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Science 10 PreAP
Energy and Matter in Chemical Change
Understanding Matter
A. WHMIS
 workplace hazardous materials information system
 all chemicals are treated with respect
 WHMIS has been developed to provide guidelines for handling,
storage and disposal of reactive materials
compressed gas
corrosive
poisonous and
infectious material
causing immediate and
serious toxic effects
poisonous and
infectious causing
other toxic effects
dangerously
reactive material
flammable and
combustible
oxidizing material
biohazardous
infectious material
B. Properties
 physical properties are properties you can see and measure
eg) colour, state, density, ductility, malleability, boiling point
 chemical properties are properties used to describe how substances
will react with each other
eg) combustion, rusting, decomposition
C. Classification of Matter
All Matter
Mixtures
Homogeneous
Alloys
Solutions
Pure Substances
Elements
Compounds
Colloids
Metals
Ionic
Suspension
Non-Metals
Molecular
Mechanical
Mixture
Metalloids
Heterogeneous
 homogeneous mixture is a mixture of 2 or more substances that has
uniform properties (that appear as one)
eg) salt water, kool-aid, air
 an alloy is a homogeneous mixture of 2 or more metals
eg) brass = copper + zinc
steel = iron + chromium + carbon
 heterogeneous mixture is a mixture of 2 or more substances and
individual components are visible
eg) cereal, chocolate chip cookies, bread
 suspension is a mechanical mixture in which the components are in
different states
eg) mud
 colloids are mechanical mixtures in which the suspended substances
cannot be easily separated from the other substances
eg) milk
 elements are substances (metals, non-metals or metalloids) that are
pure and cannot be broken apart
eg) hydrogen, helium, carbon etc.
 compounds are two or more elements combined; cannot be
separated by physical means
eg) water, carbon dioxide etc.
 we will be focusing on elements and compounds in this course!
D. Atomic Structure
 atoms are the building blocks of ALL matter
 consist of a tiny nucleus and a huge “cloud” region
eg) ball bearing in middle of football field
nucleus
cloud region
 nucleus makes up 99% of the mass of an atom
 cloud region makes up most of the volume of an atom
Subatomic Particles
1. Proton (p+)
 positive charge
 found in nucleus
 determines the type of element
2. Neutron (n0)
 no charge
 found in nucleus
 used to hold nucleus together
3. Electron (e)
 smallest subatomic particle
 negative charge
 found in “cloud” region
 arranged in energy levels
 maximum # of electrons in each level (for first 20 elements only!):
Level 1 = 2 e
Level 2 = 8 e
Level 3 = 8 e
Level 4 = we will only go up to 2 e
Atomic Mass
 atomic mass unit (amu) is the mass of a proton or neutron =
1.7  1024 g
 net charge for all atoms is zero
 # e = # p+
 mass number is the sum of the protons and neutrons (rounded to
the nearest whole number)
 used to find the number of neutrons
eg) lithium
atomic number = 3
atomic mass = 6.94 = 7
# protons = 3
# neutrons = 7 – 3 = 4
Isotopes
 atoms that have the same number of protons but a different number
of neutrons
 atomic mass on the periodic table is an average mass based on the
percentage abundances of all naturally occurring isotopes of the
element
isotope notation:
A
X
Z
X = symbol
A = mass # (#p+ + n°)
Z = atomic # (#p+)
Examples
copper - 64
64
Cu
29
# p+ = 29
# e- = 29
#n° = 64 - 29 = 35
copper - 62
62
Cu
29
# p+ = 29
# e- = 29
#n° = 62 - 29 = 33
Your Assignment: pg 1
E. Atomic Theory
1. Dalton’s Atomic Theory: 1808
 all matter is composed of tiny, indivisible particles called atoms
 atoms of an element have identical properties
 atoms of different elements have different properties
 atoms of two or more elements can combine in constant ratio to
form new substances
eg)
H:O ratio 2:1 → H2O → water
H:O ratio 2:2 → H2O2 → hydrogen peroxide
2. J.J. Thompson: 1897
