Chapter 2: Atoms, Molecules,

and Ions

1. Atomic Structure

2. Atomic Number & Mass

3. Isotopes

4. Atomic Weight

5. Periodic Table

7. Molecular Formulas & Nomenclature

8. Molecular Compounds

9. Atoms, Molecules & Mole

10. Compound Formulas

11. Hydrated Compounds

2.1 Structure of The Atom

Modern View

Particle Relative Mass Relative Charge

(g) (amu)

Proton 1.6726 x 10-24 1.0073 +1

Neutron 1.6749 x 10-24 1.0087 0

Electron 9.1094x10-28 0.0005486 -1

Structure of The Atom

Modern View

Three Fundamental Particles

1. Proton (Positive Charged Particle)

Nucleus - Region of Atomic Mass Density,

Contains Neutrons and Protons

2. Neutron (Neutral Particles)

3. Electron - Negative Charged Essentially

Mass Less Particle, Occupies Majority of

Atomic Space

Protons, Neutrons & Electrons:

Historical Perspective

Ben Franklin (1706-1790)

Labeled 2 types of charge as

Positive and Negative

Like Charges Repel

Opposite Charges Attract

Noted that charge is balanced, so for

every [+] charge there exists a [-] charge

Electricity

Types of Radioactive Processes

1. Alpha Decay

-Heavy [+] Particle

2. Beta Decay

-Light [-] Particle

3. Gamma Decay

-Charge less

-Release of Energy

Some History on The Discovery of the Atom

Discovery of The Electron:

J.J. Thomson and the Cathode Ray Tube

(1897)

Calculated

e-/m

Some History on The Discovery of the Atom

Discovery of The Electron:

Robert Millikan (1911) Measured the

Charge of the Electron

Charge of electron

= 1.6 x 10-19C.

Some History on The Discovery of the Atom

Devised an Atomic Theory From the Following

Postulates:

1. Atoms Are Neutral and Have Negative

Tiny Electrons.

2. Atoms Must Therefor Have Positive Particles to

Balance the Charge

3. Electrons Can Not Account for the Mass of the Atom

Plumb-Pudding Model, Small Electrons Are in the

Middle of a Large Sphere of Positive Charge

Some History on The Discovery of the AtomRutherford’s Gold Foil Experiment and

the Discovery of the Nucleus

Believed in Plum Pudding Model

Wanted to see How Large Atoms are

Atomic Number & Mass Number

The Mass of an Atom Can Be Determined From the

Masses of the Neutrons and Protons in It’s Nucleus

This Is a Very Small Number

1 Amu = Mass of 1 Carbon Atom With

6 Neutrons & 6 Protons

Atomic Mass Unit (AMU)

1amu = 1.66 X 10-24g

Atomic Number

# of Protons in an Atom

Know How to Determine Atomic

Numbers & Masses From the

Periodic Table

Atomic Masses

The Mass of an Atom Can Be Determined From the

Masses of the Neutrons and Protons in It’s Nucleus

-This is a very small number

One 12C atom weights 1.99 X 10-23 grams

Atomic Mass Unit (AMU)

1amu = 1.66 X 10-24g

Atomic Masses

Note: the Atomic Weights of the Elements (the Mass

of 1 Atom in Amu’s) Are Given in the Periodic

Table.

Note: Carbon has 3 Isotopes

12C, 13C, & 14C

12.01 Is the Weight of the

Natural Occurring Mixture of

These Isotope

Why Is the Atomic Weight of a Carbon = 12.01 Amu?

Atomic Masses

What is the Mass of a 100 Iron Atoms?

1 Iron atom = 55.85 amu

amuFe

amuFe 585,5)

85.55)(100(

Always Include Units in Your Calculations!

Atomic Masses

Why is This So Important?

Because by knowing the Mass of an

Element (a Measurable Quantity), We can

Determine the Number of Atoms Present!

How Do We Do This On the

Macroscopic (Everyday) Scale?

The Mole

Mole = SI Standard

Unit for Quantity

1 Mole = 6.022 x 1023

Avogadro’s Number = 6.022 x 1023

The Atomic Weight of an Element Is the Number of

Grams One Mole of That Element Would Weight

1 mole of Carbon weights 12.011 grams!

