summary of definitions

Summary of Definitions• Atomic number (Z) = # protons• Mass number (A) = # protons + # neutrons• Atomic mass (u) : the weighted average of the mass

number of all natural isotopes of the element.– Mass numbers (A) are whole– Atomic masses (u) are decimal numbers

• In a neutral atoms: # electrons = # protons• In an ion (charged atom)

– The ionic charge* = # protons - # electrons

* The ionic charge is sometimes called the “oxidation number”

Summary of Models• Know the Bohr-Rutherford Model / Simplified Model• Know

– Electron: negatively charged particle, very light– Proton: positively charged particle ~1 AMU– Neutron: neutral particle ~1 AMU

Module 1, Lesson #8

• Chemistry notations• Periodic Table

Symbols and Meaning

Se2

2

79

34

–

Number of atoms per molecule

Atomic number (Z)Number of protons

Isotope mass number (A)(if no decimal)

Valence number(without sign) most likely # bonds

With sign, ionic charge or oxidation number

Symbol

45

subtract

Number of neutrons (A-Z)

78.9

Average atomic mass (u) if writtenwith a decimal

Uncharged # electrons = ZActual # electrons = ( Z – oxidation #)

Symbols and Meaning(Summary for Notes)

Se2

2

79

34

–

Number of atoms per molecule

Atomic number (Z)Equals # of protons

Mass number (A) Valence number(no integer sign)# of bonds

With sign, ionic charge or oxidation number

Symbol

45Number of neutrons(Subtract: A-Z)

Symbols and Meaning

If it is a decimal number it is called atomic mass (u)

Uncharged # electrons = ZActual # electrons = ( Z – oxidation #)

Quick exercise: What does each notation below mean?

• Magnesium-25 ion, 12p, 13n, 10e

• Sulfur molecule containing 8 atoms of sulfur (each atom has 16p, 16e)

• Bromine-80 ion, 35p, 45n, 36e

• Typical uncharged Chlorine atom, 17p, 17e and on the average, 18.5 neutrons

• A molecule containing 6 atoms of carbon, 12 atoms hydrogen, 6 atoms oxygen

Mg25 2+

S8

80 Br 1-

?17

35.5

C6H12O6

Cl

Representative (A Families)and Transitional (B Families) 1 2 13 14 15 16 17 18

1 H He

2 Li Be B C N O F Ne

3 Na Mg Al Si P S Cl Ar

4 K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr

5 Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe

6 Cs Ba Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn

7 Fr Ra Rf Ha Sg Bh Hs Mt Ds Rg Cn ? ? ? ? ? ?

6R La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu

7R Ac Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr

Active Metals

Families in the Periodic Table(Traditional numbering in Roman numerals)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18

1 H He









The lower families are sometimes called “B” families.

Families, or groups run vertically in the table.

Fam

ily

1: A

lkal

i m

etal

Fam

ily

2: A

lkal

ine

Ear

ths

Fam

ily

3: B

oro

n f

amil

y

Fam

ily

4: C

arb

on

Fam

ily

Fam

ily

5: N

itro

gen

fam

ily

Fam

ily

6: O

xyg

en f

amil

y

Fam

ily

7: H

alo

ge

ns

Fam

ily

8: N

ob

le g

ases

Coin

meta

ls

I II III IV V VI VII VIII

Scan

diu

m

Tit

an

ium

Van

ad

ium

Ch

rom

ium

Man

gan

ese

Zin

c F

am

ily

I

B

II

B

III

B

IV

B

V

B

VI

B

VII

B

VIII

B

Iron Triad*

Lanthanides (period 6 rare earths)

Actinides (period 7 rare earths)

Palladium

Platinum

Triads

*Iron Triad AKA magnetic metals

Rows or Periods (number of shells / energy levels)

1 2 13 14 15 16 17 18

1 H He









1st Period = One Shell

2nd Period = Two Shells

3rd Period = Three Shells

4th Period = Four Shells

5th Period = Five Shells

6th Period = Six Shells

7th Period = Seven Shells

Six Shells

Seven Shells

Regions of the Periodic Table(metals vs. non-metals)

1 2 13 14 15 16 17 18

1 H He

2 Li Be Heavy line is the traditional boundary line between metals and non-metals. The gray elements are the metals, the purple are non-metals. The lighter ones are sometimes called metalloids or semimetals.

B C N O F Ne








Active Metals

METALS , NON-METALS and METALLOIDS

METALLO

IDS METALS

NON-METALS

Modern Regions of the Periodic Table(divided into more groups)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18

1 H He









Rare Earth Elements (metals)

Other Non metals

Transition Elements (metals)

Alkaline Earth (metals)

Alkali Metals

Other Metals

Halogens

Noble (Inert) gas

Metalloids

Activity in the Periodic Table 1I

2II

3 4 5 6 7 8 9 10 11 12 13III

14IV

15V

16VI

17VII

18VIII

1 H He


3 Na Mg

Al Si P S Cl Ar





6R La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm

Yb Lu

7R Ac Th Pa U Np Pu Am Cm Bk Cf Es Fm Md

No Lr

Active non-metals

Active Metals

Inactive

F

Most Active

*Most Active Metal and Non-metal

Fr Most Active Metal

Atomic Size in Periodic Table1 2 13 14 15 16 17 18

1 H He



4 K Ca Sc Ti V Cr Mn Fe Co Ni Cu zn Ga Ge As Se Br Kr


6 Cs Ba Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At R

7 Fr Ra Rf Ha Sg ? ? ? ? ? ?



Big Atoms

Smallest Atom

Bigger Atoms Smaller Atoms

B

igg

er A

tom

s

Sm

alle

r A

tom

s

Biggest Atom

Valence Electrons(Electrons in the outer shell)

