Pages 3 to 33 “Quantum Chemistry” Target Completion Date: October 1

Pages 3 to 33“Quantum Chemistry”Target Completion Date: October 1

ChemistryReview Section

About Slide IconsVery Important Points

• You should either note or highlight items from this slide. Some items from this slide WILL be on tests!

Important Sample Problems• Always hand-copy important sample problems in your note book,

and refer back to them when doing assignments. Similar problems will be on tests!

Look at this! (usually charts, diagrams or tables)• You don’t need to copy this, but you must read and understand the

diagrams or explanations here. Concepts will be tested, but not the details.

Information only. Don’t copy!• This is usually background information to make a topic more

interesting or to fill in details, or to give examples of how to use a table. Not directly tested.

Review Stuff• Not part of the material you will be tested on, but you are expected

to remember this from grade 10. It may be indirectly tested.

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Pages with a PINK background are

supplementary . Not material for a test!

Conversions

• You must be able to do ALL standard metric conversions, especially:– Litres to millilitres, millilitres to litres– Grams to kilograms, kilograms to grams

• Other conversions you will learn during the course of the year:– Temperature: degrees celcius (℃) to kelvins (K)– Pressure: kilopascals (kPa) to millimetres (mmHg)

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Quick ConversionsPrefix means

mega (M) million

… 100000

… 10000

kilo (k) 1000

hecto (h) 100

deca (da) 10

… unit 1

deci (d) 0.1

centi (c ) 0.01

milli (m) 0.001

… 0.0001

… x10-5

micro (μ) X10-6

The table on the left gives the eight most commonly used prefixes in the metric system.

It also includes five rows that do not have prefixes.

The middle row is for the unit: metre, litre, gram, newton, or any other legal metric unit.

This table can be used to quickly convert from one metric amount to an equivalent. Make a copy of this table on the margin of the front cover of your notebook, and learn how to use it. Lets do an example. Let’s find how many centimetres there are in 2.524 kmConversion: 2.524 km ? cm

2 524 km 00 cm Add extra zeros if necessary

There are five steps in the table between “kilo” and “centi”, so we have to move the decimal five places to the right. If we were going up the table we would move left. Answer: 2524 km = 252 400 cm

Density• Density is the relationship between the volume of an object and

its mass. Density is an important characteristic property of matter.

• This is a review formula from last year:

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Where: ρ = the density of the object, in g/cm3 or g/mLm = the mass of the object, in gV = the volume of the object, in cm3 or mL

ρ Vmor

ρw = 1 g/mL = 1 g/cm3 The density of water is 1 g/mL. This is not true of other substances. Objects with less density than water will float. Objects with greater density will sink.

Solving Problems• When solving Chemistry problems on a test or exam,

it is important not only to find the correct answer, but to justify it. While solving the problem you should:1. Show your data, the information you used to solve the

problem.2. Show your work, including the formulas you used and

the substitutions you made.3. Write an answer statement, a sentence that clearly

states your final answer.4. Include the correct units for your answer. Never just give

a number—you must specify what the number means!

Suggested Solution MethodProblem: A block of material has a length of 12.0 cm, a width of 5.0 cm,

and a height of 2.0 cm. Its mass is 50.0 g. Find its density.Arrange your solution like this:

List all the information you find in the problem, complete with units, and the symbols.

Data:l = 12 cm.w = 5.0 cmh =2.0 cmm =50gV = ?

To Find:ρ (density)

Write down all the formulas you intend to use: Formulas: V = lwhρ =Show the substitutions you make, and enough of your calculations to justify your solution:

V =12cm x 5cm x 2 cm= 120 cm3𝜌 = 50 g / 120 cm3

= 0.46 g/cm3

Answer: The density of the block is 0.46 g/cm3 (or 0.46 g/mL)

Always state your answer in a complete sentence, with appropriate units.

Problems on Conversions and Density

1. Convert the following:a) 125 mL to L d) 30 mL to L g) 75 mL to Lb) 450 g to kg e) 4500 mL to L h) 0.035L to mLc) 2.5 L to mL f) 1.35 kg to g i) 0.56L to mL

2. Find the density of a 4cm x 3cm x 2cm block that has a mass or 480 g. Justify your solution.

3. Find the width of a cube whose density is 5 g/cm3 and whose mass is 135 g. Justify your solution.

Also: Do the worksheets entitled “Density” and “Metric Conversions”

Overview: Significant FiguresKnowing how much to round an answer.

In the sciences, we have an particular way of determining how much precision we need in the observations and answers we record. The method of rounding is called significant digits or significant figures. There is a detailed section in the appendix to your textbook on pages 394 to 397. Unfortunately, a few of the details given there are, well… I won’t say wrong, let’s just call them “uncertain”.

UncertaintyIn math, numbers are considered pure, abstract things. In math, 2.00, 2.0 and 2 are considered the same, they all represent perfect number 2. In science, numbers are considered to be measurements, and all measurements have some degree of uncertainty. They are almost never considered perfect!The absolute uncertainty of a measurement is usually ½ of a measuring instruments smallest gradation. If a graduated cylinder is marked in millilitres, then each measurement taken with that cylinder has a ±0.5 mL uncertainty.In science, 2 mL, 2.0 mL and 2.00 mL are different!

Correct precision

• It is considered improper in science to imply that a measurement is more precise than it really is.

• If you have a graduated cylinder that is marked in 1 mL increments, you can record it to between the two smallest marks: eg. 20.0 ±0.5 mL or 25.5 ±0.5 mL are acceptable readings. • With the same graduated cylinder, it would be wrong to

write 20 ±0.5 mL or 25 ±0.5 mL or even 20.00 ±0.5 mL

• In science 20 mL, 20.0 mL and 20.00 mL have different meanings with respect to precision.

Rules for Significant Figures Interpreting Significant Digits

1. All non-zero digits are ALWAYS significant2. Zeros between significant digits are ALWAYS

significant.3. Zeros at the beginning of a number are NEVER

significant.4. Zeros at the end of a number MAY be

significant.5. Exponents, multiples, signs, absolute errors

etc. are NEVER significant.

Examples of Rule 1, 2 and 3

Rule 1. Non-zero digits are ALWAYS significant.1.234 has 4 significant digits145 has 3 significant digits19567.2 has 6 significant digitsRule 2. Zeros between significant digits ARE significant.1001 has 4 significant digits5007.4 has 5 significant digits20000.6 has 6 significant digitsRule 3. Zeros at the beginning are NEVER significant.007 has 1 significant digit0.0000005 has 1 significant digit0.025 has 2 significant digits

Explaining Rule 4Rule 4. Zeros at the end of a number MAY be significant.Your textbook says that they are ALWAYS significant, but this is contrary to what most textbooks say. If there is a decimal point, there is no problem. All textbooks agree, the zeros are ALL significant.

3.00000 has 6 significant digits5.10 has 3 significant digits10.00 has 4 significant digitsIf there is NO decimal, the situation is ambiguous, and a bit of a JUDGEMENT CALL. If you trust the source to be precise, then you count all the zeros at the end. If you have reason to believe the person was estimating, then you don’t count any of the zeros at the end.

5000 has 1 or 4 significant digits250 has 2 or 3 significant figures123 000 000 has 3 or 9 significant figures

In a test situation, assume the numbers are precise, unless something in the question states otherwise.

Estimated source Trusted precise source

Rule 5Rule 5: Exponents and their bases, perfect multiples, uncertainties (error values), signs etc. are NEVER significant.6.02 x 1023 has 3 significant digits504.1 mL x 3 has 4 significant digits5.3 ±0.5 mL has 2 significant digits–5.432 x 10-5 has 4 significant digitsIn each case, the blue part is significant, the green part is NOT significant.

Note: The term Significance in this usage is not the same as importance. A digit may be “insignificant” but still very important. The significant digits guide you to the correct way of rounding numbers to show precision. The insignificant digits may serve as “placeholders”, making sure the decimal point is in the right place. An important job, but not one that adds to the precision of the answer.