(Amadeo Avogardo)

The Mole

How Many atoms are in 48.044

grams of Carbon?

OR

CatomsxCmol

CatomsxCmols

24

23

10409.2)10022.6

(4

CmolsCg

CmolCg

0000.4)

011.12

1(044.48

The Mole

Group Exercises:

How many mols are in 48 g Tin?

How many atoms is this?

How much does 4.395 mols gold

weight?

What is the mass of a billion

radon atoms?

Atomic Masses

How Can Chemists Measure Atomic Masses?

Mass Spectroscopy

2.3 IsotopesThe Identity of an Element Is Determined by the

Number of Protons It Possesses

Isotopes - Elements With the Same Number of

Protons but Different Numbers of Neutrons.

Different Atoms of Same Element Can Have

Different Masses.

AzX

A= Mass Number (# of P and N)

Z=Atomic Number (# of P)

X= Elemental Symbol

•Isotopic Notation:

Isotopes#Protons = Z, the Atomic Number, and Can

Also Be Determined by X, the Atomic

Symbol

(All Carbon Isotopes Have Z=6 and X=C)

#Neutrons = A - Z

AzX

A= Mass Number (# of P and N)

Z=Atomic Number (# of P)

X= Elemental Symbol

Isotopes

Determine the Number of Protons , Neutrons and

Electrons for the 3 Isotopes of Oxygen

168O:

178O:

188O:

P = 8, N = (16-8) = 8, e = 8

P = 8, N = (17-8) = 9, e = 8

P = 8, N = (18-8) =10, e = 8

Note: For Neutral Atoms, the # of

Protons Equals the # of Electrons

Isotopes of Monatomic Ions

Determine the Number of Protons ,

Neutrons and Electrons for the 2 Ions:

168O-2 :

5525Mn+7

P = 8, N = (16-8) = 8, e = 10

P = 25, N = (17-8) = 30, e = 18

Note: For ions there is one less electron than

proton for each [+] charge and one more

electron than protons for each [–] charge.

Radioactive Isotopes

Many Isotopes Are Unstable and Undergo

Radioactive Decay.

If the Number of Protons Changes During

a Radiative Process, the Identity of the

Element Changes.

We Will Cover These in Chapter 23 of the

Text

2.4 Atomic Weight

Given in the Periodic Table.

Note: Carbon has 3 Isotopes

12C, 13C, & 14C

12.01 Is the Weight of the

Natural Occurring Mixture of

These Isotope

Why Is the Atomic Weight of a Carbon = 12.01 Amu?

% Isotopic

Abundance

# of Atoms of Given Isotope(100)

Total Atoms of All Isotopes=

Isotope Abundance

Natural Samples of Pure Elements Contain a

Mixture of that Elements Isotopes

Magnesium has 3 Isotopes with the

following masses and % Abundances,

23.985 amu (78.99%),

24.986 amu (10.00% )

25.983 amu (11.01%)

Determine the average atomic weight of

Magnesium

Class Problem

.7899(23.985) .1000(24.986) .1101(25.983)

24.305amu

Lithium has has only 2 isotopes, 6Li

(with mass of 6.0151) and 7Li (with

mass of 7.0160)

What is the % Abundance for

Each Isotope?

Class Problem

2 Unknowns Requires 2 Equations

X = fraction of 6Li

Y = fraction of 7Li

X(6.0151) +Y(7.0160) = 6.941

Let

X(6.0151) +(1-X)7.0160 = 6.941

X+Y=1 Y=1-X

6.941 7.0160.07493; 1 1 .07493 .9251

6.0151 7.0160X Y X

6Li = 7.5% 7Li = 92.5%

Note, 7.941-7.0160 = -0.075

2.5 Periodic Table

Dmitri Mendeleev, noted

various elements had

similar properties, with a

periodic relationship to the

increasing atomic number.