1 2 13 14 15 16 17 18

1 H He









Active Metals

*Filled orbitals: full s-orbital for He, full s and p-orbitals for all others, (but not necessarily filled d or f orbitals )

On

e

Two

Th

ree

Fo

ur

Fiv

e

Six

Sev

en

Fu

ll s

& p

-orb

ital

s (8

)(2

)

1e- 2e- 3e- 4e- 5e-6e-7e- Full*

Mostly 2 electrons in outer shell (but it can vary from 1 to 5)

Most Probable Valences (Actual valences can vary. Valence with sign is called oxidation number)

1 2 13 14 15 16 17 18

1 H He









Active Metals

*Filled orbitals:full s-orbital for He, full p-orbitals for all others

On

e

Two

Th

ree

Fo

ur

Th

ree

Two

On

e

ZE

RO

(2)

1+ 2+ 3+ 4± 3- 2- 1- 0

Many transition element have a valence of 2+, but it can vary from 1+ to 5+. Some

transition elements are polyvalent

Relationship betweenperiodic table and orbitals

1 2 13 14 15 16 17 18

1 H He









Active Metals

*For these elements, the outermost electrons are in an s orbital

s-o

rbit

als

p-orbitals

d-orbitals

f-orbitals

s

*For these elements, the outermost electrons are in a p orbital *For these elements, the outermost electrons are in a d orbital *For these elements, the outermost electrons are in an f orbital

Summary

• Know several different ways of viewing the periodic table, for example:– By representative vs. transitional regions– By metals vs. non-metals (vs. metalloids)– By major Families and regions

• Know the relationship between the periodic table and:– Valence electrons– Valence numbers (oxidation numbers)– Number of shells

Module 1, Lesson 9

• Naming compounds• Common names vs. systematic names• Binary compounds• Covalent vs. Ionic Compounds

Compound Names

• Every compound must have a unique name.

• Many compounds have common names, but some don’t, and some common names are misleading.

• A systematic name is devised for each compound by following sets of rules for systematic names. (AKA scientific names)

Examples of Common and Systematic names

• Common name formula systematic name• Water H2O dihydrogen monoxide• Lime CaO calcium oxide• Slaked lime Ca(OH)2 calcium hydroxide• Lye NaOH sodium hydroxide• Potash K2CO3 potassium carbonate• Table salt NaCl sodium chloride• Laughing gas N2O dinitrogen monoxide• Oxygen gas O2 dioxygen • Ozone gas O3 trioxygen• Ammonia NH3 nitrogen trihydride• Methane CH4 carbon tetrahydride

– Some compounds only have systematic names:• Carbon dioxide CO2 carbon dioxide• Carbon tetrachloride CCl4 carbon tetrachloride

– Organic compounds use a completely different naming system:• Table sugar C12H22O11 sucrose

(glucose-fructose disaccharide)

Don’t copy,

Read!

Some common, systematic and other acceptable names

Common Name Formula Systematic Name Other IUPAC names used

Water H2O Dihydrogen monoxide Hydrogen oxide, dihydrogen oxide

Lime CaO Calcium oxide

Slaked Lime Ca(OH)2 Calcium hydroxide

Lye NaOH Sodium hydroxide

Table salt NaCl Sodium chloride

Potash K2CO2 Potassium carbonate Potassium carbonate

Laughing gas N2O Dinitrogen monoxide Nitrous oxide

Oxygen gas O2 Dioxygen Oxygen

Ozone gas O3 Trioxygen Ozone

Ammonia or Azane NH3 Nitrogen trihydride, Trihydrogen nitride, azane

Methane CH4 Carbon tetrahydride Methyl hydride, methane

Carbon dioxide CO2 Carbon dioxide Dioxidocarbon

Carbon tet. CCl4 Carbon tetrachloride Tetrachloromethane

Hydrazine N2H4 Dinitrogen tetrahydride Hydrazine, diamine, diazane

Sewer gasStink damp, egg gas

H2S Dihydrogen sulphide Hydrogen sulphide, sulfane, Sulphur hydride

Binary Compounds vs. Ternary Compounds

• Binary Compounds– Contain only two elements– Systematic names always end in –ide.– Systematic names are usually quite simple,

containing the names of the two elements which make the compound• Example: sodium chloride (NaCl)

• Ternary Compounds– Contain more than two elements– Contain “radicals”

Rule Sets• There are Three main sets of rules for

naming compounds (with some additional variations)

The Covalent Rules are for naming covalent compounds, not including organic covalent compounds

The Ionic Rules are for naming ionic compounds with a variation for ternary ionic compounds

The Organic Rules are for naming organic compounds, but these are not studied this year

Ionic vs Covalent

• Ionic compounds contain ions (charged particles)– A metal atom with a non-metal (in salts) eg: NaCl

– A hydrogen ion with a non-metal (in acids) eg: HCl

– An ammonium salt starts with NH4… eg: NH4Cl

• Covalent compounds contain molecules (uncharged)– They are also called molecular compounds.– Include most compounds with only non-metals.

