charges attracted and that like charges repelled.0071

Number two, a subatomic particle existed that was negatively charged which had been called the electron; commonly abbreviated e^-.0075

One of the first accepted models of the atom was called the plum pudding model.0086

What the plum pudding model was the following.0090

It was essentially a sphere of positive charge.0092

Embedded in the sphere of positive charge were the subatomic particles, the electrons.0099

These were completely stationary if you will.0107

But there was no accounting for any other type of subatomic particle.0113

Once again, this is called the plum pudding model.0119

A sphere of positive charge that contained electrons scattered throughout.0122

We now get into one of the earliest theories that was accepted to explain for atomic structure.0129

This was called John Dalton's atomic theory.0138

John Dalton's atomic theory contained the following premises.0141

Number one was that matter was composed of individual units called atoms.0146

Number two, that these atoms were the smallest unit possible and that they were indivisible.0171

That is we could not break them down any further; atoms were indivisible.0177

Number three, the third premise of John Dalton's atomic theory was that atoms of one element were all identical to each other.0188

Number four, atoms of one element are going to be different than atoms of another element.0213

That is what we reference as chemical identity.0218

Atoms of one element different than atoms of another element.0224

Finally the fifth and final premise was that when elements combine, they can combine in small whole number ratios to form compounds.0239

Atoms combine in small whole number ratios to form compounds.0250

This fifth premise here is going to be looked at again.0268

That is what we call the law of multiple proportions.0273

I am going to go ahead and put an asterisk next to0284

each of the premises which we actually find not be so true today.0286

Atoms were indivisible; that is not true because we can divide an atom in a process called nuclear fission.0292

So premise two is not completely accurate.0298

The next one is atoms of one element are all identical to each other.0302

We are going to talk about this--is that not all atoms of the same element are identical.0307

Because atoms can exist in different forms which we call isotopes.0312

That is going to be coming on later on.0317

Following John Dalton's atomic theory was one of the first major experiments on atomic structure.0323

This is what we call Ernest Rutherford's gold foil experiment.0331

What Ernest Rutherford did was the following.0334

He shot a beam of high energy particles that were positively charged.0337

These positively charged particles were what we call alpha rays.0343

This was shot at a very thin piece of gold foil.0348

What was the expected result?--the expected result was the following.0355

That all of the alpha rays should have gone straight through.0360

Why?--because the positive sphere was mostly empty space.0366

And that the positive ray should have been attracted to the electrons that are embedded in the sphere.0370

This is expected result due to plum pudding.0377

However the actual result was the following.0387

From the gold foil, yes the majority of the rays did indeed go straight through.0390

However what was determined was that some rays actually got deflected at various angles.0397

You would actually have a couple rays get deflected completely backward.0405

This was highly unusual because this was not predicted by the plum pudding model.0411

Instead what came about was the following implication; an implication was the following.0416

That matter and atoms contained a dense center of positive charge.0424

Because it is positive charge that was making the alpha rays deflect.0441

This positive charge is what we today refer to as the nucleus.0449

The nucleus is at the center of all atoms and is positively charged.0456

To resummarize the implications of the gold foil experiment was that all atoms0464

have a dense center that is positively charged which we call the nucleus.0469

Inside the nucleus, two subatomic particles exist.0474

The proton is positively charged; the neutron has zero charge.0477

Number three, the nucleus is responsible for the mass of the atom; that is what gives an atom weight.0482

Finally electrons, which were already known to have existed, exist outside the nucleus and have negligible weight.0489

Let's go ahead and summarize these subatomic particles--neutron, proton, and electron.0499

Please make a note that these relative masses and the relative charges are indeed in relation to each other.0510

A neutron is going to be assigned a mass of 1 and a charge of 0.0520

A proton is assigned a mass of 1 and a charge of +1.0524

An electron has a negligible mass, a mass of 0, and a charge of -1.0529

Of course, we symbolize the neutron lower-case n, proton p, and the electron is going to be e.0535

One fallacy in John Dalton's atomic theory was that atoms of the same element are all identical to each other.0547

But we are going to learn that atoms of the same element can exist in different forms; these are called isotopes.0555