This resulted in the

Modern Periodic Table

Periodic Table

Periodic TableBe Able to Identify the Following Groups or

Families From the Periodic Table

1. Alkali Metals (Group 1)

2. Alkaline Earth Metals (Group 2)

3. Halogens (Group 7)

4. Noble Gases (Group 8)

5. Transition Metals

6. Lanthanides

7. Actinides

Metals, Nonmetals & Metalloids

Periodic Table

Physical Properties of Metals

1. Conduction of Heat and Electricity

2. Malleability (can be hammered into

thin sheets)

3. Ductility (can be pulled into wires)

4. Lustrous (shiny) appearance

Nonmetals

-Poor Conductors

-Nonlusterous

-Nonmalleable

-Solids are brittle

-Some are brightly

colored

Natural States of The Elements

Liquids - Hg, Br2

Gasses - Noble Gasses and lighter

diatomics, H2 , N2 , O2 , F2 , Cl2

Solids - All other elements

Natural States of The Elements

Elements Which Often Appear in Pure Form

A) Noble Metals (Ag, Au, & Pt)

B) Noble Gases (He, Ne, Ar, Kr Xe & Rn)

“Noble” Implies “Unreactive”

All elements with Atomic # > 83

are Radioactive

7 DiatomicsKnow the Diatomic’s and Their Natural States

Hydrogen (H2) gas (colorless)

Nitrogen (N2) gas (colorless)

Oxygen (O2) gas (pale blue)

Fluorine (F2) gas (pale yellow)

Chlorine (Cl2) gas (pale green)

Bromine (Br2) Liquid (red/brown)

Iodine (I2) Solid (purple)

Group 1A: Alkali Metals

H, Li, Na, K, Rb, Cs & Fr

Very Reactive – Do Not Exist as

Pure Elements in Nature

Form Oxides of Formula

A2O

Group 2A: Alkali Earth Metals

Be, Mg, Ca, Sr, Ba & Ra

Reactive – Form Alkaline

Solutions (Basic)

Form Oxides of Formula

AO

Group 3A

B, Al, Ga, In, Tl

Form Oxides of Formula

A2O3

Boron comes from Borax

Group 3A

B, Al, Ga, In, Tl

Form Oxides of Formula

A2O3

Boron comes from Borax

Group 4A

C, Si, Ge, Sn, Pb

C-non metal

Si, Ge – metalloids

Sn, Pb - metals

Group 4A

C, Si, Ge, Sn, Pb

Carbon forms Allotropes

BuckminsterfullereneDiamond Graphite

(Different Forms of the Same Element)

Group 5A

N, P As, Sb, Bi

N, P – Non metals (essential to life, N is most

abundant element in atmosphere

As, Sb – metalloids

Bi – metal, heaviest non radioactive element

Group 6A

O, S, Se, Te, Po

O, S, Se – nonmetals

Te – metalloid

Po – Radioactive metal

Chalcogens

Chalk Formers

Group 7A

F, Cl,Br, I, At

F,Cl, Br, I – nonmetal diatomics

At – often classified as

radioactive metalloid

Halogens

Salt Formers

Group 8A

He, Ne, Ar, Kr, Xe

“Inert Gases” – do not tend to

form compounds

Noble Gases

Part II: Compounds

4 Types of Chemical Bonds

1. Ionic Bonds - Ionic compounds between metals and

nonmetals consisting of [+] cations and [-] anions.

2. Covalent Bonds - Classical “molecules” where

valence electrons are shared between two

nonmetals

4. Metallic Bonds - Pure metals and alloys where

delocalized free electrons hold together the positive

nuclei.

3. Polar Covalent Bonds - Covalent bonds with ionic

character in that the electrons are not equally shared.