• Excluding acids and ammonium salts, which are ionic• Organic compounds are also contain molecules,

but they have lots of carbon (usually more than one carbon atom)

• Sugars, alcohols, hydrocarbons,aldehydes, ketones, proteins, nucleic acids, etc.

The Covalent Rules

• Covalent (or molecular) compound names often include prefixes.– Mono, di, tri, tetra, penta, hexa, hepta, octo…

• The name is always based on the formula– Eg:

• N2O4 dinitrogen tetroxide

• SF6 sulfur hexafluoride

• H2S dihydrogen monosulphide• The “mono” prefix is normally left off of the first element (thus we

say sulphur hexafluoride rather than monosulphur hexafluoride). Sometimes it is even left off the second element (but don’t do that unless you are sure!)

Covalent Formulas are NOT Simplified

• Formulas like N2H4 (systematic name: dinitrogen tetrahydride, common name: hydrazine) are not simplified by cancellation (ie. Don’t write NH2),

• Why? Because the bonding pattern is important in many covalent molecules

A Bad Habit that is Spreading• Some textbooks, including “Quantum Chemistry”,

have a habit of dropping the prefixes from the names of the most common covalent compounds:

• Eg. The book calls SF2 “sulphur fluoride” instead of “sulphur difluoride” and H2S is called “hydrogen sulphide” instead of “dihydrogen sulphide”.

• IUPAC allows the use of some alternate names, but I disagree with this practice.

• Why? Because there is another fluoride of sulphur—SF6, “sulphur hexafluoride” so we should specify carefully!

• If you encounter a covalent compound that is named this way, use the cross-over rule to determine the most likely formula (although this is not the only possible formula)

Another Bad Habit of Mine• I usually use the spelling “Sulfur” (an

Americanization) rather than the accepted Canadian spelling “Sulphur”

• It’s the result of reading way too many American textbooks and seeing far too many periodic tables printed in the U.S.A.

• I will try to remember the correct spelling, but cut me a bit of slack here! Hey, I will forgive you if you misspell sulf..., I mean, sulphur too.

Quick Exercisedo on a sheet of lined paper.

• What is the systematic name of:a. N2H4 d. H2O g. SF6

b. CF4 e. CO h. SBr2

c. N305 f. PO3 i. CS2

What is the formula of:a. Dinitrogen pentoxideb. Sulphur dioxidec. Oxygen difluorided. diphosphorus tetrachloride

a dinitrogen tetrahydride dihydrogen monoxide sulphur hexafluoride

b carbon tetrafluoride carbon monoxide sulphur dibromide

c trinitrogen pentoxide phosphorus trioxide carbon disulfide

N2O5

SO2

OF2

P2Cl4

The Ionic Rules (part 1)Naming Binary Ionic Compound

• Ionic compound names do NOT contain prefixes. They simply start with the metallic elements name, followed by the non-metallic portion.

• Eg:• Na2O sodium oxide

• BaF2 barium fluoride

• NaCl sodium chloride• AlCl3 aluminum chloride

The Formulas of Ionic Compounds• The name of an ionic compound does NOT tell

you its formula.– But, you can always work out the formula from the

name by using the “crossover rule.”• Eg.

– What is the formula of potassium oxide?• Potassium has a valence of 1 (ie. 1+)• Oxygen has a valence of 2 (ie 2-)• Write the valences above and to the right:

– K1 and O2

• Cross them over and simplify. That gives the formula– K2O1 simplify K2O

Ionic Formulas ARE Simplified

• Formulas like Mg4C2 (systematic name: magnesium carbide) ARE usually simplified by cancellation (ie. Mg4C2

becomes Mg2C)

• Why? Because the bonding pattern is not as important between ions as it is in molecules. What is important is the ratio of ions.

Quick Exercise (binary ionic)do on lined paper

• What is the systematic name of each compound?– CaO Na2S CaF2

– Al2O3 MgBr2 Ca3N2

– K3N Li4C BeCl2• What is the formula of:

– Lithium oxide barium nitride– Calcium fluoride sodium phosphide– Aluminum sulphide calcium sulphide– Magnesium carbide lithium chloride

Calcium oxide sodium sulphide calcium fluorideAluminum oxide magnesium bromide calcium nitridePotassium nitride lithium carbide beryllium chloride

Li2O Ba3N2

CaF2 Na3P

Al2S3 CaS

Mg2C LiCl

The Ionic RulesTernary Ionic Compound

• Ternary ionic compounds contain 3 or more elements

• Their systematic names often end in either –ate or –ite.

