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Lecture Comments (30)

0 answers

Post by Parth Shorey on April 18 at 04:37:01 PM

How did you get 2.02 moles for chlorine ? and 2.02 moles for C?

3 answers

Last reply by: Professor Franklin Ow
Sat Apr 9, 2016 3:35 PM

Post by Parth Shorey on April 7 at 03:17:38 PM

Is this Q&A active?

0 answers

Post by Peter Ke on September 7, 2015

At 47:10, can you show me the process of converting 1mg and 2.73 mg into 1.44 x 10^-4?

0 answers

Post by Peter Ke on September 7, 2015

Why ammonium has a net charge of 1+ while phosphate has a net charge of 3-?

Also is there like a table where it shows all the different net charges in an ionic compound?

0 answers

Post by Akilah Futch on July 29, 2015

Professor, are atom/ion charges and oxidation number the same?

0 answers

Post by Kate Danielle Rapinan on December 5, 2014

should i use 1.01 or 1.008 for the mass of hydrogen? does it make a lot of difference in calculations?

(i have 2 periodic tables, one has mass as 1.01, the other has 1.00798)

0 answers

Post by Kate Danielle Rapinan on December 5, 2014

are there rare times when a metal would have a negative charge than a positive charge? If so, how would you denote/write it as a roman numeral?

2 answers

Last reply by: John-Paul Kliebert
Thu Dec 4, 2014 1:22 PM

Post by John-Paul Kliebert on November 7, 2014

With the Empirical Formula I don't understand when and when not to round. I looked online and saw that it said "If the number is too far to round (x.1 ~ x.9), then multiply each solution by the same factor to get the lowest whole number multiple." What counts as too far or not? Could you please help me?

1 answer

Last reply by: Professor Franklin Ow
Mon Nov 3, 2014 10:57 PM

Post by Hannah Duncan on November 2, 2014

How do you determine the empirical formula if you're not given any masses? Such as this question:
Predict the Empirical Formulas for the Ionic Compounds formed from the following pairs of elements. Name each compound.
a.) Li and N
b.) Ga and O
c.) Rb and Cl
d.) Ba and S

1 answer

Last reply by: Professor Franklin Ow
Tue Oct 14, 2014 6:54 PM

Post by Kirk Graham on October 13, 2014

Isn't the naming of iron backwards, the 2+ would be like this:  Fe2+ and the Roman Numeral is used when the name of the compound is written out, like this:  iron (II).  That's what our textbook in Gen Chem shows.

1 answer

Last reply by: Professor Franklin Ow
Wed Sep 24, 2014 2:44 AM

Post by Ahmed alkarkhi on September 2, 2014

Hey prof do you recommend using Classical names of the transition metals over the IUPAC names?

2 answers

Last reply by: Christine Park
Mon Dec 21, 2015 8:17 PM

Post by William Kinne on August 5, 2014

Where can I find chemistry practice problems?

1 answer

Last reply by: Professor Franklin Ow
Tue Jun 24, 2014 1:32 PM

Post by brandon joyner on June 21, 2014

I thought kg was the base unit for SI like sample problem1?

1 answer

Last reply by: Professor Franklin Ow
Sat Feb 15, 2014 4:06 AM

Post by Aniket D on February 14, 2014

Rule 2 of naming ionic compounds; stated that we should not use prefixes such as; mono, di, tri etc.to denote subscripts in a formula. However doesn't the ionic compound of Carbon dioxide use the prefix "di" to represent the 2 oxygen atoms in the formula?

1 answer

Last reply by: Professor Franklin Ow
Thu Nov 7, 2013 5:08 PM

Post by Nada A. on November 2, 2013

should t Halogens and Nobel gasses be group 17 and 18... why are they referred to as group 7 and 8?

1 answer

Last reply by: Professor Franklin Ow
Thu Nov 7, 2013 5:10 PM

Post by chisom anyanwu on October 23, 2013

is the molar mass same as the atomic mass?

Related Articles:

Atoms, Molecules, and Ions

  • Rutherford’s gold foil experiment suggested the presence of the nucleus.
  • Protons and neutrons reside inside the nucleus, while electrons are outside.
  • Isotopes are atoms of the same element with different number of neutrons.
  • The mole allows for conversion to/from number of units of anything.
  • Dalton’s law of multiple proportions is demonstrated when solving for empirical and molecular formulas.