2.6 Molecules & Compounds

-Pure Substance Composed of More Than 1 Atom

-Formulas Describe the Elemental Composition

-Common Names Like Baking Soda or Sugar

Tell Us Nothing About Their Composition

-Their Formulas Provide Information on Their

Chemical Composition

Chemical Formulas

Sugar = C12H22O11

Baking Soda = NaHCO3

This tells us that:

-one molecule of sugar has 12 carbon atoms,

22 hydrogen atoms and 11 oxygen atoms

-Baking Soda has 3 atoms of Oxygen and 1

atom each of Sodium, Hydrogen and Carbon

Molecular Formulas

Molecular Formulas - only tell of the

elemental composition

-both ethanol and diethyl ether have the same

molecular formula, C2H6O

-Yet they are different

Structural Formula

Ethanol = CH3CH2OH

Dimethyl Ether = CH3OCH3

H-C-C-O-H

H H

H H

H-C-O-C-H

H H

H H

bonds

Molecular Models

Ball & Stick Persective Drawing

behind

plane

in front of

plane

METHANE (CH4)

Molecular Models - Water

space-filling models

ICE

2.7 Ions & Nomenclature

Charged Atoms Occur When an Atom Losses

or Gains Valence Electron(s)

Two Types:

Anions - Negative Ions (Gain

Electrons)

Cations - Positive Ions (Lose Electrons)

Note: Charge Is Indicated by a +/- Post

Superscript

Ions

Ionic Charge and the Periodic Table:

Metals form Cations (tend to lose electrons)

Nonmetals for Anions (tend to gain electrons)

Noble Gasses Do Not Tend to Form Ions

Cations Anions

+1 Alkali Metals, H, Ag -1 Halogens, Hydrogen

+2 Alkaline Earths, Zn, Cd -2 Group VI Nonmetals

+3 Aluminum -3 Group V Nonmetals

Note - Monatomic Ions Tend to Have Charge Required to

Become Isoelectronic With Nearest Nobel Gas

(Have Same # of Electrons As the Noble Gas)

Ions

Polyatomic Ions

-Charged Covalent Compounds

H-N-H

O

O-S-O

H

H+

NH4+

O -2

SO4-2

Ionic Compounds & Nomenclature

Ionic Compounds Form Crystal Lattices

Resulting From Electrostatic Interactions

Which Maximize the +/- Attractions While

Minimizing the +/+ and -/- Repulsions

Formula Unit

-Ionic Compounds Are Identified by

Their Formula Unit

-Formula Unit Represents the Lowest

Whole # Ratio of Ions in the

Crystal Lattice

Nomenclature

1. Ionic Compounds (between metals and

nonmetals)

2. Covalent Compounds (between two

nonmetals)

3. Compound with Polyatomic Ions

4. Acids

5. Compound Formulas from Names

Binary Compounds

Binary Compounds - Contain two types

of atoms, can be ionic or covalent

Examples:

H2O, CO, CO2, CoCl2, FeO, Fe2O3, NaCl

Binary Covalent Binary Ionic

Binary Ionic Compounds

Form Between a Metal and a Nonmetal, -Have Monatomic Ions

-Form Crystal Lattice Type Solid Structures

-Oxidation # of a Monatomic Ion Is It’s Charge

Cation - [+] ion, (Metals Form Cations)

Type 1 Metal (forms only 1 cation) Na+,

Type 2 Metal (forms 2 or more cations) Fe+2, Fe+3,...

Anion - [-] ion, (Nonmetals Form Anions)

ex: Cl-, O-2, N-3,...

Binary Ionic Compounds

Type I Metals -have single charge in all ionic

compounds (Invariant Oxidation #)

Know Following Invariant Oxidation #’s:

+1 Alkali Metals, H, Ag -1 Halogens, Hydrogen

+2 Alkaline Earths, Zn, Cd -2 Group VI Nonmetals

+3 Aluminum -3 Group V Nonmetals

Note - “A” Group Monatomic Ions Tend to Have Charge

Required to Become Isoelectronic With Nearest Nobel

Gas (Have Same # of Electrons As the Noble Gas)

Binary Ionic Compounds

Compounds with Type I Metals

Rules:

1. Name Metal First

2. Name Nonmetal 2nd With -ide Suffix

Ex: Table Salt, NaCl = Sodium Chloride

BaCl2 = Barium Chloride

Binary Ionic CompoundsIonic Formulas from Names

Rules:

1. Ionic Formula Must Be Neutral (Principal of

Charge Balance)