• They always include a polyatomic ion (also called a radical)

Some polyatomic ions

• OH- hydroxide H- hydride• CO3

2- carbonate HCO32- bicarbonate*

• SO42- sulfate (sulphate) SO3

2- sulfite (sulphite)

• PO43- phosphate PO3

3- phosphite

• ClO31-chlorate ClO2

1-chlorite

• NO31- nitrate NO2

1- nitrite

• NH4+ ammonium CN- cyanide

• See more radicals on page 97of textbook• * A.K.A. hydrogen carbonate (the term “bi” was once used to

indicate a hydrogen atom in the radical)

Finding Formulas from Names

• Binary covalent (molecular) formulas– Use the name as a guide

• Binary ionic formulas– Use the “crossover” rule (or balance charges)

• Ternary ionic formulas– Find the charge of the polyatomic ions.– Then use the “crossover” rule.

• Ternary covalent compounds (organic)– A separate, complex systems of names is

used (not studied in detail this year)

FYI: A few organic-covalent compounds can be known by two or more different names, for example, CH4 can be called carbon tetrahydride (its systematic name) or methane (its organic name). Since methane is shorter, most people use that name. Also NH3 can be called nitrogen trihydride (systematic name) or ammonia (the name used for an NH3 group in organic chemistry)

Quick Exercise (ternary ionic)do on lined paper

• What is the formula of– Calcium carbonate potassium cyanide– Sodium phosphate barium phosphate– Magnesium nitrate aluminum sulfate

• What is the name of– AgNO3 NH4CN– Na2CO3 (NH4)3Cr2O7

– Na2SiO3 KMnO4

– CsHCO3 Al(NO2)3

Ammonium cyanide

Ammonium dichromate

Potassium permanganate

Aluminum nitrite

Silver Nitrate

Sodium carbonate

Sodium silicate

Cesium bicarbonate

CaCO3

Na3PO4

Mg(NO3)2

KCNBa3(PO4)2

Al2(SO4)3

The Organic Rules

• You don’t need to know the organic rules

FYI:

The names of organic compounds are based on the hydrocarbons: methane (1C), ethane(2C), propane(3C), butane(4C), pentane(5C), hexane(6C), heptane(7C), octane(8C), etc. combined with a lot of descriptive endings. Thus an alcohol with one carbon is called methanol, an alcohol with 2 carbons is ethanol, and so on.

Further complicating this is that some traditional names are given to organic compounds that don’t match the system above, so there is a lot of memory work involved in assigning names to organic compounds.

Alkanes: the simplest example of organic compound naming.

• Alkanes are chains of carbon atoms with hydrogen atoms attached, with the general formula CxH2x+2

• CH4 = methane C8H18 = octane

• C2H6= ethane C12H26 = dodecane

• C3H8= propane C13H28 = tridecane

• C4H10=butane C20H42 = eicosane

• C5H12= pentane C23H48 = tricosane

1 2 3 4 5 6 7 8 9 10 11 12 13

meth eth prop but pent hex hept oct non dec Un-dec

Do-dec

Tri-dec

14 15 ... 20 21 22 23 ... 30 31 32 33

Tetra-dec

Penta-dec

Eicos Un-cos

Do-cos

Tricos non dec Un-dec

Do-dec

• For example, the molecule shown below is called:

18-bromo-12-butyl-11-chloro-4,8-diethyl-5-hydroxy-15-methoxytricos-6,13-dien-19-yne-3,9-dione

http://upload.wikimedia.org/wikipedia/commons/5/5b/IUPAC_naming_example_without_carbons.png

Polyvalent elements

• Some elements have more than one possible valence– Transition elements in particular are this way

(but there are a few representative elements that do this too).– For example, copper can be Cu1+ or Cu2+.– In the old days, Cu1+ was called “cuprous

copper” and Cu2+ was called “cupric copper”, but this system proved confusing.

(the old system is called the classical name)– Now, Cu1+ is called Copper(I) and Cu2+ is called

Copper(II), but some chemists still use the old names.

(the new system is called the stock name)– See the chart on page 95 for other examples

Classical Names and the metals they are associated with

metal lower oxidation state higher oxidation state

• Copper cuprous 1+ cupric 2+

• Iron ferrous 2+ ferric 3+

• Mercury mercurous 1+ mercuric 2+

• Lead plumbous 2+ plumbic 4+

• Tin stannous 2+ stannic 4+

• Chromium chromous 2+ chromic 3+

• Manganese manganous 2+ manganic 3+

• Cobalt cobaltous 2+ cobaltic 3+

Copper(I) 1+ Copper(II) 2+

Iron(II) 2+ Iron(III) 3+

Mercury(I) 1+ Mercury(II) 2+

Lead(II) 2+ Lead(IV) 4+

Tin(II) 2+ Tin(IV) 4+

Chromium(II) 2+ Chromium(III) 3+

Manganese(II) 2+ Manganese(III) 3+

Cobalt(II) 2+ Cobalt(III) 3+

Eight questions that can identify which rules to use from a compound’s formula

• Is it a common exception? H2O, H2O2

• Does the formula begin with H?• Does it end with COOH?• Does it end with CH2OH?• Does it end with OH?• Does it end with a radical?• Does it begin with a metal?• Does it begin with NH4?

Org

Ionic

Ionic

Ionic*

Ionic

Ionic*

Covalent (exceptions)

Org

acid

alcohol

base

Organic acid

salt

Ammonium salt

covalent

NO

ternary

YES

YES

YES

YES

YES

YES

YES

YES

You must ask these two question beforeAsking the next one!