Atoms, Molecules, and Ions

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.

  1. Intro
    • Lesson Overview
      • Introduction to Atomic Structure
      • Introduction to Atomic Structure Cont'd
      • Introduction to Atomic Structure Cont'd
      • Introduction to Atomic Structure Cont'd
      • Isotopes
      • Introduction to The Periodic Table
      • Periodic Table, cont'd
      • Periodic Table, cont'd
      • Ionic Compounds: Formulas, Names, Props.
      • Ionic Compounds: Formulas, Names, Props.
      • Ionic Compounds: Formulas, Names, Props.
      • Molecular Compounds: Formulas and Names
      • The Mole
      • The Mole, cont'd
      • The Mole, cont'd
      • Percentage Composition
      • Empirical and Molecular Formulas
      • Empirical and Molecular Formulas, cont'd
      • Summary
        • Sample Problem 1: Determine the Empirical Formula of Lithium Fluoride
          • Sample Problem 2: How Many Atoms of Carbon are Present in 2.67 kg of C₆H₆?
            • Intro 0:00
            • Lesson Overview 0:08
            • Introduction to Atomic Structure 1:03
              • Introduction to Atomic Structure
              • Plum Pudding Model
            • Introduction to Atomic Structure Cont'd 2:07
              • John Dalton's Atomic Theory: Number 1
              • John Dalton's Atomic Theory: Number 2
              • John Dalton's Atomic Theory: Number 3
              • John Dalton's Atomic Theory: Number 4
              • John Dalton's Atomic Theory: Number 5
            • Introduction to Atomic Structure Cont'd 5:21
              • Ernest Rutherford's Gold Foil Experiment
            • Introduction to Atomic Structure Cont'd 7:42
              • Implications of the Gold Foil Experiment
              • Relative Masses and Charges
            • Isotopes 9:02
              • Isotopes
            • Introduction to The Periodic Table 12:17
              • The Periodic Table of the Elements
            • Periodic Table, cont'd 13:56
              • Metals
              • Nonmetals
              • Semimetals
            • Periodic Table, cont'd 15:57
              • Group I: The Alkali Metals
              • Group II: The Alkali Earth Metals
              • Group VII: The Halogens
              • Group VIII: The Noble Gases
            • Ionic Compounds: Formulas, Names, Props. 17:35
              • Common Polyatomic Ions
              • Predicting Ionic Charge for Main Group Elements
            • Ionic Compounds: Formulas, Names, Props. 20:36
              • Naming Ionic Compounds: Rule 1
              • Naming Ionic Compounds: Rule 2
              • Naming Ionic Compounds: Rule 3
              • Naming Ionic Compounds: Rule 4
            • Ionic Compounds: Formulas, Names, Props. 22:50
              • Naming Ionic Compounds Example: Al₂O₃
              • Naming Ionic Compounds Example: FeCl₃
              • Naming Ionic Compounds Example: CuI₂ 3H₂O
              • Naming Ionic Compounds Example: Barium Phosphide
              • Naming Ionic Compounds Example: Ammonium Phosphate
            • Molecular Compounds: Formulas and Names 26:42
              • Molecular Compounds: Formulas and Names
            • The Mole 28:10
              • The Mole is 'A Chemist's Dozen'
              • It is a Central Unit, Connecting the Following Quantities
            • The Mole, cont'd 32:07
              • Atomic Masses
              • Example: How Many Moles are in 25.7 Grams of Sodium?
              • Example: How Many Atoms are in 1.2 Moles of Carbon?
            • The Mole, cont'd 34:25
              • Example: What is the Molar Mass of Carbon Dioxide?
              • Example: How Many Grams are in 1.2 Moles of Carbon Dioxide?
            • Percentage Composition 36:43
              • Example: How Many Grams of Carbon Contained in 65.1 Grams of Carbon Dioxide?
            • Empirical and Molecular Formulas 39:19
              • Empirical Formulas
              • Empirical Formula & Elemental Analysis
            • Empirical and Molecular Formulas, cont'd 41:24
              • Example: Determine Both the Empirical and Molecular Formulas - Step 1
              • Example: Determine Both the Empirical and Molecular Formulas - Step 2
            • 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

            Transcription: 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 CO32- which is called carbonate or SO42- which is called sulfate.1094

            NO31- 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-, O2-, 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