2. Subscripts After Elemental Symbol Indicate #

of That Element in Formula

3.Correct Formula Indicates Lowest Whole #

Ratio of Anions to Cations

Magnesium Chloride = MgCl2 not Mg2Cl4

Binary Ionic CompoundsIonic Formulas from Names

Trick: Set # of Anions = Charge of Cation

and # of Cations = Charge of Anion

Beware That This Is the Lowest Whole # Ratio

Consider Aluminum Oxide,

Charge of Aluminum =+3 and Oxygen = -2,

Al2O3

(2x+3)+(2x-3)=0

Binary Ionic Compounds

-1 -2 -3

+1

+2

+3

1. In first column, chose cations of given charges and

place symbol underneath charge, in first row, place

symbol of appropriate anions underneath each charge.

2. Work in Groups

of 3, each student

picks one row, for

each cell of their

row, write name in

upper half, and

formula in lower

half.

3. Check each others

work!

Na

Cl

Sodium Chloride

NaCl

Binary Ionic Compounds

Type II Metals -have multiple charges in ionic

compounds (Variable Oxidation #’s)

Know Following Variable Oxidation State Ions

Fe+3 Iron(III) - ferric Sn+4 Tin(IV) - stannic

Fe+2 Iron(II) - ferrous Sn+2 Tin(II) - stannous

Cu+2 Copper(II) - cupric Pb+4 Lead(IV) - plumbic

Cu+1 Copper(I) - cuprous Pb+2 Lead(II) - plumbous

Co+3 Cobalt(III) - cobaltic Hg+2 Mercury(II) - mercuric

Co+2 Cobalt(II) - cobaltous Hg2+2 Mercury(I) - mercurous

Binary Ionic Compounds

Type II Metals

Rules:

1. Name Metal First

2. Indicate Metal’s Charge (Oxidation State)

in Roman Numerals

2. Name Nonmetal 2nd With -ide Suffix

FeCl2 = Iron(II)Chloride

FeCl3 = Iron(III)Chloride

Naming Covalent Compounds

-Compounds Between Nonmetals Do Not Form

Ions but Share Electrons (Forming Covalent Bonds)

- These Are the Classical “Molecules” (in

Contrast to Ionic Crystal Lattices)

- Use Greek Prefixes to Indicate Number of

Each Type of Atom in Molecules

- Write “More Metallic” Nonmetal First

- If More Than 2 Atoms, Write in Order

of Connectivity (SCN vs. SNC)

Naming Covalent Compounds

Greek Prefixes

1 mono- 6 hexa-

2 di- 7 hepta-

3 tri- 8 octa-

4 tetra- 9 nona-

5 penta- 10 deca-

Note: often do not use prefix mono-

Naming Covalent Compounds

Group Problems: Name or write the

following formula’s

CCl4

SO3

Dinitrogen trioxide

Diphosphorous pentaoxide

- Carbon tetrachloride

- Sulfur trioxide

- N2O3

- P2O5

Polyatomic Ions

Many Ionic Compounds Have Covalently

Bonded Polyatomic Ions

Rules Are Essentially Same, but Use the Name

of the Polyatomic Ion, That Is;

Cation First (Charge If Variable) Anion Last

Polyatomic Ions

NH4+ Ammonium C8H4O4

-2 Phthlate

CH3COO- Acetate SO4-2 Sulfate

NO2- Nitrite SO3

-2 Sulfite

NO3- Nitrate CrO4

-2 Chromate

OH- Hydroxide Cr2O7-2 Dichromate

ClO4- Perchlorate PO4

-3 Phosphate

ClO3- Chlorate PO3

-3 Phosphite

ClO2- Chlorite AsO4

-3 Arsenate ClO-

Hypochlorite CN- Cyanide

MnO4- Permanganate SCN- Thiocyanide

CO3-2 carbonate C2O4

-2 Oxalate

HCO3- bicarbonate O2

-2 peroxide

Polyatomic Ions

Tip for Remembering Charges - Many Oxyanions (Polyatomic Ions With Multiple Oxygens) Have the

Same Charge As the Monatomic Nonoxygen Ion

Sulfide = -2 As Does Sulfate & Sulfite

Chloride = -1 As Does Perchlorate, Chlorate, Chlorite

& Hypochlorite

Phosphide = -3 Does Phosphate & Phosphite

Note, Nitrogen Is an Exception:

Nitride = -3, Nitrate & Nitrite = -1

Polyatomic Ions

Note: Oxyanions of Heavier Halogens

(Bromine and Iodine) Form Same Types

of Compounds As Chlorine

BrO4- perbromate IO4

- periodate

BrO3- bromate IO3

- Iodate

BrO2- bromite IO2

- Iodite

BrO- hypobromite IO- hypoiodite

Acid Nomenclature

Compound which can release H+ in water are

called Acids. (aqueous, aq.,- means dissolved in water)

Nomenclature is related to that of Ionic Compounds

-If salt is binary, it forms binary acids :

-ide => hydro(anions elemental name)-ic acid

If salt has oxyanions it forms oxyacids with endings:

ate => ic acid

ite => ous acid

Sulfur’s Sodium Salts and Acids

Na2S - Sodium Sulfide

H2S(aq) - Hydrosulfuric AcidNote: H2S(g) - Hydrogen Sulfide (gas)

Na2SO4 - Sodium Sulfate

H2SO4(aq)- Sulfuric Acid

Na2SO3 - Sodium Sulfite

H2SO3(aq)- Sulfurous Acid

Acid Nomenclature

Name or give Formulas for Following Acids:

Acetic Acid (vinegar) -

Chloric Acid -

Hydrochloric Acid -

H3PO4 -

HCN -

CH3COOH

HClO3

HCl

Phosphoric Acid

Hydrocyanic Acid

Acid Salts

-Ions with more than one negative site can

combine with more than one cation.

- Acid salts result when one cation is a

proton

Lets look at all of the compounds that

can form from phosphate

Phosphate = PO4-3

Na3PO4= Sodium Phosphate

Na2HPO4= Sodium Hydrogen Phosphate

NaH2PO4= Sodium Dihydrogen Phosphate

H3PO4 = Phosphoric Acid

Name the compounds that can form from

hydrogen, potassium and phthlate; and

identify them as salts, acids or acid salts

Hydrated Salts

-Salts which incorporate water into their

crystal structure are hydrated salts

-Include the water in the compound formula

-Greek prefixes indicate the # of waters of

hydration

CaCl2.6H2O = Calcium Chloride hexahydrate

What is the name of

CuSO4.5 H2O ?

Copper(II) sulfate pentahydrate

2.9 Molar Mass

Molar Mass - Is the Mass (in Grams) of

One Mole of a Substance (often called molecular weight or formula weight)

What is the Molar Mass of Sulfate?

1 mole SO4-2: has 1 mole S and 4 moles O

1 mole S = 32.07 g

4 mole O = 4(16) = 64.00g

1 mole SO4-2 = 96.07g

Molar Mass

How many moles are in 5.86g of AgCl?

1 mol AgCl = 143.3g, ie.,

MM(AgCl)=143.3g/mol

AgClmolAgClg

AgClmolAgClg

041.0)

3.143)(86.5(

Note how we use Dimensional Analysis and the

formula weight as a mass-->mole conversion factor

Molar Mass

Calculate the Molar Mass of:

1. Benzene (C6H6)

2. Acetylene (C2H2) :

3. Water

4. Hydrogen Peroxide

78g/mol

26g/mol

18g/mol

34g/mol

Molar Mass

What is the Molar Mass for

Aluminum Sulfate?

Al2(SO4)3 has 2 Al Cations and 3 Sulfate

Anions2(Al) = 2(26.98) = 53.96g

3(SO4)3 = 3(96.07) = 288.21g

1 mol Al2(SO4)3 = 342.17g

Note: We used the Molar Mass of Sulfate which we

determined in the Previous Problem

Molar Mass

Calculate Molar Mass for:

1. Sodium Bicarbonate

2. Ammonium Nitrate

3. Potassium Perchlorate

4. Silver Chloride

84.02 g/mol

80.08 g/mol

138.54g/mol

143.32g/mol

Molar Mass

1. What is the mass of 5.80 moles

Sodium Bicarbonate?

