A. WHMIS workplace hazardous materials

information system

all chemicals are treated with

WHMIS has been developed to provide guidelines for of reactive materials

Energy and Matter in Chemical Change Understanding Matter

respect

handling, storage and disposal

compressed gas corrosive

flammable and combustible

poisonous and infectious

material causing immediate and serious toxic

effects

biohazardous infectious material

oxidizing material

dangerously reactive material

poisonous and infectious causing other

toxic effects

B. Properties are

properties you can

eg) colour,

physical propertiessee and measure

boiling pointmalleability,ductility,density, state,

are properties used to describe how substances will

eg) combustion,

chemical properties react with each other

decompositionrusting,

C. Classification of Matter

Matter

Pure Substance Mixtures

Elements Compounds Heterogeneous Homogeneous

Metalloids

Non-metals

Ionic

MolecularAlloys Solutions

Colloids SuspensionMetals

Mechanical Mixture

mixture is a mixture of that has

eg)

is a homogeneous mixture of

eg) brass =

steel =

homogeneous 2 or more substances uniform properties (that appear as one)

alloy 2 or more metals

copper + zinc

iron + chromium + carbon

is a mixture of and

eg)

heterogeneous mixture 2 or more substances individual components are visible

is a mechanical mixture in which the

eg) mud

are mechanical mixtures in which the suspended substances

eg) milk

suspensioncomponents are in different states

colloids cannot be easily separated from the other substances

are substances (metals, non-metals or metalloids)

eg)

are cannot be

eg)

elements that are pure and cannot be broken apart

compounds two or more elements combined; separated by physical means

D. Atomic Structure are the building blocks of consist of a and a

atoms

ALL matter

tiny nucleus huge “cloud” region

makes up of the of an atom

makes up most of the of an atom

nucleus 99% mass

cloud region volume

Subatomic Particles

1. Protons (p+) charge found in determines the

2. Neutron (n0) charge found in used to

positive

nucleus

type of element

no

nucleus

hold nucleus together

3. Electron (e) subatomic particle charge found in arranged in

maximum # of electrons in each level:

Level 1 =Level 2 =Level 3 =Level 4 =

smallest

negative

“cloud” region

energy levels

2 e

8 e

8 e

we will only go up to 2 e

Atomic Mass is the

mass of a

for all atoms is

is the (rounded to the nearest whole number)

used to find thenumber of neutrons

atomic mass unit (amu)proton or neutron =1.7 1024 g

net charge zero

# e = # p+

mass number sum of the protons and neutrons

eg) lithium

atomic number =

atomic mass =

# protons =

# neutrons =

3

6.94 = 7

3

7 = 4– 3

Isotopes atoms that have the

but a

on the periodic table is an based on the of all of the element

same number of protonsdifferent number of neutrons

atomic massaverage mass percentage abundances naturally occurring isotopes

isotope notation:

A XZ

X =A =Z =atomic # (#p+)

mass # (#p+ + n°)symbol

Examples

64 Cu29

# p+ =# e- =#n° =

62 Cu29

# p+ =# e- =#n° =

29 2964 - 29 = 35

2929

62 - 29 = 33

copper - 64

copper - 62

Dalton’s Atomic Theory: 1808

E. Atomic Theory

all matter is composed of tiny, indivisible particles called

atoms of an element have

atoms of different elements have

atoms

identical properties

different properties

atoms of two or more elements can combine in

eg) H:0 ratio 2:1H:0 ratio 2:2

waterhydrogen peroxide

constant ratio to form new substances

H2O H2O2

J.J. Thompson: 1897credited with discovery of

model or modelatom is a which is ,

with like raisins in a bun most of the is associated with the

electrons

“raisin bun” “plum pudding”

sphere positivenegative electrons embedded in it

masspositive charge

Ernest Rutherford: 1911atoms have a which is

and has most of the

most of the atom is occupied by the moving

proposed the existence of

nucleus positivemass

empty spacenegatively charged electrons

protons

Neils Bohr: 1913electrons move in

around the

cannot exist between orbits

circular orbitsnucleus

James Chadwick: 1932 showed that the nucleus must contain

to account for all of the atom’s

heavy neutral particlesmass(neutrons)

exact locations of electrons are , but the probable location in a can be predicted

electrons have

model Schrodinger/de Broglie: 1930

quantum mechanical

distinct energy levels

not definedregion of space

F. The Periodic Table the periodic table was developed by

Dmitri Mendeleev in the mid

1. Atomic Number

number of in one atom of an element from left to right and top to bottom

1800’s

protons

increases

http://www.chemicool.com/

3 major categories:

1. are are found on

metals

have high luster;left side of stair case

good conductors,strong,malleable(pound into thin sheet),ductile(can draw into a wire, bendable),

potassium copper

mercury

2. Properties

2. are properties to metals; found on of staircase

non metalsopposite

right sidenon-lustrous,

poor conductors,weak, etc…

sulphur

bromine

iodine

3. show properties of ; found staircase

above and below both metals and nonmetals metalloids

arsenic

siliconboron

The Periodic Table

3. Groups

18 vertical are called

2 labelling systems: Roman Numerals with letters or ordinary numbers

IA, IIA, IIIB, IV, V, VI, VII, VIII etc

1, 2, 3, 4, 5, 6, 7, 8

columns groups or families

Group 1 (IA)Alkali Metals

Group 2 (IIA) Alkaline Earth Metals

Group 17 (VIIA) Halogens

Group 18 (VIIIA)Noble (Inert) Gases

Lanthanide Series (57-71)Rare Earths

Groups 3-12 (B series)Transition Metals

The Element Song!

elements in each group share although changes

reactivity for and for

similar chemical properties (reactivity) intensity

increases down groupmetals up group nonmetals

indicates how many electrons are in thegroup number

outermost energy level(ignoring the “1” in groups 10 and above)

Reaction of Alkali Metals with Water

http://www.youtube.com/watch?v=m55kgyApYrY

http://www.youtube.com/watch?v=jhg0WsINmPc&feature=related

4. Periods

show a which

change from each time you move to a new period you

horizontal rows

trend in reactivityleft to right

start the trend over

G. Electron Energy Level Representations (EELR)

nucleus –

energy levels – shows # of in each level (# of levels = )

are the e in energy level ( ignoring the 1 in front of groups 13-18 )

shows # p+ and n0

e

period element is in

valence electrons outermostsame as group #,

Examples sodium

# p+ = # e- =#n° = 23 – 11 = 12

p+ =n° =

2 e-8 e-

1 e- 11 e-

atomic # mass #

1122.99 = 23

1111

1112

Na

argon

# p+ = # e- =#n° = 40 – 18 = 22

p+ =n° =

2 e-8 e-

8 e- 18 e-

atomic # mass #

1839.948 = 40

1818

1822

Ar

Draw the EELR for the following:

1. potassium

2. chlorine

3. beryllium

4. calcium

potassium

p+ = 19n° = 20

2 e-8 e-

8 e-

19 e-

1 e-

atomic # mass #

1939.10 = 39

# p+ = # e- =#n° = 39 – 19 = 20

1919

K

chlorine

p+ = 17n° = 18

2 e-8 e-

7 e- 17 e-

atomic # mass #

1735.45 = 35

# p+ = # e- =#n° = 35 – 17 = 18

1717

Cl

beryllium

p+ = 4n° = 5

atomic # mass #

49.01 = 9

# p+ = # e- =#n° = 9 - 4 = 5

44

2 e-2 e- 4 e-

Be

calcium

p+ = 20n° = 20

2 e-8 e-

8 e-

20 e-

2 e-

atomic # mass #

2040.08 = 40

# p+ = # e- =#n° = 40 – 20 = 20

2020

Ca

H. Octet Rule atoms tend to be when the

outer energy level is

the octet rule states that atoms bond in a way to have a (there are exceptions)

atoms will either in order to satisfy this octet rule

compounds are formed when

stablefull of electrons

full valence energy level

share electrons, or gain or lose electrons

two or more different elements bond together either by sharing or transferring electrons