Same Number, Different PrecisionNumber Precise to

200.000 6 significant digits200.00 5 significant digits200.0 4 significant digits200* Ambiguous* 1 to 3 SD2.00 x 102 3 significant digits2.0 x 102 2 significant digits2 x 102 1 significant digit

*Your textbook says to call this 3 significant figures. Traditional measurement would call it 1 significant figure. Written this way it is ambiguous. Avoid writing answers that end with zeros and no decimal!

Try to avoid the “ambiguous” situation in your answers. If an answer ends in zero, or worse, in several zeros, indicate whether it should be interpreted as “exactly” or “approximately”.

Better still, convert it to scientific notation, and leave only the zeros you know are accurate.

Eg. If your answer is 2500 mL, but you only measured to the nearest 10mL, then write 2.50 x 103 mL. That way every one will know its accurate to 3 significant figures

Math with Significant Figures

• The basic rule for math is that you do not improve significance by multiplying or dividing numbers:53.81 m x 2.43 m = 131 m2 NOT 130.7583 m2 !!!Why? Because the least precise measurement had 3

significant digits, so our answer should not have more than 3 significant digits!

The technique for addition and subtraction is slightly different (see p.396 ) but the concept is the same. You cannot make your result better than your measurements!

Topic 1: Organization of Matter

• 0.1.1 Atoms and Molecules– All matter is composed of atoms.– The atoms that make up most matter are

assembled into molecules. • A molecule may contain one atom, or it may contain

several thousand atoms, or any number between.

– A molecule is represented by its formula• Water molecules, for example, are represented by the

formula H2O, shown below:

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H2O2 atoms of hydrogen

1 atom of oxygen

0.1.1

18

R

H

CO

N

SCl

O

O

H

H

H

Cl

NH3

CO2

SCl2H

Ne

One atom

several thousand atoms

Ne

DNA

Na Cl

• Chemical Formulas and Ions– Some matter is formed from ions instead of

normal atoms or molecules.• For the most part, we treat ions the same way as

regular atoms, but there are a few very technical differences.

– Ions are atoms (or clusters of atoms) that have become positively or negatively charged by losing or gaining one or more electrons.• Positive ions are called cations (ca+ions), • Negative ions are called anions (aNions)• Metals almost always form cations (+), non-metals

may form anions (-)

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Notice the slightly

stronger wording with

respect to metals than nonmetals!

Na+ Cl–cation anion0.1.2

19

Differences between ionic and covalent compounds

Ionic Compounds Covalent (molecular) Compounds

Ionic bonds “give” or “take” electrons Covalent bonds “share” electrons

Ionic compounds don’t have distinct molecules. Clusters of ions are sometimes referred to as “formula units” rather than “molecules”.

Covalent compounds have distinct, strongly bonded molecules. This is why some people call covalent compounds “molecular” compounds.

Most ionic compounds are solid at room temperature.

Covalent compounds may be solid, liquid or gas at room temperature.

Ionic compounds usually have a high melting point. That’s why they are solid.

Covalent solids usually have a low melting point.

Ionic solids are usually hard, but brittle Covalent solids are usually softer

Ionic compounds are usually more soluble in water, but less soluble in non-polar solvents like acetone.

Covalent solids are usually less soluble in water, but more likely to dissolve in non-polar solvents like acetone.

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Sample IonsElement ions Ion names alternate names

Sodium Na Na+ Sodium ion

Calcium Ca Ca2+ Calcium ion

Aluminum Al Al3+ Aluminum ion

Tin Sn Sn4+,Sn2+ Tin(IV) ion, Tin(II) ion Stannous, Stannic

Copper Cu Cu2+, Cu+ Copper(II) ion, Copper(I) ion Cuprous, Cupric

Iron Fe Fe3+, Fe2+ Iron(III) ion, Iron(II) ions Ferrous, Ferric

Carbon C C2+, C4+, C4- Carbon(II), Carbon(IV), Carbide

Carbon can form both anions and cations as well as covalent bonds

Nitrogen N N3- Nitride ion,

Phosphorus

P P3- Phosphide ion,

Oxygen O O2- Oxide ion,

Sulphur S S2- Sulphide ion Sulfide ion

Fluorine F F1- Fluoride ion

Chlorine Cl Cl1- Chloride ionNotice that some elements can form more than one type of ion. Compounds of the same element can differ quite a bit, for example, red iron oxide (rust) has Fe3+ ions, black iron oxide (wustite) contains Fe2+ ions. Note also, that most negative ions have the name ending changed to –ide.

Meta

l Io

ns

(+)

Non-M

eta

l Io

ns

(-)

+ Ca

tions

– An

ions

21

H

H

O OO

O

Big Fat Ions(Polyatomic Ions)

• Polyatomic ions are ions that are composed of a cluster of atoms, instead of a single atom.

• For example, the nitrate ion (NO3–) looks like

this:• But it acts like a single, negatively charged

particle in reactions.• Polyatomic ions are sometimes called radicals.• They are not the same as molecules.

O

O

ON-

SO

O

PO

O

N HH+

2–

3–Cl O

–

O

O

ON- Na+ + NO3- NaNO3

Na+

Common Polyatomic Ions (p.422)Formula Name (ionic charge) Formula Name (ionic charge)

PO4 3- Phosphate ion (3-) NO3 - Nitrate (1-)

PO3 3- Phosphite ion (3-) NO2- Nitrite (1-)

SO4 2- Sulphate ion (2-) ClO4 - Perchlorate (1-)

SO3 2- Sulphite ion (2-) ClO3 - Chlorate (1-)

CO3 2- Carbonate (2-) ClO2 - Chlorite (1-)

CrO4 2- Chromate (2-) ClO - Hypochlorite (1-)

C2O4 2- Oxalate (2-) MnO4 - Permanganate (1-)

SiO3 2- Silicate (2-) H2PO4 - Dihydrogen phosphate (1-)

HPO4 2- Hydrogen phosphate (2-) . H2PO3- Dihydrogen phosphite (1-)

HPO3 2- Hydrogen phosphite (2-) HSO4- Hydrogen sulphate (AKA: bisulphate) (1-)

Cr2O7 2- Dichromate (2-) HSO3 - Hydrogen sulphite (AKA: busulphite) (1-)

C2H3O2 - Acetate (AKA: ethanoate) (1-) HCO3 - Hydrogen carbonate (AKA: bicarbonate) (1-)

OH- Hydroxide (1-) NH4 + Ammonium (1+)

CN - cyanide (1-) H3O+ Hydronium (also written as H+) (1+)

3-

2-

1-

1-

23This information is important when naming ternary ionic compounds. Click to skip ahead to Ionic Naming Rules

Representation of Atoms

• Early Representations– Democritus (c.450 BC) suggested that matter

was made of particles.– John Dalton (1800) represented the atoms as

spheres (like microscopic bowling balls)– J.J. Thomson represented the atom as a “plum

pudding” of positive charge with negative charged electrons scattered inside “like rasins”

– You studied the historic importance of these models last year, so you will not be tested on them this year. We will concentrate on the three most widely used representations on the slides that follow.

H

C

O

N

+-

- -

-

-

0.2.0

P

S

Cl

24

Dalton modelsOriginal

and Modern

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1. Rutherford-Bohr Model– Rutherford discovered that the atom has

a dense nucleus containing positively charged protons.

– Negatively charged electrons move around this nucleus in paths that resemble an orbit.

– Later, Bohr calculated that there were different orbital energy levels or “shells” that could hold different numbers of electrons.

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Early Rutherford model

Revised Bohr model

0.2.1

25

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2. The Simplified Atomic Model– The simplified atomic model that we often use

today adds neutrons (discovered by James Chadwick after the Bohr-Rutherford model had been proposed) to the protons in the nucleus.

– We often draw this in a simplified way, showing the nucleus as a full circle, and the electron “shells” as half-circles.

Page 5

11p+

12n0Na2e- 8e- 1e-

Z=11, configuration: 2,8,1

Nucleus: If asked for a complete simplified model, give the #protons and #neutrons (if known) in the nucleus. Otherwise, just draw a full circle.

The Atomic Number, Z, is the number of protons in the element. The configuration is the arrangement of the electrons in the shells

Symbol: The symbol of the element Electrons: 2 in first shell, 8 in 2nd 1 in 3rd

0.2.2

26

WARNING

• Be careful how you draw them!• The diagram must show the nucleus!