Textbook Exercise

• Page 110 #18 to 32

Summary of Rules for Naming Covalent Compounds

• Covalent Compounds contain NO metals.• They don’t start with H or NH4 either.• Names end in “-ide” • Names depend on formulas

– Use prefixes to indicate #atoms of each element• mon(o) =1 penta = 5• di = 2 hexa = 6• tri = 3 hepta =7• tetra = 4 oct(o) = 8

• Examples:carbon dioxide: CO2 carbon monoxide: COcarbon tetrachloride: CCl4 diphosphorus pentoxide: P2O5

sulfur hexafluoride: SF6

Summary of Rules for Binary Ionic Compounds• Binary Ionic Compounds contain a metal

and a non-metal– Names do not contain prefixes– Metal name goes first, followed by non-metal,

with ending changed to “-ide”– If the metal is polyvalent, insert a roman numeral

in brackets to give its valence.– Examples

sodium oxide: Na2O calcium chloride: CaCl2

copper(II)oxide: CuO iron(III)oxide: Fe2O3

-To find a formula from the name, use CROSSOVER.

Name Rules for Ternary Ionic Compounds• Ternary Ionic Compounds always contain a

radical (A.K.A. polyatomic ion)

– Names do not contain prefixes– Name of the metal is followed by the name of the

radical part.• Exception: If the radical is ammonium (NH4) then it goes

first!– Examples:

sodium carbonate: NaCO3

calcium hydroxide: Ca(OH)2

ammonium sulfite: (NH4)2SO3

copper(II)phospate: Cu3(PO4)2

-To find a formula from the name, use CROSSOVER.

The Eight Questions

• Is it water (H2O)? covalent• Does the formula begin with H? ionic• Does it end with COOH? ionic• Does it end with CH2OH? covalent• Does it end with OH? ionic• Does it end with a radical? ionic• Does it begin with a metal? ionic• Does it begin with NH4? ionic• It has only non-metals covalent

Quick Exercise

N2S5 carbon tetrafluoride

SO2 diphosphorus trisulfide

Na2S calcium bromide

Li3N iron(III)sulfide

Ca(NO3)2 iron(II)sulfide

Na2SO4 magnesium hydroxide

Mg3P2 nitrogen trichloride

AlPO4 phosphorus pentoxide

BeCl2 ammonium sulfide

NF3 magnesuium phosphate

Dinitrogen pentasulfideSulfur dioxideSodium sulfideLithium nitrideCalcium nitrateSodium sulfatemagnesium phosphideAluminum phosphateBeryllium chlorideNitrogen trifluoride

CCIIItItIItIC

CF4

P2S3

CaBr2

Fe2S3

FeSMg(OH)2

NCl3PO5

(NH4)2SMg3(PO4)2

Indicate if the compound is Covalent or IonicThen give the systematic name.

Give the formula of each of the compounds named below

Answers to Ex. P.110• 18. Classify the elements:• a) Al = metal b) Ag = metal c) Si = metalloid d) He =

non-metal e) Zn = metal

• 19. Which groups do they belong to?• A) Aluminum = other metals (or group IIIA or boron fam.) • B) Silver = transition metals (or group IB or coin metals)• C) Silicon = metalloids (or group IVA or carbon family)• D) Helium = noble gases (or Inert gases or group VIIIA) • E) Zinc = transition metals (or group IIB)

• 20. State the electrons gained or lost:• a) S2- = gain 2 d) Ba2+ =lost 2• b) K+ = lost 1 e) Li+ = lost 1 • c) Cl- = gain 1 f) H- = gain 1

• 21. Name and identify the ions in Q.20– A) S2- sulphide ion: anion– B) K+ potassium ion: cation– C) chloride ion: anion– D) barium ion: cation– E) lithium ion: cation– F) hydride ion: anion

• 22. Ionic or molecular?– A) CO = molecular* C) C3H8 = molecular*

– B) KBr = ionic D) SO3 = molecular*

* here molecular means the same as covalent, so you could answer “

• 23. Molecular compounds have low melting and boiling points, and are often soft and seldom dissolve in water. Ionic compounds are hard and crystalline and usually soluble in water.

• 24. a) molecule b) formula unit c) molecule d) molecule

• 26. Lithium ion: Li1+, Oxide ion: O2-, Barium ion: Ba2+, Fluoride ion: F1-, Potassium ion: K1+, Neon atom: Ne0

• Note: the digit “1” may be omitted. Neon does not form a stable ion.

• 27. name the ions.• A) H+ hydrogen ion• B) CN- cyanide ion• C) Cr3+ chromium(III) ion or chromic ion• D) Cr2O7 dichromate ion

• E) H2PO4 dihydrogen phosphate ion• F) Sn4+ tin(IV) ion or stannic ion• G) SO3

2- sulfite ion• H) Se2- selenide ion

Older equivalent names are in italics.