I. Ions are

that have a number of p+ and e are equal

most atoms try to achieve the electron configuration of a

ions atoms or groups of atomsnet charge

not

noble gas

means having the as another atom or ion eg) fluorine to be

isoelectronic with

isoelectronic same number of e

gains an electronneon

neon atom

eg) potassium to be isoelectronic with

loses an electronargon

potassium atom argon atompotassium ion

+

Cations charged ions electrons to obtain a

electron configuration form

eg) sodium to completely empty the last energy level

Na+ has

on a metal ion is the for groups (ignore the 1 in front of 13 and 14)

positively

lost stable

METALS cations

loses one e

10 e and 11 p+

charge same as the group number 1,2,3,13,14

Anions charged ions to obtain a

electron configuration (full energy level) form

have ending

eg) oxygen to completely fill the last energy level

O2- has and is called the

on a non-metal is (not Roman Numerals)

negativelygained electrons stable

NON-METALS anions

“ide”

gains two e

10 e and 8 p+

oxide ion

charge 18 - group number

EELR’s for Ions

number of p+ =number of e =number of n0 =

atomic numbernumber of p+ – chargeatomic mass – atomic number (# of p+)

nitride ion

# p+ = # e- =#n° = 14 – 7 = 7

p+ = 7n° = 7

2 e-8 e-

10 e-

atomic #

mass #

Examples 7

14.01 = 14

77 – (-3) = 10

N3-

calcium ion

# p+ = # e- =#n° =

p+ = 20n° = 20

2 e-8 e-

8 e-

18 e-

atomic #

mass #

20

40.08 = 40

2020 – (+2) = 1840 – 20 = 20

Ca2+

Draw the EELR for the following:

1. sulphide ion

2. aluminum ion

3. chloride ion

4. magnesium ion

sulphide ion

# p+ = # e- =#n° =

p+ = 16n° = 16

2 e-8 e-

8 e- 18 e-

atomic #

mass #

16

32.07 = 32

1616 – (-2) = 1832 – 16 = 16

S2-

aluminum ion

# p+ = # e- =#n° =

p+ = 13n° = 14

2 e-8 e-

10 e-

atomic #

mass #

13

26.98 = 27

1313 – (+3) = 1027 – 13 = 14

Al3+

chloride ion

# p+ = # e- =#n° =

p+ = 17n° = 18

2 e-8 e-

8 e- 18 e-

atomic #

mass #

17

35.45 = 35

1717 – (-1) = 1835 – 17 = 18

Cl-

magnesium ion

# p+ = # e- =#n° =

p+ = 12n° = 12

2 e-8 e-

10 e-

atomic #

mass #

12

24.30 = 24

1212 – (+2) = 1024 – 12 = 12

Mg2+

J. Elements elements exist as

chemical formula is simply the followed by the at room temperature

eg) sodium mercurycopper

Na(s)Hg( )Cu(s)

symbolstate

metallic single atoms (monatomic)

(not including ) and are called

non-metals noble gases do not exist as single atoms

molecular elements(diatomic, polyatomic)

chemical formula is the and the at room temperature memorize the subscripts (flagpole)!!!

H2

N2 O2 F2

P4 S8 Cl2

Br2

I2

symbol with the subscript state

Monatomic

Summary:

Polyatomic

Diatomic

C(s), noble gases, all metals

H2(g), N2(g), O2(g), F2(g), Cl2(g), Br2(), I2(s)

P4(s), S8(s)

K. Molecular Compounds are formed

when bond together bonded by which is the force of attraction between atoms that are electrons

molecular compounds two or more nonmetals

covalent bondssharing

properties:

1. when

2. in water to form either a or an3. at room temperature

do not conduct electricity dissolved in water

dissolve neutral molecular solutionacidic solution

solids, liquids or gases

Naming (when only two elements combined) give the for the

(with the if there is ) then give the and include the

Note: if the first element is , do put a prefix (these are acids!)

atom name first elementprefix more than one

name for the second element with “ide” ending prefix

hydrogen not

1 = 6 =

2 = 7 =

3 = 8 =

4 = 9 =

5 = 10 =

mono

di

tri

tetra

penta

hexa

hepta

octa

nona

deca

Prefixes

Examples

1. CO(g)

2. CO2(g)

3. P4O10(s)

4. BrH7(s)

carbon monoxide

carbon dioxide

tetraphosphorus decaoxide bromine heptahydride

Writing Formulas simply write each followed by the

Examples

1. oxygen dibromide

2. diphosphorus pentasulphide

3. carbon tetraiodide

4. phosphorus pentachloride

OBr2

P2S5

CI4

PCl5

symbolsubscript (from prefix)

some molecular compounds have classical names…memorize them!!!