Unacceptable!

Nucleus is not shown. Nucleus is confused with 1st shell

Nucleus shown as solid circle.Labelled with element symbol beside.

Nucleus shown as full circle.Labelled with #protons and neutrons.

Unacceptable!

ACCEPTABLE ACCEPTABLE

The Sub-atomic ParticlesParticle Symbol Charge Actual Mass

(g) Rounded mass

(amu)Location in atom

Proton p+ 1+ 1.672x10 -24 ≈1 u (1679/1680) Nucleus

Neutron n0 0 1.674x10 -24 ≈1 u (1680/1680) Nucleus

Electron e- 1- 9.109x10 -28 ≈0 u (1/1680) Shells

nucleons

Page 6

28

3. Lewis Model: (AKA Lewis electron dot notation)– Lewis notation is a way of drawing a representation

of the valence electrons of an atom– When sketching an atom, write the symbol, and

then arrange dots around it to represent its valence electrons.

– Example: N has 5 valence electrons N– The “odd” or unpaired electrons are available for

the purpose of bonding. – When bonding, atoms gain, lose or share electrons

in order to get a total of 8 electrons around each atom.

1

2

3

4

5 2 paired electrons

3 “odd” unpaired electrons

0.2.3

29

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The preferred way of drawing Lewis diagrams of the first ten elements is shown below:

However, the dots may be moved around to show different arrangements. All of the drawings of Beryllium shown below might be correct in some circumstances.

Sometimes showing the bonding between atoms requires clever movement of dots, as in the drawing of a nitrogen molecule (N2) shown here:

30

Sometimes electrons are removed from one atom to others in order to get 8

The Modern Model(Optional Enrichment)

• The Modern Model of the Atom– Of course, the Rutherford-Bohr model and the

Simplified Model do not perfectly represent what happens inside the atom. No model can!

– A more complete model, The Modern or Electron-Cloud model exists, but is more complicated and extremely difficult to draw.

– The Modern Model more accurately explains the relationship between the atom and the periodic table, and allows you to produce simplified models of elements in the transition area of the periodic table.

31


• The 2-8-8 vs. 2-8-18 problem.– You have probably been taught how to draw

Simplified Models for the first 20 elements – If so, you have noticed that for the elements

potassium and calcium, the third shell only holds 8 electrons—but Bohr said it should hold up to 18!

– The models you have been taught can’t explain why, but the modern model includes a concept called “orbitals” or subshells, and a filling pattern called the “aufbau diagram” that explain this .

32


• You are not required to learn the Aufbau diagram or the modern electron cloud model, but if time permits, I will show you how it works near the end of the review section. In the meantime:– You must know that the third shell CAN hold up to

18 electrons, but often doesn’t.– And you must learn how the periodic table can be

used to figure out the electron arrangement of many elements past the first 20. • But that is part of the next lesson…

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Atomic Model Exercises

1. Draw Simplified Models of the first 20 elements.2. Draw Lewis Models of the first 20 elements.3. Convert the following:

a) 125 cm to m d) 320 mL to L g) 750 mL to Lb) 280 g to kg e) 45000 mm to km h) 0.0035km to cmc) 4.63 L to mL f) 5.52 kg to g i) 0.45L to mL

Periodic Classification

OverviewThe periodic table is a useful arrangement of the elements, into regions, families and periods that have important meanings. It is also a source of much additional information about the elements. With careful interpretation of the table, we can find the number of protons an atom has, the approximate number of neutrons, and the arrangement of electrons in the atom and in its ions.

In-line Notation of Element Information

• An alternative to the periodic table is in-line notation of elements and isotopes. Note that the arrangement of information in this notation system is not the same as the arrangement in most periodic tables.

• Examples of inline notation:

• In-line notation is designed to be more compact, but less complete presentation of the information in a full periodic table.

In-Line Notation for a Carbon-14 atom(carbon-14 is an isotope, or alternate form of normal carbon)

C14

6

4–

2

Isotope or Mass number. Represents

the number of nucleons in a

particular atom

Atomic number “Z” represents the

number of protons in this atom

Subtracting the Mass # and the Atomic #

“Z” gives the number of neutrons in the

atom

8

Valence (4)or

Ionic Charge (4–)or

Oxidation # (–4)

Number of atoms in a molecule, such as

C2H4

Valence is the number of bonds the atom is likely

to form. Ionic charge is the most

likely charge an ion will have.

An average carbon atom weighs 12.01 amu according to the periodic table. But no atom

of carbon has that exact weight. For every thousand atoms that weigh exactly 12 amu, a few weigh more. This one weighs 14 amu

(6p+, 8n0, 10e-)

Configuration of this atom

Information in your Periodic Table

O8 2- Ionic charge

Atomic Radius

Density (g/L gas)

(g/mL solid/liquid)

Oxygen

15.999

Name

Atomic weight (amu)

0.65

1.43

3.44

1314

-218.3-182.9

Atomic number (Z)

Electronegativity

Ionization EnergyMelting Point (°C)

Boiling Point (°C)

(or g/mol)

Symbol

The symbol is a 1 or 2 letter abbreviation of the element’s name, or sometimes its Latin name. The first letter is always uppercase. If there is a second letter it MUST be written in lowercase. (eg. For sodium, Na is correct, na or NA are absolutely unacceptable!)

The number of protons

The English name of the element

Also the molar mass in g/mol

Electronegativity is a rating of how well the atom attracts electrons, on a scale from 0 to 4

Ionization Energy is how much energy it takes to remove an electron (kj/mol)

38

The Periodic Tablewith Regions shaded

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

dGas VIII

1 H II Synthetic III IV V VI VII He

2 Li Be metal Metal-oid

Non-metal B C N O F Ne

3 Na Mg

III B

IV B

V B

VI B

VII B

VIIIB

IB

IIB

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 Db Sg Bh Hs Mt Ds Rg Cn Uut Uuq Uup Uuh Uus Uuo

6 La Ce Pr Nd Pm

Sm

Eu Gd

Tb Dy Ho Er Tm

Yb Lu

7 Ac Th Pa U Np Pu Am

Cm

Bk Cf Es Fm

Md

No Lr

↑ The properties and region associations of these 10 elements are hypothetical ↑

39The heavy “staircase” line was the traditional separation between metals & non-metals but we now know it is not a sharp division.

The Periodic Tablewith Families Shaded

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18I III IV V VI VII VIII

1 H II Transition Elements He

2 Li Be B C N O F Ne

3 Na Mg

III B

IV B

V B

VI B

VII B

VIIIB

IB

IIB

Al Si P S Cl Ar



5 Rb Sr Y Zr Nb Mo




6 La Ce Pr Nd Pm

Sm

Eu Gd

Tb Dy Ho Er Tm

Yb Lu


Cm

Bk Cf Es Fm

Md

No Lr

↑ The properties and family associations of most elements in period 7 are hypothetical↑

IA:

Alka

li M

etal

s

IIA

: Al

kalin

e Ea

rths

VIIIA

: N

oble

Gas

es

VIIA

: H

alog

ens

VI: O

xyge

n Fa

mily

V: N

itrog

en F

amily

IVA:

Car

bon

Fam

ily

IIIA:

Bor

on F

amily

Lanthanides

Actinides

IB:

Coi

n M

etal

s

Iron Triad

40

The Periodic Tableand Valence Electrons (electrons in outermost shell)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18I III IV V VI VII VIII

1 H II Transition Elements He


3 Na Mg

III B

IV B

V B

VI B

VII B

VIIIB

IB

IIB

Al Si P S Cl Ar



5 Rb Sr Y Zr Nb Mo




6 La Ce Pr Nd Pm

Sm

Eu Gd

Tb Dy Ho Er Tm

Yb Lu


Cm

Bk Cf Es Fm

Md

No Lr

↑ The properties and family associations of these synthetic elements are hypothetical ↑

ON

E (

I)

TW

O (

II)

EIG

HT

(VIII

)

SE

VEN

(V

II)

SIX

(V

I)

FIV

E (

V)

FOU

R (

IV)

TH

REE

(III)

41

TH

REE

(III)

SE

VEN

(V

II)

SIX

(V

I)

FIV

E (

V)

FOU

R (

IV)

TWO

ON

E

If the square is the same colour as the arrow above, it obeys its family with respect to valence. If the square is rainbow shaded, it is polyvalent, and not obeying its family rules. If the square is partly shaded, then it obeys its family rules most of the time.