Isotopes are atoms of the same element; they differ only in their number of neutrons.0563

Number of protons is going to be equal to the number of electrons; this is because of charged neutrality.0570

But the only difference is the number of neutrons; number of n differs.0577

We are going to introduce something called isotope format.0586

For example, carbon has two isotopes, carbon-12 and carbon-13.0589

Anytime you see the number immediately following the hyphen, this is what we call the mass number of the isotope.0597

The mass number of the isotope is simply equal to the number of neutrons plus the number of protons.0607

When we write an isotope in a specific format, it is always going to have the following design.0618

Element symbol, X; A on top and Z on the bottom as a subscript.0628

X is simply your element symbol; A is your mass number; Z is simply your atomic number.0634

The atomic number is strictly equal to the number of protons of the element.0648

A mathematical relation follows that if you take the difference of A minus Z, you are going to get the following.0659

The number of neutrons plus the number of protons and then minus the number of protons; as you can see protons cancels.0668

When you take the difference of A minus Z, you simply get the number of neutrons.0677

For example, if you take 12, 6 and carbon.0683

We have 12 is equal to the number of neutrons plus the number of protons.0689

6 is on the bottom; that is the number of protons which is also equal to the number of electrons.0698

Therefore A minus Z is equal to 6; that is equal to the number of neutrons.0704

So carbon-12 has six protons, six neutrons, and six electrons.0712

However if you do the same analysis for carbon-13, we are going to get six protons, seven neutrons, and six electrons.0721

As you can see, isotopes strictly differ by the number of neutrons present inside the nucleus.0730

Now we get introduced to the periodic table.0739

The periodic table is essentially organized into different rows and columns.0742

Each column is also referred to as a group; each row is referred to as a period.0749

Typically there is always a staircase that is embedded in all of our periodic tables.0766

Anything to the left side of the staircase is what we call a metal.0776

Anything to the right side of the staircase is what we call a nonmetal.0785

When we look at the periodic table, you see always the following.0797

You see a number on top; the number of top is what we call the element number.0802

The element number is actually the atomic number, the number of protons.0810

The number on the bottom is the atomic mass of the element.0816

Finally right in the middle is your element symbol.0823

We are going to talk now more about different areas of the periodic table.0830

We talked about metals; metals have the following characteristics.0839

They appear lustrous which is shiny; they are good conductors of heat and of electricity.0843

They are malleable which means we can form dinged sheets out of them.0850

And they are ductile which means we can pull them into fine wires.0856

Nonmetals are going to be the opposite.0866

They are going to be electrically and thermally insulated which is what we mean by poor conductivity.0869

In addition, they are brittle which means that if we take a hammer and hit them, they are going to crumble.0878

Instead are not too malleable; so will crumble.0885

But right in the middle of the border between metals and nonmetals, that is what we find semimetals.0891

Semimetals are also known as metalloids.0900

They border that staircase that we drew previously; border the staircase.0904

Because they are called semimetals or metalloids, they are going to have properties of both.0915

They tend to be good conductors of electricity but poor conductors of heat; they tend to be malleable.0923

The typical examples are of course silicon and germanium.0928

We see these two being used in a lot of electrical devices and what we call semiconductors.0938

What we are going to analyze next are some of the major columns of the periodic table.0960

Group 1 is what we call the alkali metals such as lithium and sodium and potassium, etc.0966

Group 1 metals tend to have the following characteristics.0974

That is they are moisture sensitive which means they have a tendency to react with water.0977

They can be explosive when they react with water.0982

Group 2 are what we call the alkaline earth metals such as calcium and magnesium and barium.0986

These are harder than group 1 and less moisture sensitive.0995

Group 6 are the halogens which are fluorine, chlorine, bromine, and iodine.1001

All of these occur naturally in nature diatomically which means they occur in pairs.1007

When someone says fluorine, it is really F₂, Cl₂, Br₂, and I₂; the halogens are incredibly reactive.1015

Finally group number 8, the rightmost column of the periodic table are what we call the noble gases such as helium, argon, and neon.1028

Unlike the group 7s, these are going to be monoatomic; that is they occur just as individual units.1039