ammoniawater

NH3 =H2O =

H2S =HF, HCl, HBr, HI =

CH4 =CH3OH =C2H6 =C2H5OH =

C6H12O6 =

hydrogen sulphideno prefixes

methane

methanolethane

ethanol

glucose

sucroseC12H22O11 =

hydrogen peroxideozoneO3 =

H2O2 =

L. Ionic Compounds are formed when

electrons are , allowing to bond together

is the force of attraction between

ionic compoundstransferred

oppositely charged ions

ionic bondoppositely charged ions

properties:

1. when

2. when

3. at room temperature

conduct electricity dissolved in water

separate into ions dissolved in water

crystalline solids

1. Monovalent Ionic Compounds means there is

on the metal

eg)

Naming give the name for each are

have ending eg) NaF

Na2S

the 2 means that two sodium ions are bonded with one sulphide ion… this doesn’t matter for naming

sodium fluoridesodium sulphide

monovalent one charge

Na+, Ca2+

metal + nonmetal

ion…non-metals

normal,“ide”

metals

Try These:

1.     LiF

2.     KCl

3.     BeS

4.     Rb3P

5.     MgF2

6.     Na2O

7.     CsBr

8. BaCl2

lithium fluoride

potassium chloride

beryllium sulphide

rubidium phosphide

magnesium fluoride

sodium oxide

cesium bromide

barium chloride

Writing Formulas look up the and

write them listing the

using numbers ***in ionic compounds, the total positive charges must equal the total negative charges…the net charge is zero

eg) sodium oxide Na O

1+ 2 1+ 2 = 2+ 2 1 = 2

2

symbol for each ionmetal ion first

balance the charges subscript

calcium phosphide Ca P

2+ 3 2+ 3 = 6+ 3 2 = 6

3 2

1.     magnesium chloride

2.     calcium chloride

3.     zinc sulphide

4.     silver sulphide

5.     germanium oxide

6.     calcium arsenide

7.     magnesium nitride

Try These:

MgCl2

CaCl2

ZnS

Ag2S

GeO2

Ca3As2

Mg3N2

2. Multivalent Ionic Compounds metal ions that have

eg)

the listed is the

more than one possible charge

Cu2+, Cu+, Fe3+, Fe2+

transition metal + nonmetal

first charge most common

Naming same rules as before the difference is you must include

containing the of the metal ion in

figure out the charge on the metal by using there are in the nonmetal ions

eg) CuI

TiBr4

Ti3P4

copper (I) iodide

titanium (IV) bromide

titanium (IV) phosphide

how many negative charges

bracketscharge

Roman Numerals (I,II,III,IV,V,VI,VII)

1.     AuBr

2.     CrCl2

3.     Co2O3

4.     VS2

5.     PuN2

Try These:

gold (I) bromide

chromium (II) chloride

cobalt (III) oxide

vanadium (IV) sulphide

plutonium (VI) nitride

Writing Formulas same rules

the on the metal is given to you in the

eg) iron (II) oxide

tin (II) chloride

chromium (III) sulphide

FeO

SnCl2

Cr2S3

chargebrackets

1.     chromium (II) sulphide

2.     nickel (III) chloride

3.     vanadium (IV) phosphide

4.     gold (III) iodide

Try These:

CrS

NiCl3

V3P4

AuI3

3. Mixed Ionic Compounds

eg) PO43, SO4

2, HCO3 etc.