The Periodic Table with Periods shaded

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

1 H II III IV V VI VII He


3 Na Mg

Al Si P S Cl Ar



5 Rb Sr Y Zr Nb Mo




6 La Ce Pr Nd Pm

Sm

Eu Gd

Tb Dy Ho Er Tm

Yb Lu


Cm

Bk Cf Es Fm

Md

No Lr

↑ The properties and family associations of these 10 elements are hypothetical ↑

3rd Period = 3 shells

5th Period = 5 shells





1st Period = 1 shells

2nd Period = 2 shells


42The periods of the table show how many shells of electrons an element normally has.

Use the Periodic Table to Find the Electron Arrangement of an Atom

Eg. Find the electron arrangement of Iodine (I)

Iodine is at the intersection of Period 5 and Family VII. Its number is 53. It has a total of five shells, 7 electrons in the outermost shell, and will have 53p+, and normally 53 e-. From this we can USUALLY figure out the electron arrangement.

53p+

2 7e-8 18 Total 53, So far: 35, left: 1818

1 H II A

III A

IVA

VA

VIA

VIIIA

He


3 Na Mg

III B

IV B

V B

VI B

VII B

VIIIB

IB

IIB

Al Si P S Cl Ar



5 Rb Sr Y Zr Nb Mo



5th Period = 5 shells 53

Five shells

SE

VEN

(V

II)

Periodic Table Exercises

• Write the name and symbol of each of the first 20 elements. (bragging rights if you can do it without looking!)

Naming Compounds

• There are four sets of rules for naming compounds:– The binary ionic rules:

• For compounds containing only two elements, joined by an ionic bond.

– The ternary ionic rules:• For compounds containing 3 or more elements, including a

polyatomic ion.

– The covalent rules:• For two elements joined by covalent bonds (usually two non-

metals)

– The organic rules:• Used for compounds that contain carbon atoms bonded to

each other covalently.45

The Binary Ionic Rules

– First name the element on the left side of the compound’s formula.

– Then name the element on the right hand side of the compound’s formula, but change the suffix to “ide”

• For example:NaCl sodium chloride BaCl2 barium chloride

CaO calcium oxide K2S potassium sulphide

Al2O3 aluminum oxide Ca2C calcium carbide

Na+ Cl-

Ca2+ O2-

Ba2+Cl- Cl-

Al AlO2-

O2-

O2-

S2- K+K+

Ca2+ Ca2+C4-

1+ 1–

3+3+

46

Non-metal ion endingsElement (symbol) Negative Ion (charge) Element (formula) Negative Ion (charge)

Boron (B) Boride (B5-) Phosphorus (P4)

Phosphide (P3-)

Carbon (C) Carbide (C4-) Sulphur (S8) Sulphide (S2-)

Silicon (Si) Silicide (Si4-) Fluorine (F2) Fluoride (F–)

Arsenic (As) Arsenide (As3-) Hydrogen (H2) Hydride (H–)

Selenium (Se)

Selenide (Se2-) Chlorine (Cl2) Chloride (Cl–)

Other monatomic negative ions occur rarely. If you encounter one, use the atomic name, with the last syllable altered to ide as sounds best. Eg. Antinide or Polonide

Bromine (Br2) Bromide (Br–)

Iodine (I2) Iodide (I–)

Nitrogen (N2) Nitride (N3-)

Oxygen (O2) Oxide (O2-)

47

Ionic Rules No No!• When naming an ionic compound (and that

includes most compounds that contain a metal)

YOU SHOULD NOT USE A PREFIX!• Do NOT say: calcium difluoride for CaF2

• It’s Wrong. The correct name is just calcium fluoride.

• Do NOT say: dialuminum trioxide for Al2O3 • It’s Wrong. The correct name is aluminum oxide.

There are, or rather there USED to be, a few exceptions to this. Chromium dioxide was an acceptable name for CrO2, and is still used occasionally. Now the name chromium(IV)oxide is preferred for the compound, since it obeys the ionic rules. Monosodium glutamate is an organic compound that does not follow the rules.

Dealing with Polyvalent Metals

• Some metal elements have more than one possible valence. Copper, for example, can have a valence of 1+ or 2+, depending on what compound it is in (eg. CuCl or CuCl2). Since we don’t use prefixes in naming ionic compounds, we shouldn’t use copper dichloride. We need a new rule!– If a metal is polyvalent, we include its current valence in

roman numerals inside parenthesis within an ionic compound name, for example:

– CuCl = Copper (I) chloride (not copper monochloride)

– CuCl2 = Copper (II) chloride (not copper dichloride!)This copper ion must have a charge of 2+

This copper ion has

a charge of 1+

49

Polyvalent ElementsThe elements with flashing circles have more than one positive valence.

1+ 2+ 3+ 4+ 5+ 6+ 7+ 4- 3- 2- 1- 0I VIII

1 H II III IV V VI VII He

2 Li Be Non-metal B C N O F Ne

3 Na Mg

III B

IV B

V B

VI B

VII B

VIIIB

IB

IIB

Al Si P S Cl Ar



5 Rb Sr Y Zr Nb Mo



7 Fr Ra U n c e r t a i n

6 La Ce Pr Nd Pm

Sm

Eu Gd

Tb Dy Ho Er Tm

Yb Lu


Cm

Bk Cf Es Fm

Md

No Lr

FeCr

Sn

Pb

Sm Eu

Cu

Au

Mn Co

Hg

Ru Pd

Pt

Ti Ni

U

C

Tl

Sb

Bi Po

50

The Common Polyvalent IonsFormula Charge Stock Name (new name) Classical Name (old name)

Cu+ 1+ Copper (I) ion Cuprous ionCu2+ 2+ Copper (II) ion Cupric ionFe2+ 2+ Iron (II) ion Ferrous ionFe3+ 3+ Iron (III) ion Ferric ionSn2+ 2+ Tin (II) ion Stannous ionSn4+ 4+ Tin (IV) ion Stannic ionPb2+ 2+ Lead (II) ion Plumbous ionPb4+ 4+ Lead (IV) ion Plumbic ionMn2+ 2+ Manganese (II) ion Manganous ionMn3+ 3+ Manganese (III) ion Manganic ionCr2+ 2+ Chromium (II) ion Chromous ionCr3+ 3+ Chromium (III) ion Chromic ion

Hg+, Hg22+ 1+ Mercury (I) ion Mercurous ion

Hg2+ 2+ Mercury (II) ion Mercuric ion

51

Examples of Ionic Compounds with Polyvalent Elements

Formula Common Name

Stock (new) Name Classical (old) Name or Incorrect name

ions

FeO Wustite Iron(II)oxide Ferrous oxide Fe2+, O2-

Fe2O3 rust Iron(III)oxide Ferric oxide Fe3+, O2-

Fe3O4 Iron(II,III)oxide Ferosso Ferric oxide* Fe2+, Fe3+, O2-

Cu2O cuprite Copper(I)oxide Cuprous oxide Cu+, O2-

CuO Copper(II)oxide Cupric oxide Cu2+, O2-

CrO Chrome black Chromium(II)oxide Chromous oxide Cr2+, O2-

Cr2O3 Chrome green Chromium(III)oxide

Chromic oxide Cr3+, O2-

CrO2 Crolyn Chromium(IV)oxide

Chromium dioxide Cr4+, O2-

CrO3 Chromic acid Chromium(VI)oxide

Chromium trioxide Cr6+, O2-

PbCl2 cotunnite Lead(II)chloride Plumbous chloride Pb2+, Cl-

PbO2 platternite Lead(IV)oxide Plumbic oxide Pb4+, O2-*Ferrosso ferric oxide is a unique combination of Iron(II)oxide and Iron(III)oxide together in a crystalline ionic structure Its formula can also be given as (FeO Fe∙ 2O3) 52

The Ternary Ionic Rules

– First name the metallic element (or ammonium ion) on the left of the formula.

