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