Naming give the name for the then give

the name for the

***NH4+ (ammonium ion) is the only

positive complex ion…you’ll see it in the place of a metal

metal ion + polyatomic ion (complex ion)

first ioncomplex ion

potassium iodate

sodium acetate

ammonium nitrate

1.     KIO3

2.     NaCH3COO

3.    MgSO3

4. NH4NO3

5.     Ca3(PO4)2

Try These:

magnesium sulphite

calcium phosphate

Writing Formulas …look up the

for each ion then the charges using subscripts if you must multiply one of the complex ions, put around it first then write the subscript eg) aluminum phosphate

aluminum chlorate

calcium sulphite

scandium acetate

ammonium sulphate

AlPO4

Al(ClO3)3

CaSO3

Sc(CH3COO)3

(NH4)2SO4

same as before symbolbalance

brackets

Summary: Ionic vs. Molecular

Ionic Molecular

metal + non-metal (or polyatomic ions)

no prefixes prefixes

all non-metals

solids

solutions conduct

solutions are usually basic or neutral

solids, liquids or gases

solutions do not conduct

solutions are usually acidic or neutral

4. Acids and Bases matter can be subdivided into

based on its properties:

1. acids

2. bases

3. neutral substances

three groups

let’s look at the of acids and bases:

properties

Acids Bases usually soluble in

H2O

taste sour

taste bitter

litmus - litmus -

neutralize bases neutralize acids

reacts with metals to produce H2(g)

vinegar, lemon juice Tums, ammonia

usually soluble in H2O

conduct electricity

conduct electricity

feel slippery

Indicators Indicators

bromothymol blue - bromothymol blue -

phenophthalein - phenolphthalein –

pink blue

yellow blue

colourless bright pink

0 1 5 7 11 14

strong acid strong base

the was devised to indicate how a substance is

stomach acid coffee water antacid

drain cleaner

Acids BasesNeutral

pH scaleacidic or basic

Naming and Writing Formulas for Bases same rules as

many of them contain the

eg) NaOH

KOH

Ba(OH)2

NaHCO3

calcium hydroxide

calcium carbonate

sodium hydroxide

potassium hydroxide

barium hydroxide

Ca(OH)2

CaCO3

sodium hydrogen carbonate

ionic compounds

hydroxide ion (OH)

Naming acids contain

(almost always as the )

Rules1. hydrogen becomes acid2. hydrogen becomes acid3. hydrogen becomes acid

______ide hydro____ic

______ate ______ic______ite ______ous

acids are always

eg) HCl(aq), H2SO4(aq), HNO3(s)

not an acid!!!

always hydrogenfirst element

aqueous (dissolved in water) (aq)

Examples

1.   HF(aq)

2.   H2SO3(aq)

3.   H3BO3(aq)

4.   HCl(g)

hydrofluoric acid

sulphurous acid

boric acid

not an acid!!!!!!!!

Writing Formulas use the naming acids rules in the

opposite direction come up with the then

the charges and add to the end

eg)

hydrosulphuric acid

carbonic acid

chlorous acid

hydrogen sulphide H2S(aq)

hydrogen carbonate H2CO3(aq)

hydrogen chlorite HClO2(aq)

“ionic” name write the formula, balance “(aq)”

M. States and Solubility acids –

elements –

molecular compounds –

ionic compounds – either look up on the

always (aq)

can be (s), (l) or (g)…

can be (s), (l), or (g)

solubility chart

look up state on

periodic table

the question usually tells you (or use common sense!)

(s) or (aq)…

Examples

1. LiCl( )

2. AgCl( )

3. NaNO3( )

4. Ba(OH)2( )

5. BaSO4( )

6. K2S( )

aq

s

aq

aq

aq

s

N. Chemical Reactions chemical reactions can cause both

and involve the formation of a

Evidence

1. temperature change2. colour change3. solid (precipitate) produced4. gas produced

physical and chemical changes always

new substance

states

states

balancing

Reactants Products

1 H2(g)+1 ZnCl2(aq)2 HCl(aq)1 Zn(s) +

that occur with chemical reactions can be:

1. endothermic = energy is

2. exothermic = energy is

reactants + energy products

reactants products + energy

energy changes

absorbed (enters)

released (exits)

2 SO3(g) + 197.8 kJ 2 SO2(g) + O2(g)

Mg(s) + ½ O2(g) MgO(s) + 601.6 kJ

O. Law of Conservation of Matter

Law of Conservation of Matter states thatmatter cannot be created or destroyed, it only changes forms

mass of reactants = mass of products

Counting Practice! How many of each element are in the following compounds?