– Then name the polyatomic ion on the right side of the formula. • If the compound is an ammonium salt, then name the

non-metal ion, changing it to end in “ide”

• Examples:– NaNO3sodium nitrate CaCO3calcium carbonate

– K2SO4potassium sulphate Ba(CN)2barium cyanide

– Al2(CrO4)3aluminum chromate NH4Cl ammonium chloride

Polyatomic ions: See Table 8.10 on p. 422

53

Covalent Rules– Name the less electronegative element on the left.– Name the more electronegative element on the

right, changing its suffix to “ide”– Add prefixes to each element to indicate the

number of atoms in the formula:• Mono*=1, di=2, tri=3, tetra*=4, penta*=5, hexa*=6

• Examples:– CCl4 carbon tetrachloride** N2H4 dinitrogen tetrahydride

– PF3 phosphorus trifluoride** P2O5 diphosphorus pentoxide

– CO2 carbon dioxide** CO carbon monoxide**

* The last “o” in mono or the “a” in tetra, penta, or hexa is usually dropped before “oxide” to sound better. (eg. “Carbon monoxide”, not “carbon monooxide”)** The “mono” prefix is usually dropped from the first element of the compound, except when that would cause confusion between two similar compounds. 54

Electronegativity(how much an atom attracts electrons)

I VIII

1 2.2H

II III IV V VI VII 0.0He

2 1.0Li

1.6Be

0.70.9

1.01.4

1.51.9

2.02.4

2.52.9

3.03.1

3.23.5

3.64.0

2.0B

2.6C

3.0N

3.4O

4.0F

0.0

Ne

3 0.9Na

1.3Mg

III B

IV B

V B

VI B

VII B

VIIIB

IB

IIB

1.6Al

1.9Si

2.2P

2.6S

3.2Cl

0.0Ar

4 0.8K

1.0Ca

1.4Sc

1.5Ti

1.6V

1.7Cr

1.6Mn

1.8Fe

1.9Co

1.9Ni

1.9Cu

1.7Zn

1.8

Ga

2.0

Ge

2.2As

2.6Se

3.0Br

0.0Kr

5 0.8Rb

0.9Sr

1.2Y

1.3Zr

1.6Nb

2.2Mo

2.1Tc

2.2Ru

2.3Rh

2.2Pd

1.9Ag

1.7Cd

1.8In

2.0Sn

2.0

Sb

2.1Te

2.7I

0.0Xe

6 0.8Cs

0.9Ba

1.3Hf

1.5Ta

1.7W

1.9Re

2.2

Os

2.2Ir

2.2Pt

2.4Au

1.9Hg

1.8Tl

1.8Pb

1.9Bi

2.0Po

2.2At

0.0Rn

7 0.7Fr

0.9Ra

Rf Db Sg Bh Hs Mt Ds Rg Cn Uut Uuq Uup Uuh Uus Uuo

One of the uses of electronegativity is to decide which element goes first in a formula or name. Usually the element with the lowest electronegativity goes first. Therefore it is called carbon dioxide (CO2), NOT dioxygen carbide (O2C).

There are a few exceptions, like CH4 and NH3, where the more electronegative elements are written first. These formulas have been used for years, and are based on organic chemistry concepts, so it’s unlikely we will change them.

DO use prefixes with covalent compounds

# atoms Covalent prefix Examples

1 Mono…, mon… carbon monoxide (CO), mononitrogen monoxide(NO)

2 Di… carbon dioxide (CO2), dihydrogen dioxide* (H2O2)

3 Tri… nitrogen trichloride (NCl3)

4 Tetra…, tetr… carbon tetrachloride (CCl4), tetramethyl lead ((CH3)4Pb)

5 Penta…, pent… diphophorus pentoxide (P2O5), nitrogen pentafluoride (NF5)

6 Hexa…, hex… sulphur hexafluoride (SF6)

7 Hepta…, hept… bromine heptafluoride (BrF7), heptose (C7H14O7)

8 Octo…, oct… diphosporus octafluoride (P2F8) , octane (C8H18)

9 Nona…, non… nonane (C9H20)

10 Deca…, dec… Decane (C10H22)

*commonly called hydrogen peroxide.

Simplification of Covalent Names• IUPAC (The International Union of Physicists and Chemists) which

oversees naming conventions, allows some simplifications to the systematic names of covalent compounds.– The “mono” prefix may be dropped from an element, unless doing so

could result in confusion.• We usually say “carbon dioxide” rather than “monocarbon dioxide”• However, we always say “carbon monoxide” for CO, since there are two common

oxides of carbon (CO2 and CO)

– A prefix may be dropped from a formula if there is no ambiguity in the formula• Many chemists simply say “hydrogen sulphide” instead of “dihydrogen sulphide”

for the compound H2S. Since H2S is the only common sulphide of hydrogen, this doesn’t cause confusion.

– Knowing when simplification is allowed is a matter of experience. Until you become familiar with the conventions, it is safer to use all the prefixes. It’s not wrong to include them all.• Water can be called hydrogen oxide, but it is perfectly acceptable to use

“dihydrogen oxide” or even “dihydrogen monoxide” 57

Finding Formulasfrom Compound Names

• For covalent compounds, the name usually tells you the formula:

• For example: dinitrogen pentoxide = N2O5

• However:• If the name has been simplified by dropping a prefix you may have

to use the crossover rule, discussed later.• For example: “sulphur fluoride” has had a prefix dropped, so

S(valence=2) F(valence=1) crossover SF2

• “Sulphur fluoride” is the short name for the compound more accurately called sulphur difluoride.

• For ionic compounds, the name never tells you the formula.

• You always use the crossover rule to find the formula.• Example: Sodium oxide is Na1 and O2crossoverNa2O

58

The Crossover Ruleand simple ionic compounds

• The crossover rule is used to find the formula of a compound when the name has no prefixes (ie. all ionic compounds and some covalent compounds that have had a prefix removed)

• Example 1: What is the formula of aluminum sulphide?• Aluminum sulphide : Al S• Ions: Al3+ S2- • Valences (remove signs): Al3 S2

• Cross over: Al2S3

• The formula of aluminum sulphide is Al2S3

59

60

The Crossover Ruleand covalent compounds

• The crossover rule can also be used for covalent compounds if prefixes have been dropped from a name. When a covalent compound’s name has no prefixes at all, check it with the crossover rule.

• Example 1: What is the formula of “sulphur chloride”?• Sulphur chloride: S Cl• Oxidation numbers: S2- Cl- • Valences (remove signs): S2 Cl1

• Cross over: S1Cl2 or SCl2

• The formula of “sulphur chloride” is Al2S3

Notes: 1) The compound “sulphur chloride” should properly be called sulphur dichloride2) The prefixes trump the crossover rule. If any prefixes were used in the name, then they take precedence over whatever formula the crossover rule would give you.

The Crossover Rulesimplifying ionic compounds

• Ionic compounds can often be simplified• Example 1: What is the formula of the compound

made from Barium ions (Ba2+) and Carbide ions (C4-)?• Ions: Ba2+ C4- • Remove the signs Ba2 C4

• Cross over: Ba4C2

• Cancel (divide both by 2) Ba2C

• The formula of barium carbide is Ba2C

Note: Do not simplify covalent compounds by cancellation. Covalent compound formulas must reflect the compound names that include prefixes.

61

Reverse Crossover Rulefor finding the valence of uncertain ions

• Sometimes we can use the crossover rule in reverse to find the valence or ionic charge of an ion we are not certain of, such as an ion of polyvalent metal.

• For example, what is the name of Fe2O3? FeO– They are both iron oxide, but which iron oxide (there are

several types!)

– Fe O Fe O2 3 1 1

Fe has a valence of 3, so the name of the compound is:Iron(III)oxide

There’s a problem here! Oxygen hardly ever has a valence of 1. Let’s double both valences.