They are called noble gas species because they are relatively inert and have a smaller tendency to react.1045

Now that we have been introduced to the periodic tables and the elements that comprise them, let's now talk about different types of compounds.1057

The first type of compound we are going to study are called ionic compounds.1066

Ionic compounds are formed between a metal and a nonmetal or they contain what are known as polyatomic ions.1069

Because of the strong electrostatic attractions that exist between opposite charges, ionic compounds tend to have very high melting points.1078

Below are a couple of these polyatomic ions; polyatomic could be the following.1087

Something like CO₃^2- which is called carbonate or SO₄^2- which is called sulfate.1094

NO₃^1- which is called nitrate; you see why these are called polyatomic ions.1107

Because it is several atoms coming together to form one entity; those are what we call polyatomic.1111

That is different than monoatomic ions of course which are just formed from one unit.1120

Something like Cl^-, O^2-, Al³⁺, etc.; polyatomic versus monoatomic.1124

It turns out that we can actually use the periodic table to predict the ionic charge for main group elements.1135

Main group elements are going to be group 1 and group 2, then group 3 to 6, and then 7 and 8.1141

It turns out that all of the group 1's tend to form a 1+ charge.1160

All of the group 2's tend to form a 2+ charge; group 3's tend to form a 3+ charge.1167

We are going to skip 4 and 5 because they tend not to form ionic compounds to such a large degree.1174

Column 6, we are going to get a 2- charge; column 7 is going to be a 1- charge.1183

Group 8, we say no charge because group 8 are the noble gases and they tend not to react with other species.1188

You noticed also very importantly that the positive charges, which are what are called cations, these are going to be formed by metals.1196

The negative charges are what are called anions; these are all nonmetals.1208

I am sorry; I am leaving out group 5.1219

Group 5 is of course... this is going to be a 3- charge.1222

So you have your cations and then your anions.1226

Once again cations are going to be formed by metals and anions are going to formed by nonmetals.1229

Now we have to learn how to name ionic compounds, such as NaCl, such as BaF₂.1237

We need to now go over the systematic procedure for doing so.1247

You are going to name the cation first followed by the anion and the anion will end in ?ide.1253

Something like NaCl; the cation is simply going to be sodium.1259

Cl is usually chlorine; but now we are going to become chloride.1267

So the correct name of NaCl is simply sodium chloride; very simple.1275

You do not use prefixes mono, di, tri, etc, to denote subscripts in a formula.1284

Something like for example BaF₂; let's go ahead and name it.1289

That is going to be barium followed by fluoride; you do not say difluoride.1294

No, no difluoride, for example; again we ignore the prefixes.1302

For certain transition metals, you are going to use a Roman numeral in parentheses to indicate the metal charge.1310

If we want Fe²⁺, we are going to denote it as iron(II); Fe³⁺ cation is going to be Fe(III).1317

Again you should definitely ask your instructor as to which transition metals he or she1331

would like you to know for sure where you have to use a parentheses to indicate charge.1336

Finally you can also have a hydrate; a hydrate is when you have water attached.1343

For example, CoCl₂·5H₂O, we are going to use the prefixes followed by the word hydrate.1350

Five using the Greek prefix would be penta; we would just say that this is the pentahydrate compound of CoCl₂.1360

Let's go ahead and put everything together and do a couple of examples.1372

Al₂O₃, FeCl₃, CuI₂·3H₂O; we are going to name the cation first.1378

Al₂O₃, that is going to be aluminum followed by oxygen becoming oxide.1394

FeCl₃, Fe is one of those we need to know.1402

Iron, something in parentheses, and then chlorine becomes chloride.1406

Now we have to come up with the charge; each chlorine is a 1- charge.1411

There is three of them, giving me a 3- overall.1416

For iron, iron is going to be 3+ here, giving us iron(III) chloride in the end.1422

Don't forget for the ionic compounds, the net charge should equal zero to maintain charge neutrality.1431

Right here we are going to have copper.1442

Copper is one of those where you need a Roman numeral for, followed by iodide.1444

Then we have three waters which becomes trihydrate.1451

Here iodine is a 1- charge; there is two of them giving me a -2 overall.1457

Copper must be a +2 to balance it; I am going to indicate that in the parentheses as Roman numeral.1464