 credited with discovery of electrons
 “raisin bun” model or “plum pudding” model
 atom is a sphere which is positive, with negative electrons
embedded in it like raisins in a bun
 most of the mass is associated with the positive charge
3. Ernest Rutherford: 1911
 atoms have a nucleus which is positive and has most of the mass
 most of the atom is empty space occupied by the moving negatively
charged electrons
 proposed the existence of protons
4. Neils Bohr: 1913
 electrons move in circular orbits around the nucleus
 cannot exist between orbits
5. James Chadwick: 1932
 showed that the nucleus must contain heavy neutral particles to
account for all of the atom’s mass (neutrons)
6. Schrodinger/de Broglie: 1930
 quantum mechanical model
 electrons have distinct energy levels
 exact locations of electrons are not defined, but the probable location
in a region of space can be predicted
F. The Periodic Table
 the periodic table was developed by Dmitri Mendeleev in the mid
1800’s
http://www.chemicool.com/
1. Atomic Number
 number of protons in one atom of an element
 increases from left to right and top to bottom
2. Properties
 3 major categories:
1. metals are good conductors, strong, malleable (pound into thin
sheet), ductile (can draw into a wire, bendable), have high
luster; are found on left side of stair case
2. non metals are poor conductors, non-lustrous, weak,
etc…opposite properties to metals; found on right side of
staircase
3. metalloids show properties of both metals and nonmetals; found
above and below staircase
3. Groups
 18 vertical columns are called groups or families
 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







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 Elements
elements in each group share similar chemical properties
(reactivity) although intensity changes
 reactivity increases down group for metals and up group for
nonmetals
 group number indicates how many electrons are in the outermost
energy level
4. Periods
 horizontal rows
 show a trend in reactivity which changes from left to right
 each time you move to a new period you start the trend over
G. Electron Energy Level Representations (EELR)
 nucleus – shows # p+ and n0
 energy levels – shows # of e in each level (# of levels = period element
is in)
 valence electrons are the e in outermost energy level (same as group
#, ignoring the 1 in front of groups 13-18 )
Examples
sodium
- atomic # 11
- mass #
22.99 = 23
+
# p = 11
# e- = 11
#n° = 23 - 11 = 12
1 e8 e2 e-
11 e-
p+ = 11
n° = 12
argon
- atomic # 18
- mass #
39.948 = 40
# p+ = 18
# e- = 18
#n° = 40 – 18 = 22
8 e8 e2 ep+ = 18
n° = 22
18 e-
Draw the EELR for the following:
1. potassium
2. chlorine
3. beryllium
4. calcium
potassium - atomic # 19
- mass # 39.10 = 39
# p+ = 19
# e- = 19
#n° = 39 – 19 = 20
1 e8 e8 e2 e-
chlorine
- atomic # 17
- mass # 35.453 = 35
# p+ = 17
# e- = 17
#n° = 35 – 17 = 18
7 e8 e2 e-
19 e-
p+ = 19
n° = 20
p+ = 17
n° = 18
beryllium - atomic # 4
- mass # 9.012 = 9
# p+ = 4
# e- = 4
#n° = 9 – 4 = 5
2 e2 ep+ = 4
n° = 5
17 e-
calcium
- atomic # 20
- mass # 40.08 = 40
# p+ = 20
# e- = 20
#n° = 40 – 20 = 20
2 e8e8 e2 e-
4 ep+ = 20
n° = 20
20 e-
H. Octet Rule
 atoms tend to be stable when the outer energy level is full of
electrons
 the octet rule states that atoms bond in a way to have a full valence
energy level (there are exceptions)
 atoms will either share electrons, or gain or lose electrons in order
to satisfy this octet rule
 compounds are formed when two or more different elements bond
together either by sharing or transferring electrons
I. Ions
 ions are atoms or groups of atoms that have a net charge
 number of p+ and e are not equal
 most atoms try to achieve the electron configuration of a noble gas
 isoelectronic means having the same number of e as another atom
or ion
eg) fluorine gains an electron to be isoelectronic with neon
eg) potassium loses an electron to be isoelectronic with argon