2 2

Fe’s proper valence here is 2Iron(II)oxide

The Organic Rules(not studied this year)

A system of names for organic compound exists that is based on the number of carbon atoms they have (as a prefix), and the type of compound they are (as a suffix): alkane (…ane), alkene (…ene) alcohol (…ol), aldehyde (…hyde), ketone (…tone), organic acids, etc.# carbons Prefixes examples

1 Methyl, Formyl Methane, methanol, formaldehyde, formic acid

2 Ethyl, Acetyl Ethane, ethanol, acetaldehyde, acetone, acetic acid

3 Propyl, Propane, propanol, propanoic acid

4 Butyl Butane, butanol, butanoic acid

5 Pentyl Pentane, pentanol, pentanoic acid After this the prefixes resemble those for inorganic compounds, 6=hex, 7=hept, 8=oct, etc.

As you may notice, the common names of some chemicals come from the organic system, such as methane, the common name of carbon tetrahydride (CH4) . For more information on organic nomenclature, see the wikipedia article.

http://en.wikipedia.org/wiki/IUPAC_nomenclature_of_organic_chemistry

Practice

• Page 12, Question #9• Practice sheets:

• Naming ionic compounds• Naming covalent compounds• Naming mixed compounds

64

The Mole Conceptand the Enumeration of Matter

• The Mole: The mole is a unit used to count atoms, ions, molecules, and other fundamental particles.

• A mole corresponds to Avogadro’s Number of particles: 6.02 x 1023 particles.

NA =6.02 x 1023 = 602 000 000 000 000 000 000 000= six hundred and two sextillion

0.5.1

65

Molar Mass• Molar mass is the mass of one mole of atoms or

molecules.• The symbol for molar mass is M (not MM!)

• For elements, molar mass corresponds to the atomic mass found in the periodic table, but expressed in grams/mol rather than amu. For example, the molar mass of carbon, M(C )= 12.011 g/mol, (frequently rounded to 12.0 g/mol)

• For compounds, M is the sum of the masses of all the atoms in the molecule or all the ions in the formula. For example, the molar mass of carbon dioxide molecules is :

M(CO2) =44.009 g/mol, (frequently rounded to 44.0 g/mol)

• that is: 2M(C)+2M(O) or 12.001 +2(15.999) g/mol

0.5.2

66

Diatomic and Polyatomic Elements• Diatomic elements: There are seven elements

whose molecules normally contain two atoms: I2, H2, N2, Br2, O2, Cl2 and F2.

• If finding the molar mass of these elements, remember to double the mass of one atom.

• M (I2) = 253.808 g/mol (not 126.904 g/mol!)

• Polyatomic elements: a few elements, such as sulphur and phosphorus, occur in larger molecules (eg. S8 or P4)

• If a formula like this has been used in a balanced equation, remember to multiply the atomic mass by the appropriate amount (eg. M(S8)=256.52 g/mol)

67How to Remember the Diatomic Elements: I Have No Bright Or Clever Friends

The Mole FormulaThe mole formula is used to convert from grams to moles and vice-versa

n

m

M

Actual mass(g)

# moles(mol)

Molar mass(g/mol)

𝑚=𝑛𝑀

𝑛=𝑚𝑀 𝑀=

𝑚𝑛

Actual mass = # moles x molar mass

# moles = Molar mass =

68

Practice

• Page 14, #12, 13, 14• Practice sheet:

• Moles and Molar mass

69

Physical Changes• A physical change occurs when a substance undergoes

a modification in its appearance or form, but does not alter its nature or characteristic properties.

• In a physical change the molecules or ionic formula of the substance do not change.

• There are 3 main categories of physical change• Change of form, caused by crushing, cutting, grinding, bending,

denting, etc.• Change of phase or state, caused by melting, boiling, freezing,

evaporation, condensation, sublimation, etc.• Change of mixture, caused by dissolving (dissolution without

reaction), blending, stirring together dry ingredients, mixing paints, etc.

70

Phase Change• As a pure substance is heated, its particles move faster.

It changes from a solid state to a liquid state and then to a gaseous state. Your textbook refers to this as “phase change”

• Change of phase is a physical change, since the particles of the pure substance do not (usually) change.

Picky note: What your textbook calls “phase change” should more properly be called “change of state”. Although “phase” and “state” are frequently used as synonyms, the word phase has a broader meaning in chemistry. There are three main states of matter (solid, liquid, and gas) , “phase” includes these three, but may also apply to many other possible phases of matter– including aqueous (a solid dissolved in water), gel (a jelly-like colloidal mixture) etc. In addition, phase can refer to a boundary between two similar phases that don’t mix, for example, a liquid mixture could have an oily phase and a watery phase that contact each other but do not mix.

0.6.1

71

LiquidSolid

Gas

Melting (fusion)

Freezing (solidification)

Vaporization

Liquid

Liquid CondensationSubl

imati

on

Solid

Con

dens

ation

Terminology associated with

Change of Phase

Rapid vaporization is called “boiling”,

Slow vaporization is “evaporation”

Sublimation occurs when a material

“evaporates” from a solid straight to a gas, like dry ice or iodine.

ExothermicProcess

EndothermicProcess

72

Comparison of the States of MatterSolid Liquid Gas

Shape Definite Variable Variable

Volume Definite Definite Variable

Compressibility Incompressible Incompressible Compressible

Fluidity Not Fluid Fluid (flows) Fluid (flows)

Particle separation Close together Close together Far apart

Motion of particles Vibration only Rotation and vibration

Rotation, vibration and translation

73

Phase Markers

• During the course of the year, you will often notice small letters in parenthesis added formulas in equation. These “phase markers” are inserted whenever it is important to know what state or phase the reactants or products are.

• The most important phase markers are:• (s) = solid: the substance is a solid or a powder• (l) = liquid: the substance is a pure liquid• (g) = gaseous: the substance is a gas• (aq) = aqueous: the substance is dissolved in water

Eg: NaCl(s) H2O(l) NH3(g) NaCl(aq)

74

Dissolution and Solubility

• In dissolution, one or more solutes are mixed into a solvent to create a solution.

• During dissolution:• The mass of the substances does not change.• The total volume is usually slightly less than the sum of

the volumes of the components (since some particles pass into the spaces between other particles)• When the solvent cannot dissolve any more of the

solute, the solution is saturated.

0.6.2

75

Dissolving = Physical Change• Remember that dissolution is normally considered

a physical change, not a chemical one. The material mixes with the solvent, but is not significantly altered by it

• In a few cases a material will react with the solvent, rather than just dissolve. For example, trying to dissolve sodium in water, or baking soda in vinegar will produce a reaction. In this case a chemical change has occurred as well.

eg: Na(s) + H2O(l) NaOH(aq) + H2(g)

• Ionic compounds may “dissociate” while dissolving, that is, their ions may separate by some distance. While this may seem like a chemical change, it is not a permanent condition, and is considered to be a physical change.

eg: NaCl(aq) Na+(aq) + Cl-

(aq) (dissociation of salt)76

Dissolution of Ammonia Gas in Water (an extreme case of solubility at 25°C)

• 100g of water + 50g of ammonia 150g of ammonia solution

• 100 mL of water + 72058 mL of ammonia 101 mL of NH3(aq) solution

• If you try to dissolve more than 50g of ammonia in 100 mL of water, you won’t be able to. There will be leftover ammonia!

100

72.058 litres NH3(g)

100

150

50g of NH3(g)150

101

Ammonia is a great example, because water can absorb what seems like a huge amount of ammonia gas before it becomes saturated. Mass-wise, its actually half the weight of the water, but volume-wise its over 720 times greater!

55 g of NH31005 g

+

+

+ +

gg

g g

mL mL

77

• Solubility indicates the maximum amount of solute that can dissolve in a given volume of solvent at a given temperature.

• Solubility is usually expressed as grams of solute per 100 mL of solvent (g/100mL).

• A substance’s solubility can vary with temperature:

• Solubility of solids usually increases with temperature• Solubility of gases usually decreases with temperature• Solubility of gases can also be affected by pressure.