Now we have to do the other way.1474

Before I gave you the formula and you were asked to provide the name.1476

Now let's go ahead and write the formula given the name.1480

Let's go ahead and look at the following examples; barium phosphide; ammonium phosphate.1485

Barium phosphide; barium is Ba; phosphide is phosphorus which is P.1511

What I would like to do is I would like to put the charges up in the top just for my own purposes.1518

Barium is 2+; phosphorus is 3-; what I always tell my students to do is then I use the crossing rule.1524

Basically if you cross the charges, they become the subscripts for each of the other elements.1530

Ba is going to get subscript of 3; phosphorus is going to get a subscript of 2.1537

You can convince yourself to make sure that the charges cancel.1543

Three of the bariums times 2+ plus two of the phosphorus times 3-, yes they do cancel to zero.1547

Next one is ammonium phosphate; ammonium is NH₄; phosphate is PO₄.1556

Ammonium is a 1+ charge; phosphate is a 3- charge; let's go ahead and cross them again.1564

When we cross them, we are going to get (NH₄)₃PO₄ with a 1.1572

Let's go ahead and make sure that they all work out.1582

Three of the ammoniums times a +1 charge plus one of the phosphates and a 3- charge; yes it does cancel to zero.1585

That is some nice sample problems on naming ionic compounds.1596

The other type of compound we are going to learn about now is called a molecular compound.1603

Molecular compounds are not a metal and nonmetal but instead they are composed from two nonmetals.1608

Something like CO₂; something like CO; and of course H₂O itself.1613

To name a molecular compound, the rules are going to be very similar with one main exception.1620

We are going to name the first element first by its entire name; then the second element has a ?ide suffix.1627

This time we do pay attention to the subscripts.1636

We are going to use the Greek prefixes, mono, di, tri, tetra, etc, to determine what the subscript is.1639

However you omit mono for the first nonmetal.1647

Something like CO₂, name the first element first which is carbon.1651

The second element is oxygen becoming oxide.1656

There is two of them which is how we come up with carbon dioxide.1660

Something like N₂O₅; there is two nitrogens which is di.1665

The first element gets the full name; so dinitrogen; then five oxides or just pentoxide.1670

So actually naming a molecular compound is always easier than naming an ionic compound.1679

Why?--because you don't have to worry about charge.1685

Now onto what we call the concept of the mole.1691

When you go to the supermarket and you purchase eggs, they are always in either a half dozen or a dozen.1697

It is very hard and very seldom that you are going to purchase just an individual egg.1704

It is convenience; it is convenience for us to refer to eggs as a dozen.1708

For chemistry, we have the equivalent of our dozen.1716

The mole is our chemists' dozen for convenience.1721

When you go into lab, you don't measure just the weight of for example one atom.1730

You don't deal with an individual atom; it is not practical.1735

What the mole is all about is that it goes from the microscale to a macroscale; it is more practical.1739

Basically our dozen which is what we call the mole... the mole is abbreviated simply m-o-l.1755

It is equal to the following: 6.022 times 10²³ units of whatever you are measuring.1763

This is equivalent of saying one dozen is equal to twelve units of whatever you are measuring.1774

Again our dozen, the chemists' dozen is what we call the mole.1782

This mole is probably... this number you have to memorize I am pretty sure.1786

6.022 times 10²³ is what we call Avogadro's number; Avogadro's number.1792

The mole we are going to find is a central unit; it can connect the following quantities.1803

The mole, we can equal and calculate individual units; the mole we can also get grams.1809

When we go from mole to units, we are going to multiply this way.1823

By Avogadro's number, 6.022 times 10²³ units for every one mole.1830

But when we go the opposite direction, we are going to divide.1838

We are going to divide by Avogadro's number.1842

Mole can also be used to go to mass.1854

When we go to mass, we can actually multiply the mole by a rate conversion factor of gram for every one mole.1858

When we go the opposite direction, we are going to divide by one mole.1871

We are going to divide by grams; so we can see that mole is truly a central unit.1879