• 28. Write the formula and charge• A) Magnesium ion Mg2+

• B) Lead (IV) ion Pb4+

• C) Chromate ion CrO42-

• D) Nitrite ion NO21-

• E) Hydroxide ion OH1-

• F) Iron(II) ion Fe2+

• G) Ammonium ion NH41+

• H) Copper(I) ion Cu1+

• 29. Write the formula from the 2 ions• A) Sr2+, Se 2- SrSe C) Ca2+, N3- Ca3N2

• B) K+, O2- K2O D) Co3+,I- CoI3

• 30. Write the formulas:• A) silver sulfide Ag2S• B) stannic chloride SnCl4

{tin(IV)chloride}• C) sodium nitride Na3N• D) strontium iodide SrI2

• 31. Name the binary compounds:• A) ZnO zinc oxide• B) NaI sodium iodide• C) Cu2O copper(I)oxide (or cuprous oxide)• D) CaBr2 calcium bromide

• 32. When are parenthesis used in a formula?• Parenthesis in a formula can surround

polyatomic ions, but only if there is more than a single polyatomic ion.

• Parenthesis can also be used to surround a roman numeral indicating the charge of a polyvalent metal ion.

Module 1, Lesson 10The Mole and Stoichiometry

• The mole• Avogadro’s number• Volume of a mole of gas• Molar masses• Balanced Equations• Stoichiometry defined• Examples of Stoichiometry

The mass of a mole of any element or compound is called its molar mass.

A Mole is the amount of substance that contains 6.02 x 1023 (Avogadro’s #)

particles of that substance

mass of a substance (g) mass of a substance (g)

molar mass of substance (g/mol)

molar mass of substance (g/mol)

mn= ------

Mn M

mActual mass of substance (g)

# moles

(mol)

Molar mass

(g/mol)

Six Faces of the Mole A mole is six hundred and two sextillion particles

(6.02 x 1023) A mole is a gram-atomic-mass (for elements) A mole is a gram-molecular-mass (for compounds) A mole is 22.4 litres of an ideal gas (at S.T.P.) A mole is the basic unit of amount in chemistry The mole is the basis of the most important

measure of concentration is chemistry (molarity)

Volume of a mole of gas(preview of a coming concept)

• For most substances (ie: solids & liquids) moles are measured by mass, and there is no specific volume associated with a mole.The bigger the particles, the bigger the mole.

• For gases a mole of particles will take up exactly 22.4 L at standard pressure and temperature. Gas particles are very far apart compared to

their size, so the volume of the individual particles is negligible

Molar Mass

• The molar mass (M) is the mass of one mole of a substance.

• Each substance has a unique molar mass which is determined by its formula.

• To find the molar mass, multiply each element’s atomic mass by the number of atoms of the element (given by the formula)

Molar Mass• For example, the molar mass of water (H2O)

– The atomic mass of hydrogen is 1.00794 amu* ≈ 1.01– The atomic mass of oxygen is 15.9994 amu* ≈ 16.0– The molar mass of water is:(2 x 1.01) + (1 x 16.0) = 18.02 g/mol ≈ 18.0 g/mol*

If you prefer, you can use the method below:

H2 1.01 x 2 = 2.02

O1 16.0 x 1 = 16.0

= 18.02 g/mol ≈ 18.0 g/mol*

*For high school chemistry, we can round our masses to 1 decimal place.

Molar Mass• For Example, the molar mass of H2SO4

(2 x 1.01) + (1 x 32.1) + (4 x 16.0) = 98.12 g/mol

≈ 98.1 g/mol

If you prefer, you can use the method below:

H2 1.01 x 2 = 2.02

S1 32.1 x 1 = 32.1

O4 16.0 x 4 = 64.0

= 98.12 g/mol ≈ 98.1 g/mol*

*For high school chemistry, we can round our masses to 3 significant digits.

Balancing Equations

• It is important to be able to balance chemical equations.

• In a balanced equation the atoms of the reactants match the atoms of the products, in agreement with the law of conservation of mass.

• You probably learned to balance most types of equation last year, but we will see a few examples and do some practice ones too.

Balancing Equations

• Example:• Hydrogen + oxygen water skeleton • H2 + O2 H2O unbalanced

• H2 + O2 H2O

Left side Right side

H = 2 H = 2

O = 2 O = 1 2

22

44

Practice

• Balance the equations on page 150 of your textbook (problems 3 & 4)

• Check your own answers in the answer key (p. 723)

Stoichiometry

• Stoichiometry is a method of calculating the amount of a reactant that is needed, or of a product that can be produced in a chemical reaction.

• You must have a balanced equation to do it

• You use the ratio of moles to calculate the missing amounts.

2Mg + O2 2MgO

• Mole Ratio – – describes the ratio in which the substances

combine – Ex. for every 2 mol Mg, 1 mol O2 needed

– used to predict what will happen– Ex. for every 2 mol Mg & 1 mol O2,

2 mol MgO will be produced

2 : 1 : 2

Magnesium metal reacts with oxygen to produce magnesium oxide.

2Mg + O2 2MgO2 : 1 : 2

4 moles of magnesium metal reacts with oxygen to produce magnesium oxide. How many moles of MgO are produced?

Mole Ratio

Write actual amounts on top

4 mol (Mg) ? mol (MgO)=

2 mol (Mg) 2 mol (MgO)

from question

from mole ratio

4 mol (Mg)

? mol (MgO)

Not in Question.

Solve by cross multiplying

4 moles Mg ? moles MgO=

2 moles Mg 2 moles MgO

4 moles MgO are produced

4 x 2 ÷ 2

4Fe + 3O2 2Fe2O3

4 : 3 : 2

Iron metal reacts with oxygen to produce 3 moles of iron (III) oxide. How many moles of oxygen were needed in the reaction?