78

Solubility Curves(Graphs of Solubility vs. Temperature. See page 16)

• Notice how most of the solids become more soluble at higher temperatures – KNO3, for example, starts at a

mere 10 g/100 mL at 0°C, but goes right off the top of the chart by 70°C

• Notice that most of the gases become less soluble at high temperatures – NH3 goes from 90 g/100mL at 0°C

to less than 10 g/100 mL at 100°C

79

Concentration and Dilution• Concentration is the ratio of dissolved solute to total

amount of solution.• General formula for concentration is:

• But concentration can be expressed in many different units, including:

• g/L (grams per Litre) g/mL (grams per millilitre)• % (by volume) % (by mass)• ppm (parts per million) mol/L (molar concentration)

• Molar concentration is the most important.

0.6.3

80

𝑐𝑜𝑛𝑐𝑒𝑛𝑡𝑟𝑎𝑡𝑖𝑜𝑛=𝑎𝑚𝑜𝑢𝑛𝑡 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒𝑎𝑚𝑜𝑢𝑛𝑡 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛

Molar Concentration(molarity)

• The molar concentration is the number of moles of solute that is dissolved in one mole of the solution.

• Molar concentration can be represented by the letter C, or by square brackets [] or occasionally by a capital M used as a unit (molarity). Any of the following notations could represent a 2.0 mol/L solution of hydrochloric acid:

CHCl = 2.0 mol/L

[HCl] = 2.0 mol/LCHCl = 2.0 M

The correct unit for molar concentration is mol/L, although this is sometimes

abbreviated with a capital M for molarity

81

Molar Concentration Formula

n

C VConcentration

(mol/L)

# moles(mol)

Volume(L)

C =

Molar Concentration =

V =

n = CV82

Dilution

• Dilution is a physical change that lowers the concentration of a solution by adding more solvent.

• The dilution formula is:

C1V1 = C2V2

Where: C1 is the concentration before dilution,V1 is the volume before dilutionC2 is the concentration after dilutionV2 is the volume after dilution

83

Assignments

• Read pages 15 to 18• Do page 16

• Questions 15 to 17

• Do page 19:• Questions 18 to 23

84

Electrolytes

• Electrolytes are substances which, when dissolved in water, allow the solution to conduct electricity.

• Electrolytes are usually ionic compounds.• Electrolytes “dissociate” into positive and

negative ions when they dissolve.• There are three main types of electrolytes:

Acids, Bases, and Salts.• Most solid electrolytes do not conduct

electricity until they are dissolved.

0.6.4

85

Electrolyte CharacteristicsAcids Bases Salts

Ions: Release H+ ions Release OH-ions Metal and non-metal ions

neutralization Neutralize bases Neutralize acids Products of neutralization

pH pH is less than 7 pH is greater than 7 pH variable, close to 7*

Litmus Turn litmus paper red Turn litmus paper blue Don’t change litmus*

Phenolphthalein Stays clear Turns red/purple Stays clear*

Formula H + non-metal Metal + OH Metal + non-metal

Dissociation eg: HCl(g) H+(aq) + Cl-

(aq) NaOH(s)Na+(aq) + Cl-

(aq) NaCl(s)Na+(aq) + Cl-

(aq)

pH ScaleThe pH (positive Hydrogen potential) scale is used to measure the relative acidity or alkalinity of a solution. It is in theory open-ended, but in practice runs from 0 to 14.

Strong Acids Weak Acids Neutral Weak Base Strong Base

0 1 2 3 4 5 6 7 8 9 10 11 12 13 14

86* Some salts are slightly acidic (aluminum salts) or slightly basic (carbonates)

Assignments

• Read page 19• Do page 20

• Questions 24-27 • Question 28

87

0.7 Chemical Changes

• Chemical changes occur when substances (reactants) react to form new substances (products).

• The products differ from the reactants:• They have different characteristic properties.• They have different molecular or ionic arrangements.

Reactants ProductsReactants on the

Left side of equation

Products on the Right side of

equation

88

• Indications that a chemical change has taken place include:

• Release of a gas (effervescence)• Significant change in colour• Formation of a precipitate (solid from two solutions)• Change of energy in the form of heat, light or explosion.

• Parts of a chemical equation:

4 Fe (s) + 3 O2 (g) 2 Fe2O3 (s)

Reactants Product

Chemical equation

Coefficients4:3:2

(used for balancing)

Change to

Phases(s) Solid (l) liquid (g) gas

(aq) dissolved in water

Indexes* 2,2,3

Number of atoms in the molecules

*Yes, I am fully aware that the dictionary says that the correct plural of index is indices, but for clarity I am using the term the text uses. 89

O OHH

HH H

Conservation of Mass• During a chemical reaction, mass is neither

lost nor gained• The total mass of all the reactants is equal to

the total mass of the products.• This is because no atoms are created or destroyed

during the reactions. The atoms are just rearranged.

• The balancing of chemical equations is based on the law of conservation of mass.

H OH O

2H2 + O2 2H2OH

m reactants = m products

0.7.1

90

Balancing Equations• Balancing means adding coefficients in front of the

formulas of an equation so that it will conform to the law of conservation of mass

• A word equation names the reactants and products• A skeleton equation is an unbalanced equation• A balanced equation respects conservation of mass.

• Rules for balancing equations:• Only coefficients may be added or changed. The indexes in

formulas must not be changed.• You do not need to write the coefficient 1. It is understood.• Balanced equations should be reduced to the lowest terms.• When an equation is properly balanced, the total number of

atoms of each element on the left and right sides will be equal.

0.7.2p. 22

91

Stoichiometry

• Stoichiometry is the study of the relationships between the amounts of substances (reactants and products) that take part in a chemical reaction.

• Stoichiometry can be used to:• Calculate the amount of reactants need for a reaction• Calculate the expected amount of product from a

reaction.

0.7.3p. 23

92

Steps for Stoichiometry1. Balance the equation, or verify that the equation you have

been given is properly balanced.2. Use the coefficients to find the mole ratios3. Write the amount in moles of the known reactant under

the corresponding mole ratio number.• If the amount is given in grams, convert it to moles using the

mole formula.

4. Write an x under the mole ratio of the substance you are looking for. Ignore the other substances for now

5. Change the : to =; Solve for x by cross multiplying.6. The result is the answer in moles.• If you need an answer in grams, convert using the mole formula

(with the proper molar mass!) 93

Example1. Balance the equation, or verify

that the equation you have been given is properly balanced.

2. Use the coefficients to find the mole ratios

3. Write the amount in moles of the known reactant under the corresponding mole ratio number.• If the amount is given in grams,

convert it to moles using the mole formula.

4. Write an x under the mole ratio of the substance you are looking for. Ignore others

5. Change : to =. Solve for x by cross multiplying.

6. The result is the answer in moles.• If you need an answer in grams,

convert using the mole formula (with the proper molar mass!)

Problem: 8 grams of hydrogen are burned with oxygen to make water. How much oxygen was used?

Step 1: H2 + O2 H2O (skeleton)

2H2 + O2 2H2O (balanced)

Step 2: mole ratios 2 : 1 : 2

Step 3: known reactant is 8g hydrogen. To convert it to moles we must divide by the molar mass of hydrogen, 2.0; That gives us 4 moles of hydrogen. Write this under the corresponding mole ratio 2 : 1 : 2

4

Step 4: write an x 2 : 1 : 24 x

Step 5: cross multiply 2 = 1 so x = 2 mol 4 x

Step 6: to get the answer in grams, multiply the 2 mol by the molar mass of oxygen (32 g/mol) to give us the answer 64 g of oxygen is used.

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Assignments

• Read pages 21-24• Do Question 30 on page 23• Do Questions 31 and 32 on page 24

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0.8 Examples of Chemical Reactions

• There are many types of chemical reaction. Among the most important types are:

• Acid-base reactions• Synthesis, Decomposition and Precipitation Reactions• Endothermic and Exothermic Reactions• Oxidation and Combustion• Photosynthesis and Respiration

• These are just a few of the types. Some your textbook does not mention include:

• Single Replacement • Double Replacement

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Acid – Base Reactions• When an acid and a base are mixed:

• The H+ ions from the acid join the OH– ions from the base to make H2O, that is water.• The other ions, usually a metal and a non-metal ion,

join to form a salt whose nature depends on the reagents.• If the original solutions contained equal amounts of H+

and OH-, then the mixed solution will be neutral.• If there was a surplus of H+ or OH- ions, then the

resulting solution will be slightly acid or slightly basic.