The conversion factor of gram per mole, that is what we call the molar mass.1892

This is equal to the grams that one mole of a substance weighs.1903

We are going to come back to this quite heavily.1920

We are going to do a lot of example problems utilizing the mole.1924

Where do we get this molar mass from then?--on the periodic table.1929

On the periodic table, atomic masses... remember it is the number below the element symbol... are also called molar masses.1935

When you look at C-6, 12.101 on the periodic table, this bottom number here it the molar mass of carbon.1943

One mole of carbon weighs exactly 12.01 grams; that is how you write it out.1953

That is going to be a conversion factor.1960

How many moles are in 25.7 grams of sodium?1962

You say 25.7 grams of sodium times something over something.1967

That is going to give us moles of sodium.1973

Remembering our dimensional analysis, now grams of sodium is going to go downstairs.1978

Moles of sodium is going to go upstairs.1984

You just look up the value; that is going to be 22.99 grams of sodium for every one mole of sodium.1988

Giving you your answer in units of moles of sodium.1994

Let's now go ahead and utilize Avogadro's number.1998

How many atoms are in 1.2 moles of carbon?2002

1.2 moles of carbon times something over something is going to give us our answer in units of atoms of carbon.2008

To get cancelled, moles of carbons goes downstairs; then atoms of carbon goes upstairs.2021

We know from our flow chart with the mole that this is involving Avogadro's number.2029

One mole will go downstairs and 6.022 times 10²³ goes upstairs to give our answer in units of atoms of carbon.2034

Whenever you see the word atoms or molecules, that is pretty much a good give away2046

that you are going to be using Avogadro's number as your conversion factor.2053

Again always keep in mind that; especially if the word moles is mentioned too.2060

Now that we have been introduced to what the mole can help us get,2068

we are going to find something else that the mole can be used to calculate,2073

what is called molar mass of not only an atom but also of a compound.2079

Let me say that again; the mole can be used to calculate the molar mass of a compound.2084

The molar mass of an element is already provided to us on the periodic table.2088

Basically the molar mass is simply the net sum of all molar masses of each element.2093

The molar mass of carbon dioxide is simply equal to2111

the molar mass of carbon plus the molar mass of oxygen2114

plus the molar mass of oxygen because we have two oxygens.2122

That is going to be equal to 12.01 grams plus 16.00 grams plus 16.00 grams.2126

Giving us a grand total of 44.01 grams; one mole of CO₂ weighs exactly 44.01 grams.2135

Having said that, we can also use molar mass as a conversion factor for a compound.2149

How many grams are in 1.2 moles of carbon dioxide?2154

1.2 moles of CO₂ times something over something is going to give us our answer in units of grams of CO₂.2158

The mole is going to go downstairs; the grams is going to go upstairs.2169

Right there in the previous problem, we calculated this conversion factor or molar mass of CO₂.2175

Anytime you want to relate mass and moles, it is always molar mass to be your conversion factor.2183

It is going to be 1 on the bottom; it is going to be 44.01 on top.2190

Giving us our answer in units of grams of CO₂.2195

The next application of the mole is what we call percentage composition.2206

Percentage composition is given to us to be the following.2211

Percent composition of an element tells us basically the parts of the element divided by total parts times 100.2216

It basically tells us your relative amount that a specific element makes up of the compound.2236

This equation that we actually use is going to then be the total mass of the element2244

divided by the molar mass of the compound; all of that times 100.2256

Let's go ahead and answer this question.2269

How many grams of carbon dioxide are contained in 65.1 grams of CO₂?2270

This problem is strictly wanting to know how much carbon is going to be contained in so much CO₂.2276

Let's go ahead and first calculate the percentage composition of carbon in CO₂.2283

We are going to say 12.01 grams because there is only one carbon.2294

Divided by the entire molar mass of CO₂ which is 44.01 grams; times 100.2299

That is going to give us 27.3 percent.2306

Carbon dioxide is only 27.3 percent carbon; the majority is oxygen.2309

That tells me that any sample I hold of CO₂ in my hand, 27.3 percent of that is carbon.2317

We are going to use percentage as our conversion factor.2324

We are going to say .273 times the 65.1 grams of CO₂ is going to give me my grams of carbon.2328