1. NaCl 5. NH4CH3COO

2. BaBr2 6. 3 (NH4)2S

3.(NH4)3P 7. 2 CaCl2

4.Ba(OH)2 8. 8 PbI2

9. 4 Zn(CH3COO)2

there must be of each element on both sides of the reaction

are used to increase the number of atoms in a compound

when chemicals react they follow the Law of Conservation of Matter

equal numbers

coefficients(balancing)

P. Identifying Chemical Reactions

1. Hydrocarbon combustion

C?H? + O2(g) CO2(g) + H2O(g)

eg) CH4(g) + 2 O2(g) CO2(g) + 2 H2O(g)

a compound containing hydrogen and carbon (a hydrocarbon) burns/combusts in the presence of O2(g)

forms CO2(g) and H2O(g)

hydrocarbon

always

2. Simple Composition

element + element compound

eg) 2 Mg(s) + O2(g) 2 MgO(s)

combine to form a

will always have

are usually

more reactants than products

exothermic

elements compound

3. Simple Decomposition

compound element + element eg) 2 H2O(l) 2 H2(g) + O2(g)

will always have

are usually

compound

more products than reactants

endothermic

decomposes into its elements

4. Single Replacement

element + compound element + compound

eg) Cu(s) + 2 AgNO3(aq) 2 Ag(s) + Cu(NO3)2(aq)

an reacts with an to form a different element and a different ionic compound

Cl2(g) + 2 NaBr(aq) Br2(l) + 2 NaCl(aq)

element ionic compound

5. Double Replacement

compound + compound compound + compound

eg) Pb(NO3)2(aq) + 2 KI(aq) PbI2(s) + 2 KNO3(aq)

two ionic compounds react to form two different ionic compounds

Q. Significant Digits any digit from is significant

zeros areeg)

zeros areeg)

counted objects and constants are included in sig digs

zeros areeg)

1-9

trailing significant

“sandwich” significant

leading not significant

not

/ :

+/ :

multiply or divide then round answer to the

add or subtract then round answer to the

lowest number of sig digs

lowest number of decimal places

R. The Mole the mass of a single atom is so small that we

cannot easily measure it the is a concept that is used so

that we can actually measure the mass of elements and compounds…IT IS JUST A !!!

mole, n,

NUMBER

1. Avogadro’s Number

1 mole = atoms,

molecules etc.

you can use Avogadro’s number to calculate the number of moles in a substance if you know the number of molecules

Avogadro’s number = 6.022 1023

that’s 602 000 000 000 000 000 000 000 atoms, molecules etc!

n = # atoms NA

(NA)

Example

A diamond contains 5.0 x 1025 atoms of carbon. How many moles of carbon are in this diamond?

= 5.0 1025 atoms

6.02 1023 atoms/mol

= 83.056….mol

= 83 mol

n = # atoms NA

2. Atomic Molar Mass

the atomic masses given on the periodic table are an average of all the of each element

naturally occurring isotopes

is the

measured in

the atomic molar masses are given on the periodic table

molar mass mass of one mole of a substance

g/mol

Examples:Calculate the molar mass of the following substances:

22.99 g/mol

32.00 g/mol

Na(s) 1 x 22.99 g/mol =

O2(g) 2 x 16.00 g/mol =

44.01 g/mol

CO2(g)1 x 12.01 g/mol =

12.01 g/mol

2 x 16.00 g/mol =32.00 g/mol

Ca(OH)2(s)

101.96 g/mol

74.10 g/mol

Al2O3(s) 2 x 26.98 g/mol =

3 x 16.00 g/mol =

53.96 g/mol

48.00 g/mol

1 x 40.08 g/mol =

2 x 16.00 g/mol =

2 x 1.01 g/mol =

40.08 g/mol

32.00 g/mol

2.02 g/mol

now we can use number of moles and molar mass in a formula:

n = m M

where: n =

m =

M =

number of moles in mol

mass in g

molar mass in g/mol

m = nM

3. Mole Calculations

Example 1

How many moles are in 200 g of table salt (NaCl)? m =200 gM = 58.443 g/moln = ?

n = m M = 200 g 58.44 g/mol = 3.422… mol = 3.42 mol

Example 2

How many grams are in 62.9 mol of lead (II) nitrate?

n = 62.9 molM = 331.23 g/molm = ?

m = nM = (62.9 mol)(331.23 g/mol) = 20 834.367 g = 2.08 104 g

Pb(NO3)2