2. What is the mass of 6.82 moles

Ammonium Nitrate

3. How many moles are in 51.50 g of

Potassium perchlorate?

4. How many moles are in 100.0 g of

Silver Chloride?

420.g

546g

.3717 moles

.6977 moles

2.10 Percent Composition

Mass Fractionof a given element =

Mass of element in compound

Mass of compound

Mass Percentof a given element

Mass Fractionof a given element X 100=

What Is the Mass Fraction of an Element in a

Compound?

What is the Mass Percent of an Element in a

Compound?

Percent Composition

1. What Is the Sum of the Mass Percents

of All the Elements in a Compound?

What is the mass percent of H, C & O

in Acetic Acid (Vinegar)

CH3COOH

2. Does the Percent Composition Change

as the Quantity of Substance Changes?

100%

NO

Percent Composition

Consider 1 mole CH3COOH

MCarbon = 2(12.01) = 24.02g

MOxygen = 2(16.00) = 32.00g

MHydrogen = 4(1.01) = 4.04g

MAcetic Acid = 60.06g

60.06

%C=24.02(100) = 40.00%

60.06

%O=32.00(100)=53.28%

60.06

%H=4.04(100)=7.0%How Do We Get the

Percent Composition

From the Above Data?

Percent Composition Problems

1. Determine the % Composition of the

Sulfate Ion.

2. Determine the % Composition of

Aluminum Sulfate

3. Determine the Percent Composition

of Benzene (C6H6)

4. Determine the Percent Composition

of Acetylene (C2H2)

33.3% S, 67.7% O

15.8% Al, 28.1% S, 56.1%O

7.7% H, 92.3% C

7.7% H, 92.3% C

Compound Formulas

1. Empirical Formula - Simplest Whole #

Ratio of the Various Elements in a

Compound

Note: Benzene (C6H6) & Acetylene (C2H2)

Have Different Molecular Formulas,

but the Same Empirical Formula (CH)!

2. Molecular Formula - the Actual Ratio of

the Various Elements in a Compound or

Molecule

Empirical Formulas

-Empirical data is based on observation and

experiment, not theory

-Opposite Calculation to Percent Composition

-For Empirical Formulas, You Know the Masses

of Each Element From Experimental Data,

So Determine Simplest Formula

-In % Composition, You Start Knowing the

Formula, and Determine the Mass %

Empirical Formulas

1. Obtain Mass of Each Element (in Grams)

If given % Composition, assume 100 g of substance

2. Calculate # of Moles of Each Element Present

From Masses and Atomic Weights

3. Divide # of Moles by the Moles of Least

Present Element (Forcing It to 1)

4. Multiple the Results of Step 3 by Smallest

Integer Which Will Convert Them All to

Whole Numbers

Empirical Formulas

Determine the Empirical Formula of Aspirin,

Which Was Analyzed and Found to Have

a Mass Percent Composition of 60.0%C,

4.48%H and 35.5%O.

Empirical Formulas

1. Assume 100 grams of sample, giving

MC=60.0g, MH=4.48g, MO=35.5g

2. Calculate Moles Present

moles C = 60.0g(1mol/12.0g) = 5

moles H = 4.48g(1mol/1.00g)=4.48

moles O = 35.5g(1mol/16.0g) =2.22

Empirical Formulas

C: 5/2.22 = 2.25

H: 4.48/2.22 = 2.02

O: 2.22 /2.22 = 1.00

3. Divide By

Smallest Mole

Fraction (2.22)

4.Multiply by smallest

Integer Forcing all

results from step 3 to

be whole numbers

C: 2.25(4) = 9

H: 2.02(4) = 8

O: 1.00(4) = 4

Empirical Formula of Aspirin = C9H8O4

Molecular Formulas

In Order to Determine the Molecular

Formula, We Need to Know the

Molar Mass in Addition to the

Empirical Formula

The Molecular Formula will be an Integral

Multiple of the Empirical Formula

For an Ionic Compound, the Formula Weight

Corresponds to the Empirical Formula

Determine the molecular formula of a compound

containing H, O & C. A 50.00 g sample has 23.88g

of both carbon & oxygen and the compound has a

molar mass of 603g/mol.