We are going to get 17.8 grams as our answer.2339

Once again anytime you see a problem that relates the amount of an element2345

contained in a certain amount of compound, you want to use percentage composition as your tool.2351

Finally we are now going to see how we can determine what is known as an empirical and molecular formula.2360

An empirical formula is the simplest whole number ratio of a formula.2367

For example, CO₂, this is an exact 1:2 ratio of carbon to oxygen.2371

That is my smallest ratio; I cannot go smaller than that.2378

However let's go ahead and look at one that is maybe C₂H₆.2382

In this example, I have a ratio of 2:6 of carbon to hydrogen.2388

But you notice that I can divide; I can divide by a lowest common multiple which is 2.2395

I am going to get 1:3 ratio instead; therefore 1:3 is actually my smallest ratio.2402

That is going to be CH₃ which is my empirical formula.2410

What we are going to do is we are going to determine how to calculate the empirical formula.2415

The empirical formula is determined from elemental analysis.2422

Basically you have an unknown sample and you put it through...2428

You perform what is called an elemental analysis test on it.2434

After you perform this analysis, you are going to get a printout2440

of the percentage that a certain element makes up in your unknown.2445

You are going to get x percent of for example carbon.2450

You are going to get y percent of for example oxygen.2454

You are going to get z percent for example of hydrogen.2457

These percentages again are only to be used for empirical formulas.2461

They may or may not necessarily give you the actual formula; only the relative amounts are reported.2466

If you want to get the actual formula, which is known as the molecular formula, you can determine it.2473

But only if the molar mass is already provided for you.2478

Let's go ahead and look at a typical example.2486

A compound used as an additive for gasoline to help prevent engine knock has the following percent composition by mass.2490

71.65 percent chlorine; 24.27 percent carbon; 4.07 percent hydrogen.2497

The molar mass is known to be 98.96 grams per mole.2505

Determine both the empirical and molecular formulas.2509

The very first step is to get your percentages into grams; get percentage to grams.2516

After we get everything into grams, then we can go to our central unit from the chapter which is moles.2527

Again whenever you are in doubt, get to moles.2533

The nice thing about percentage is that percentage is always made out of 100.2536

If we assume 100 gram of compound, the percentages automatically become grams.2541

We have 71.65 grams of chlorine, 24.27 grams of carbon, and 4.07 of hydrogen.2548

Again assuming 100 gram of compound.2560

Now we are going to go ahead and get to moles.2572

This is something we know how to do by now.2575

When we go ahead and get to moles, we are going to get 2.02 moles of chlorine.2577

We are going to get 2.02 moles of carbon.2584

We are going to get 4.04 moles of hydrogen.2587

Step one is done; now step two.2596

Step two, we are going to take our moles and we are going to divide.2601

We are going to divide by the smallest moles present; divide by smallest moles present; everything.2604

We are going to get a nice whole number that is going to become our subscripts in our empirical formula.2617

The result is the empirical formula subscript.2626

When we take chlorine and divide it by the smallest number which is 2.02, we are going to get 1.2642

Carbon divided by 2.02, get 1.2648

For hydrogen, 4.04 divided by 2.02, we are going to get 2.2651

Giving us our empirical formula of CH₂Cl.2655

What happens if you don't get a perfect whole number?--what if you get 1.5, 1.5, 2?2665

Dalton's law of multiple proportions tells us that we cannot combine elements in other than a small whole number ratio.2684

Basically any time you have a decimal number, we are just going to then2694

multiply everything by the same factor to get everything in whole numbers.2697

When we do that, we can multiply 1.5 by 2 and everything by 2.2702

Giving us nothing but whole numbers; 3, 3, and 4.2707

If that happened, then those numbers would become the empirical; we would get C₃H₄Cl₃.2711

To get the molecular formula, you simply do the following.2722

You take the molar mass of the molecular formula which is given to you.2726

You divide it by the molar mass of your empirical formula.2733

That is always going to give you a nice whole number.2738

When we do this, we are going to get 98.96 divided by our molar mass of our empirical formula which is CH₂Cl.2741

We actually get 48.468; when we do this, we get approximately 2.2752

You take that whole number and you multiply the subscripts of the empirical formula by this.2758