Mole Ratio

Set up ratio

? mol (O2) 3 mol (Fe2O3)=3 mol (O2) 2 mol (Fe2O3)

from question

from mole ratio

3 mol? mol

NOT ASKED

Solve by cross multiplying

3 mol (O2)

3 mol (Fe2O3)=? mol (O2)

2 mol (Fe2O3)

4.5 moles O2 are needed

3 mol (O2) x 3 mol (Fe2O3) ÷ 2 mol (Fe2O3)

Practice

• Try the problem on page 174 of your text book (problem 2)

• Check your own answer in the answer key

Stoichiometry Problems with Mass

• If the number of moles of substance present is not given, but its mass is, then you have to change the mass into moles, using:

n M

m

4Fe + 3O2 2Fe2O3

4 : 3 : 2

Iron metal reacts with oxygen to produce 80g of iron (III) oxide. How many grams of oxygen were needed in the reaction?

Mole Ratio

Calculate g to mol

? moles O2 0.5 moles Fe2O3=3 moles O2 2 moles Fe2O3

80g = n mol ?

n M

mn= 80 g

159.6 g/mol

≈ 0.5mol

0.5 mol

m

MFe2O3

ScratchWork---Molar mass of Fe2O3

---Fe2 =55.8x2 =111.6O3= 16.0x3 = 48.0

==== 159.6

=

• 3 mol O2 x 0.5 mol Fe2O3 ÷ 2 mol Fe2O3 = 0.75

• 0.75 mol of oxygen were needed. How much is this in grams?

• m = n Moxygen

• M for oxygen ≈ 32.0 g/mol, so• 0.75 mol x 32.0 g/mol = 24.0g

? moles O2 0.5 moles Fe2O3=3 moles O2 2 moles Fe2O3

n M

m

Answer: It takes 24 g of O2 to make 80 g of Fe2O3

Practice

• Page 177 # 3, 4• Check your answers on page 724

Limiting, Excess and Sufficient Reactants (Reagents)

• If there is enough of a reactant to produce the desired amount of product, we say there is “sufficient reactant”

• If there is more than enough of one of the reactants (ie. Too much) we say there is “excess reactant”

• If there is less of one reactant than can react with another, we say that “this is the limiting reactant” (since it limits the amount of product that can be produced)

Example of Limiting and Excess Reactants

• Johann puts 24 g of oxygen gas and 10 grams of hydrogen gas into a reaction chamber and burns them. – A) Which is the limiting reactant– B) Which is the excess reactant– C) How much excess reactant is there?

• Balanced equation: 2 H2 + O2 2 H2O• Next Slide: the Stoichiometry Solution

Johann puts 24 g of oxygen gas and 10 grams (1.0x101g) of hydrogen gas into a reaction chamber and burns them.

• First calculate for the hydrogen and water (ignore oxygen)

• Molar mass (MH2) of Hydrogen is 2.0 x 1.0 g ≈ 2.0 g/mol

• 10 g H2 = 10g ÷ 2.0 g/mol = 5.0 mol H2

• 5.0 x 2 ÷ 2 = 5.0 moles of water can be made.

• Then calculate for the oxygen and water• Molar mass (MO2) of Oxygen is 2 x 16.0 g = 32.0 g/mol

• 24 g O2 = 24g ÷ 32.0 g/mol = 0.75 mol O2

• .75 x 2 ÷ 1 = 1.5 mol of water can be made.

• Oxygen is the limiting reactant• Since less product can be made with the given amount of O2

Next Slide: Calculating Excess

2 H2 + O2 2 H2O2 : 1 : 2

10 g = n mol 24 g = n mol ? mol

Mole ratios underneath:

Problem info on top:

5 = ?

2 2

0.75 = ?

1 2

5.0 .755 mol with H2

1.5 mol with O2

IgnoreFor Now

Calculating excess

• Hydrogen will be the excess reactant• How much excess?

• Since we can produce 1.5 mol of H20 from our oxygen, we can only use 1.5 x 2 ÷ 2 = 1.5 mol of hydrogen, or 3 g. Since we started with 10 grams of hydrogen, we will have:

10g - 3g = 7 g left over.

2 H2 + O2 2 H2O2 : 1 : 2

n mol = 24 g = 0.75 mol 1.5 mol

Next Slide: Percent Yield

Percent Yield

• Stoichiometry assumes that all reactions go perfectly, and will give an exact amount of product. In the real world, there is usually a slight discrepancy between the expected amount of product (theoretical yield) and the actual yield.

• The Percent yield is calculated as:

%100YieldlTheoretica

YieldActualYieldPercent

Next Slide: Assignments

Six aspects of a Mole• A mole is six hundred and two sextillion

particles (6.02 x 1023)• 602 000 000 000 000 000 000 000 particles

• A mole is a gram-atomic-mass (for elements) • Atomic mass in grams

• A mole is a gram-molecular-mass (compounds)• Molar mass in grams/mol

• A mole is 22.4 litres of an ideal gas (at STP)– Standard temperature and pressure

• A mole is the basic unit of amount in chemistry• The mole is the basis of the most important

measure of concentration is chemistry (molarity)– One mole of substance dissolved in a litre of water

nMmMmn

n M

mActual mass of substance (g)

# moles

(mol)

Molar mass

(g/mol)

The mole formula is used to find the number of moles if you know the actual mass and the molar mass of a substance. It can also be rearranged to solve other problems using these quantities

n

mM

Re-arrange

Sufficient, Limiting, and Excess Reactants• Sufficient Reactants means you have

exactly enough of all reactants to complete the reaction.