In general: ACID(aq) + BASE(aq) WATER(l) + A SALT(aq)

Example: HNO3(aq) + KOH(aq) H2O(l) + KNO3(aq)

The words Neutralization

and Titration are also

associated with this process

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Synthesis, Decomposition, Precipitation

• Synthesis is when two or more reactants combine to form a single product.

• Eg. 2Na(s) + Cl2(g) 2 NaCl (s)

• Decomposition is when a single reactant breaks into two or more products.

• Eg.

• Precipitation is when a solid powder is formed by the mixing of two solutions.

• Eg

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Endothermic and Exothermic

• Endothermic reactions are chemical reactions that absorb energy. Endothermic reactions usually make their immediate surroundings cooler.

• Reactants + Energy Products• Exothermic reactions release heat. They often

make their surroundings warmer.• Reactants Products + Energy

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Oxidation and Combustion0.8.4

• Oxidation is process where a substance combines with an oxidizer (usually O2, but O3, F2, Cl2, N2O and other substances work too).

• Your textbook incorrectly states that an oxide is always formed, but sometimes chlorides or fluorides are formed by oxidation.

• Slow oxidation takes time to happen.• Eg. The rusting of iron: 4 Fe + 3 O2 2 Fe2O3

• Combustion is rapid oxidation that produces heat and flames.

• Eg. Combustion of gasoline: 2 C8H18 + 25 O2 16 CO2 + 18H2O

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Photosynthesis and Respiration

• Life on Earth depends on two related chemical processes:

• Photosynthesis is the chemical reaction in which organisms, such as plants, transform radiant energy from sunlight into stored chemical energy.

6 CO2 + 6 H2O + energy C6H12O6 + 6 O2

• Respiration is the process by which organisms release stored chemical energy in sugars and other organic compounds in living cells.

C6H12O6 + 6 O2 6 CO2 + 6 H2O + energy

0.8.5p. 27

Error in textbook: On p. 27, respiration is referred to as a “combustion” reaction.What the textbook means, of course, is that is an “oxidation” reaction. 101

Assignments

• Page 25 #33• Page 26 #35-36• Page 27 #37-38

Chemical Bonds

• Chemical bonds are the forces that bind atoms together into larger structures, such as molecules or crystal lattices.

• Chemical bonds are the result of exchange or sharing of electrons between two atoms, which causes the formation of a compound or diatomic or polyatomic element.

• There are many types of chemical bond. The three most important are:

• Metallic: metal to metal, found in alloys• Ionic: metal to non-metal, found in salts• Covalent: non-metal to non-metal, found in molecules.

Not Studied

Studied

Studied

Sugar, covalent moleculeSalt, ionic crystal lattice

Your textbook has little about metallic bonds, but since we don’t study alloys in detail, this is not a problem.

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Ionic Bonding• An ionic bond forms when electrons are exchanged

between two atoms.

• This type of bond forms when one of the elements has a much higher electronegativity (X) than the other. This usually happens between a metal atom and a non-metal atom.

• Ionic bonds are between negative and positive ions• Ionic bonds do not form strong, distinct molecules. In most

ionic solids, the ions form a crystal lattice of alternating positive and negative particles. Some chemists prefer the term “formula units” to “molecules” when talking about ionic compounds.

Alternating particles do not overlap.

A crystal lattice structure with

alternating ions

0.9.1p. 28

NaNa+ Cl–cation anionCl

Sodium has an “extra” electron in its outer shell

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Na+ Cl–

A sodium chloride formula unit

N ClCl

Cl A covalent molecule

Chlorine “needs” another electron in its outer shell

X.= 3.16X= 0.93 ΔX = 2.23

Electronegativities(Supplemental)

• The electronegativity (X )(Greek letter chi or curly x) of an element can be found from the periodic table in front of your textbook.

• It indicates how much an element attracts electrons.• The greater the electronegativity difference between two

elements, the more likely they will form an ionic bond.• No bond is 100% ionic or 100% covalent, but we treat them that

way for simplicity.• The character of a bond is based on several things, in addition to

electronegativity, so the chart below is an approximation.

ΔX Character of bond Name of Bond type

1.7 to 3.9 > 50% ionic Ionic0.4 to 1.7 10% to 50% ionic Polar-Covalent0.0 to 0.4 <10% ionic Covalent

Covalent Bonding

• A covalent bond forms when electrons are shared between two atoms.

• This type of bond forms when two elements have similar electronegativity. This usually happens between two identical atoms, or between two non-metal atoms.• Covalent bonds can be single (sharing one pair of

electrons), double (sharing two pairs) or triple (sharing three pairs)• Covalent compounds form true, strong molecules.

They are sometimes referred to as molecular compounds.

C OO

Shared electrons in overlapping shells

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Illustrating Covalent BondsWith Rutherford-Bohr models:

With Lewis electron dot diagrams:

In either case, we draw the atoms to show a stable number of electrons (usually 8) in the outer shell of each atom involved in the covalent bond.

N ClCl

Cl Another way to illustrate covalent bonds is with overlapping circles

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Energy• Energy is the ability to do work or make a change.

• There are many types of energy, a few of which are listed in the table below:

Form of Energy Associated with Example

Kinetic Energy An object’s movement Car driving along a road

Thermal Energy Agitation of particles Boiling water

Radiant Energy Electromagnetic waves Light, microwaves, radio waves

Gravitational* Energy Object’s position above ground Water behind a dam

Elastic* Energy Compressed/stretched materials A spring that has been stretched

Electric* Energy Force between electric charges Charged particles in a storm cloud

Nuclear* Energy Stored in the nucleus of atoms Uranium in a reactor

Chemical* Energy Stored in the bonds of molecules Energy in gasoline or glucose

* The word “potential” is often inserted to indicate that these associated with potential energy.

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Kinetic Energy

• Kinetic energy is the energy associated with the movement of an object, or with the movement of its particles (molecules).

• Kinetic energy depends on the mass of the object and the velocity of its motion.

Where: Ek= kinetic energym= mass of the objectv= velocity of the object

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0.10.2 Potential Energy

• Potential Energy is energy stored in a body that can be transformed into another form of energy.

• Potential energy is sometimes referred to as “hidden energy”, since it is difficult to observe and measure.

• There are several types of potential energy, including:

• Gravitational Potential Energy (important in physics)• Chemical Potential Energy (important in chemistry)

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• Gravitational Potential Energy is the product of an object’s mass, its height above the ground, and the gravitational acceleration.

• Chemical Potential Energy (Enthalpy) is associated with the energy in the bonds between the particles of a material.

Where: Ep = Gravitational Potential Energy in joulesm = mass of the object in kilogramsg = gravitational acceleration (9.8 m/s2 on Earth)h = height of the object above a reference point (such as the ground)

We will devote a section later in the course to calculating enthalpy.

0.10.2p. 30

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Conservation of Energy• The law of conservation of energy states that

energy cannot be created or destroyed in chemical reactions, but it can be changed from one form to another.

• Potential energy can change to kinetic and vice versa

• Mechanical Energy is the total energy of an isolated system.

Where: Em = total Mechanical EnergyEp = Potential EnergyEk = Kinetic Energy

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0.10.4 Thermal Energy & Temperature

• Thermal energy or “heat” is a form of energy possessed by a substance due to the agitation of its particles. It depends on:

• The mass of the substance• The temperature of the substance• The specific heat capacity of the substance

𝑄=𝑚𝑐∆𝑇Where: Q = amount of heat energy in joules

m = mass of the substance heated in grams (usually the water in a calorimeter)c = specific heat capacity of the substance heated, in j/g∙°CΔT = the change in temperature in °C 113

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0.11 Fluids• Compressible & Incompressible Fluids

• Substances that flow, like liquids and gases, are fluids• Gases are compressible fluids• Liquids are incompressible fluids

• Pressure• Pressure is the force exerted on a surface.• The standard unit of pressure is the kilopascal (kPa)• Formula for pressure: Pressure = Force divided by Area.

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END of MODULE 1

• Prepare for the module 1 test.– Read up to page 33 in your text book– Prepare study notes.–

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Documents

Pages 3 to 33 “Quantum Chemistry” Target Completion Date: October 1