I get a molecular formula of C₂H₄Cl₂ which is the formula of the actual compound.2764

You can always double check your work.2773

You can add up the molar mass of C₂H₄Cl₂.2774

We are going to get very close to 98.96.2778

Let's go ahead and summarize.2783

Atoms are composed of a central nucleus that is positively charged.2786

Protons and neutrons reside within the nucleus; electrons are outside.2790

We saw that isotopes are the same element; differ only by their number of neutrons.2796

We have went through some specific rules for naming ionic and molecular compounds.2801

We also saw that the mole is a central unit that allows for conversion between number of atoms and molecules and for mass.2806

Finally the mole is a central unit that is required for several types of different problems2814

Including percentage composition and the empirical and molecular formula problems.2822

Let's go ahead and look at some sample problems.2831

You have 1 milligram of lithium metal reacting with molecular fluorine gas.2834

The resulting fluoride salt has a mass of 7.3 mg.2838

Determine the empirical formula of lithium fluoride; you have 1.00 mg of lithium.2842

After it reacts with the molecular fluorine gas which is F₂, you get 3.73 mg of your lithium fluoride.2850

If I start with only 1 mg of lithium and I wind up with 3.73 mg of compound, why did my mass get heavier?2865

The mass got heavier because it reacted with fluorine; fluorine came on board.2874

We can actually determine the mass of fluorine.2879

It is just the difference, 3.73 mg minus 1.00 mg.2882

That is going to give us 2.73 mg which is the mass of fluorine reacted.2888

We got all of our masses now; now let's get everything into moles.2902

2.73 mg of fluorine is going to become 1.44 times 10^-4 moles of fluorine.2908

1.00 mg of lithium is going to become 1.44 times 10^-4 moles of lithium.2919

Now that we have all of our moles, we can divide by the smallest number present.2930

For this problem, they are identical.2934

When we divide everything, we are going to get Li₁ and F₁ for our empirical formula; or just LiF.2937

This answer actually makes sense because we know that lithium is a 1+ charge and that fluorine is a 1- charge.2947

Indeed the charges do balance each other out.2954

Now moving onto sample problem two.2962

How many atoms of carbon are present in 2.67 kg of C₆H₆?2964

In this case, we are asked about an element within a certain compound.2970

That sounds a lot like percentage composition.2976

Let's go ahead and calculate the percentage composition of carbon in C₆H₆.2978

The percentage composition is equal to the total mass of carbon...2984

There is six of them by the molar mass 12.01.2988

Divided by the total mass of C₆H₆ which is 6 by 12.01 plus the 6 hydrogens by 1.008.2991

All of that times 100; when we do this, we get 92.3 percent.3001

So C₆H₆ is mostly carbon; 92 percent carbon.3007

Remember what we did last time.3013

We are going to take our percentage composition 0.923.3015

We are going to multiply it by the total mass 2.67 kg.3019

That is going to give us the 2.46 kg of carbon in this specific sample.3023

The question is asking for atoms of carbon though.3031

Somehow we have to go from kilograms to atoms.3035

Remember anytime you see the word atoms or molecules, it is going to be via an Avogadro number.3040

Our first step is to go from kg just to regular g.3047

Then from g to the central unit which is moles.3052

Then of course moles onto atoms; we can go ahead and do this.3056

We are going to say 2.46 kg of carbon times something over something.3063

G goes on top; kg on the bottom; that is going to be 10³ kg over 1 kg.3071

Now onto moles; times something over something; g goes downstairs to get cancelled.3078

Moles goes upstairs to get carried through the final answer; that is molar mass.3084

You look up the molar mass of carbon which is 12.01 grams for every one mole.3090

Then finally times something over something, giving us our answer in atoms of carbon.3095

That is Avogadro's number; mole on the bottom; Avogadro's on top which is atoms here.3102

That is 1 mole on the bottom and 6.022 times 10²³ atoms on top.3110

You should get an answer of 1.23 times 10²⁶ atoms of carbon.3116

I want to thank you for your attention.3128

This concludes our lecture on atoms, molecules, and ions.3130

Thank you for using Educator.com.3135