• Limiting Reactant means that you will run out of this reactant, and then the reaction will end (ie. it limits the reaction).

• Excess Reactant means that you will have some left over when the reaction ends (you may have to calculate how much will be left)

Percentage Yield

%100YieldlTheoretica

YieldActualYieldPercent

Stoichiometry Setup

Actual moles 1 Actual moles 2

=Molar Ratio 1 Molar Ratio 2

Skills to Practice

• Balancing equations• Finding molar masses• Solving stoichiometry problems

– These are skills that must be practiced until you are comfortable with them.

Assignment

• Read Chapter 8 “Stoichiometry”– Concentrate on pp. 169 to 184. This is what

you need to complete the exercises.– The “Energy Changes” and “Heat of Reaction”

sections will be used later in the course. You may skim them briefly now, we will review them in detail later

• Do the following Questions from pp191-193– #17, 18, 19, 20, 21, 22*, 23*, 24*, 27

* These 3 are a set. You may want to do 22a, 23a, 24a then 22b, 23b, 24b …

Answers to Exercises, p191

• 17.

5C + 2SO2 CS2 + 4CO

– A) 1.3 mol of CS2 calculation:((6.3 x 1) / 5)

– B) 18.1 mol of C calculation:((7.4 x 5) / 2)

– C) 3.05 mol of CO calculation:((0.762 x 4) / 1)

– D) 364 mol of SO2 calculation:((182 x 2) / 1)

5 : 2 : 1 : 4

6.3 mol

2

2

15

3.6

CS

CS

C

C

mol

xmol

mol

mol

• 18. Note: M(CH3OH) = 32g/mol, M(CO) =28g/mol

– A) n= 600g÷32g/mol = 18.75mol, so we will need 18.75 mol CO and 37.5 mol H2

– B) n=10 mol, so we need 10 mol CO and 20 mol H2, that is: 280g CO and 40g H2

– C) n=5.74, so we need 11.48 mol H2 or: 22.96 g

• 19. MNa202 ≈ 78.0 g/mol, MH2O ≈ 18.0 g/mol,

MNaOH ≈ 40.0 g/mol MO2 ≈ 32.0 g/mol

– A) how many grams of Na2O2 are needed to make 3.20g of oxygen?

• n= 3.20g / 32g/mol = 0.10mol, so we need 0.20mol Na2O2, or about: 0.20 mol x 78g/mol = 15.6g Na2O2

– B) how many grams of NaOH are produced?

• n=0.1 mol, so we need 0.4 mol NaOH, or about: 0.40mol x 40g/mol = 16g NaOH.

– C) When 0.480g of Na2O2 are dropped in water how much oxygen?

• n=0.480g/78 g/mol ≈ 0.006 mol, so we can produce 0.003 mol O2 or about 0.003mol x 32g/mol = 0.096g O2

2Na2O2 + 2H2O O2 + 4NaOH

• 20. The molar masses of reactants are:• M (Li3N)= 34.8g/mol M (H2O) = 18.0g/mol

• M (NH3) = 17.0g/mol M (LiOH) = 23.9g/mol

A) n=98.7g/34.8g/mol= 2.84 molLi3N so we need 8.53 molH2O or

8.53mol x18.0g/mol = 154g of H2O

B) 2.84mol (6.02 x 10 23) = 1.71 x 10 24molecules

C) n= 45L/22.4L=2.01 mol x34.8 ≈ 69.9g of Li3NMolar masses are accurate to 3 significant digits. If different accuracies are used for molar masses then answers may range from 150g to 155g, 1.69x1024 to 1.73x1024 and 69.7g to 70g.

• 21 The limiting reagent is the reactant that will produce the least amount of product.

• 22 limiting 23. product 24. excess– A) O2 4.6 mol H2O 0.4 mol H2

– B) O2 3.2 mol P2O5 0.6 mol P

– C) Cl2 5.33 mol AlCl3 0.667 mol Al

– D) P2O5 0.74 mol H3PO4 1.4 mol H2O

• 27: Fe2O3 + 3 CO 2Fe + 3CO2

1 3 2 3– Molar mass of Fe2O3: MFe2O3= 159.6 ≈160. g/mol,– So, 84.8g Fe2O3 =84.8g/160.g/mol = 0.531 mol Fe2O3

– Which produces: 0.531 x 2 /1 = 1.062 molFe

– Or 1.06mol x 55.8g/mol = 59.1gFe calculated theoretical yield

– Actual yield = 57.8gFe, from the question.

– so (57.8g ÷ 59.1g) x 100 = 97.8% yield

(actual) (theoretical) (%)

84.8g =0.53mol X mol =?g

Note: the calculations in this problem are rounded to 3 significant digits during computation. If they are not rounded, the result will be slightly different. (97.47%)

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