1. Determine moles in sample

123.88 1.988

12.01

123.88 1.493

16.00

1.0082.24 2.26

molmoles C gC mol C

g

molmoles O O mol O

g

molmoles H gH mol H

g

note: mass hydrogen = total mass-mass oxygen –mass carbon

= 50.00g-2(23.88g) = 2.24g

Step 2: Divide by Smallest number

1: 2.24 1.5

1.493

1: 1.988 1.33

1.493

1: 1.493 1

1.493

H

C

O

Trick: Convert to Fractions

1 3: 2.24

1.493 2

1 4: 1.988

1.493 3

1: 1.493 1

1.493

H

C

O

Tips in converting decimals to fractions

1 1decimal expresssion

decimal expressionX

X

Know: 0.50 = 1/2 0.33 = 1/3

0.25 = 1/3 0.20 = 1/5

Step 2: Divide by Smallest number

1: 2.24 1.5

1.493

1: 1.988 1.33

1.493

1: 1.493 1

1.493

H

C

O

C8H9O6

Trick: Convert to Fractions

1 3: 2.24

1.493 2

1 4: 1.988

1.493 3

1: 1.493 1

1.493

H

C

O

1 3: 2.24 (6) 9

1.493 2

1 4: 1.988 (6) 8

1.493 3

1: 1.493 1(6) 6

1.493

H

C

O

Step 3: Find lowest whole number ratio by

factoring out the denominators

Step 4: Calculate Molecular Formula

from Empirical Formula Using Masses

EF = Empirical Formula EM = Empirical Mass

MF = Molecular Formula MM = Molecular Mass

MF = n(EF) & MM=n(EM), n=1,2,3…

EM(C8H9O6) = 201g/mol, MM=603g/mol

6033

201

gmol

gmol

MMn

EM

MF = n(EF)=3(C8H9O6) = C24H27O18

Acidum Formicum – Formic AcidWhat is the empirical formula

of Formic Acid?

Obtained through the distillation of Formica rufa

Empirical Formula of Formic Acid

from Combustion Data

Sample (CxHyOn)

furnace H2O absorber

CO2

absorber

excessoxygen

2.4541g

formic acid0.9606g 2.3482g

mH = .9606g(2gH/18gH2O ) = .1067g H

mC = 2.3482g(12gC/44gCO2 ) = .6404g C

mO = 2.4541g - .1067g - .6404g = 1.7071gO

Step 1: Determine Mass of each element

Step 2: Determine Mole of Each Element1

0.1067 0.10671.008

10.6404 0.0533

12.01

11.7071 0.10670

15.9994

molmoles H gH mol H

g

molmoles C gC mol C

g

molmoles O mol O

g

Step 3: Divide by Smallest Number

1: 0.1067 2

0.0533

1: 0.0533 1

0.0533

1: 0.1067 2

0.0533

H

C

O

CH2O2

2.11 Hydrated Salts

Hydrated salts absorb water into

the crystal lattice- this is the

“water of hydrations”

-Stoichiometric Proportions

(water of hydration is of

integer proportions)

Anhydrous Salt – “Without Water”, Form

of Salt without water of hydration, forms

different type of crystal

Hydrated Salts

Sodium Carbonate Forms a Hydrated Salt

Na2SO4.xH2O

What is the Formula if a 0.767g

sample was dried to a constant weight

of 0.284 g?

Hydrated Salts

Mhydrated salt = Manyhdrous salt + M water

M water = Mhydrated salt - Manyhdrous salt

M water = 0.767g – 0.284g = 0.483g

2

2

2

2

2 3

2

2

32 3

, ,

0.4830.0268

18

0.2840.00268

106Na CO

H O

H O g gmol molH O

Na CO

g gmol molNa CO f

n mole M mass fw formula weight

M g gn mole H O

fw

M g gn mol Na CO

fw

Hydrated Salts

Dividing moles water and moles Sodium

carbonate by species with least number of

moles (0.00268 moles Na2CO3):

Na2CO3.10 H2O