Cations
 positively charged ions
 lost electrons to obtain a stable electron configuration (full valence
level)
 METALS form cations
eg) sodium loses one e to completely empty the last energy level
Na+ has 10 e and 11 p+
 charge on a metal ion is the same as the group number for groups
1,2,3,13,14 (ignore the 1 in front of 13 and 14)
Anions
 negatively charged ions
 gained electrons to obtain a stable electron configuration
 NON-METALS form anions
 have “ide” ending
eg) oxygen gains two e to completely fill the last energy level
O2 has 10 e and 8 p+ and is called the oxide ion
 charge on a non-metal is 18  group number (not Roman Numerals)
EELR’s for Ions
number of p+ = atomic number
number of e = number of p+ – charge
number of n0 = atomic mass – atomic number (# of p+)
Examples
nitride ion - atomic # 7
- mass # 14.007 = 14
# p+ = 7
# e = 7-(-3) = 10
#n° = 14 – 7 = 7
8 e2 e-
10 e-
p+ = 7
n° = 7
calcium ion
- atomic # 20
- mass # 40.08 = 40
# p+ = 20
# e- = 20 – (+2) = 18
#n° = 40 – 20 = 20
8 e8 e2 ep+ = 20
n° = 20
18 e-
Draw the EELR for the following:
1. sulphide ion
2. aluminium ion
3. chloride ion
4. magnesium ion
sulphide ion
- atomic # 16
- mass # 32.07 = 32
# p+ = 16
# e- = 16 – (2) = 18
#n° = 32 – 16 = 16
8 e8 e2 e-
- atomic # 13
- mass # 26.98 = 27
# p+ = 13
# e- = 13 – (+3) = 10
#n° = 27 – 13 = 14
18 e-
p+ = 16
n° = 16
chloride ion
aluminium ion
8 e2 e-
10 e-
p+ = 13
n° = 14
- atomic # 417
- mass # 35.45 = 35
# p+ = 17
# e- = 17 – (1) = 18
#n° = 35 – 17 = 18
8 e8 e2 e-
magnesium ion
- atomic # 12
- mass # 24.31 = 24
# p+ = 12
# e- = 12 – (+2) = 10
#n° = 24 – 12 = 12
18 e-
p+ = 17
n° = 18
Your Assignment: pg 2
8 e2 ep+ = 12
n° = 12
10 e-
J. Elements
 metallic elements exist as single atoms (monatomic)
 chemical formula is simply the symbol followed by the state at room
temperature
eg) sodium Na(s)
mercury Hg()
copper Cu(s)
 non-metals (not including noble gases) do not exist as single atoms
and are called molecular elements (diatomic, polyatomic)
 chemical formula is the symbol with the subscript and the state at
room temperature
 memorize the subscripts (flagpole)!!!
H2
N2
O2
F2
P4
S8
Cl2
Br2
I2
Summary:
Monatomic C(s), noble gases, all metals
Diatomic
H2(g), N2(g), O2(g), F2(g), Cl2(g), Br2(g), I2(s)
Polyatomic P4(s), S8(s)
K. Molecular Compounds
 molecular compounds are formed when two or more nonmetals
bond together
 bonded by covalent bonds which is the force of attraction between
atoms that are sharing electrons
 properties:
1. do not conduct electricity when dissolved in water
2. dissolve in water to form either a neutral molecular solution or
an acidic solution
3. solids, liquids or gases at room temperature
Naming
 give the atom name for the first element (with the prefix if there is
more than one) then give the name for the second element with
“ide” ending and include the prefix
 Note: if the first element is hydrogen, do not put a prefix (these are
acids!)
Prefixes
1 = mono
2 = di
3 = tri
Try These
1. CO(g)
2. CO2(g)
3. P4O10(s)
4. BrH7(s)
4 = tetra
5 = penta
6 = hexa
7 = hepta
8 = octa
9 = nona
10 = deca
carbon monoxide
carbon dioxide
tetraphosphorus decaoxide
bromine heptahydride
Writing Formulas
 simply write each symbol followed by the subscript (from prefix)
Try These
1. oxygen dibromide
2. diphosphorus pentasulphide
3. carbon tetraiodide
4. phosphorus pentachloride
OBr2
P2S5
CI4
PCl5
 some molecular compounds have classical names…memorize them!!!
O3 = ozone
NH3 = ammonia
H2O = water
H2O2 = hydrogen peroxide
CH4 = methane
CH3OH = methanol
C2H6 = ethane
C2H5OH = ethanol
C6H12O6 = glucose
C12H22O11 = sucrose
H2S = hydrogen sulphide
HCl, HF, HBr, HI = hydrogen halides (usually acids)
Your Assignment: pg 3
L. Ionic Compounds
 ionic compounds are formed when electrons are transferred,
allowing oppositely charged ions to bond together
 ionic bond is the force of attraction between oppositely charged ions
 properties:
1. conduct electricity when dissolved in water
2. separate into ions when dissolved in water
3. crystalline solids at room temperature
1. Monovalent Ionic Compounds
 monovalent means there is one charge on the metal
eg) Na+, Ca2+
 metal + nonmetal
Naming
 give the name for each ion…metals are normal, non-metals have
“ide” ending
*** never use capital letters in naming!!
eg) NaF sodium fluoride
Na2S sodium sulphide
the 2 means that two sodium ions are bonded with one sulphide
ion… this doesn’t matter for naming
Try These:
Name the following:
1. LiF
lithium fluoride
2. KCl
potassium chloride
3. BeS
beryllium sulphide
4. Rb3P
rubidium phosphide
5. MgF2
magnesium fluoride
6. Na2O
sodium oxide
7. CsBr
cesium bromide
8. BaCl2
barium chloride
Writing Formulas
 look up the symbol for each ion and write them listing the metal ion
first
 balance the charges using subscript numbers
***in ionic compounds, the total positive charges must equal the
total negative charges…the net charge is zero
eg) sodium oxide
Na2O
1+
1+  2 = 2+
2
2  1 = 2
calcium phosphide
2+
2+  3 = 6+
Ca3P2
3
3  2 = 6
Try these:
1. magnesium chloride
2. calcium chloride
3. zinc sulphide
4. silver sulphide
5. germanium oxide
6. calcium arsenide
7. magnesium nitride
MgCl2
CaCl2
ZnS
Ag2S
GeO2
Ca3As2
Mg3N2
2. Multivalent Ionic Compounds
 metal ions that have more than one possible charge
eg) Cu2+, Cu+, Fe3+, Fe2+
 transition metal + nonmetal
 the first charge listed is the most common
Naming
 same rules as before
 the difference…you must include brackets containing the charge of
the metal ion in Roman Numerals (I,II,III,IV,V,VI,VII)
 figure out the charge on the metal by using how many negative
charges there are in the nonmetal ions
eg) CuI
copper (I) iodide
TiBr4
titanium (IV) bromide
Ti3P4
titanium (IV) phosphide
Try These:
1. AuBr
2. CrCl2
3. Co2O3
4. VS2
5. PuN2
gold (I) bromide
chromium (II) chloride
cobalt (III) oxide
vanadium (IV) sulphide
plutonium (VI) nitride
Writing Formulas
 same rules
 the charge on the metal is given to you in the brackets
eg) iron (II) oxide
FeO
tin (II) chloride
SnCl2
chromium (III) sulphide
Cr2S3
Try these:
1. chromium (II) sulphide
2. nickel (III) chloride
3. vanadium (IV) phosphide
4. gold (III) iodide
CrS
NiCl3
V3P4
AuI3
Your Assignment: pg 4
3. Mixed Ionic Compounds
 metal ion + polyatomic ion (complex ion)
eg) PO43, SO42, HCO3 etc.
***NH4+ (ammonium ion) is the only positive complex ion…you’ll
see it in the place of a metal
Naming
 give the name for the first ion then give the name for the complex ion
Try These:
1. KIO3
2. NaCH3COO
3. MgSO3
4. NH4NO3
5. Ca3(PO4)2
potassium iodate
sodium acetate
magnesium sulphite
ammonium nitrate
calcium phosphate
Writing Formulas
 same as before…look up the symbol for each ion then balance the
charges using subscripts
 if you must multiply one of the complex ions, put brackets around it
first then write the subscript
Try These:
1. aluminum phosphate
2. aluminum chlorate
3. calcium sulphite
4. scandium acetate
5. ammonium sulphate
AlPO4
Al(ClO3)3
CaSO3
Sc(CH3COO)3
(NH4)2SO4
Your Assignment: pg 5
4. Hydrated Ionic Compounds
 ionic compounds containing water in their structure
 water is represented by “xH2O” in the formula where x is the
number of water molecules
Naming
 give the name for the first part of the compound using our previous
rules
 name the “xH2O” part as prefix + “hydrate” (same prefixes as
molecular compounds)
Try These:
1. NaF3H2O
2. FeSO46H2O
3. GaNH2O
sodium fluoride trihydrate
iron (II) sulphate hexahydrate
gallium nitride monohydrate
Writing Formulas
 same as before…look up the symbol for each ion then balance the
charges using subscripts
 for the hydrate part…add “xH2O” where x is the number given in
the prefix
Try These:
1. iron (III) nitrate nonahydrate
2. sodium chlorate tetrahydrate
3. nickel (II) sulphite heptahydrate
Fe(NO3)39H2O
NaClO34H2O
NiSO37H2O
Your Assignment: pg 6
Summary: Ionic vs. Molecular
Ionic
 metal + non-metal (or
polyatomic ions)
 no prefixes (except hydrates)
 solids
 solutions conduct
 solutions are basic or neutral
Molecular
 all non-metals




prefixes
solids, liquids or gases
solutions do not conduct
solutions are acidic or neutral
5. Acids and Bases
 matter can be subdivided into three groups based on its properties:
1. acids
2. bases
3. neutral substances
 let’s look at the properties of acids and bases:
Acids
 usually soluble in H2O
 conduct electricity
 neutralize bases
 taste sour
 react with metals to produce
H2(g)
Indicators
 litmus – pink
 bromothymol blue – yellow
 phenolphthalein – colourless
Indicators
 litmus – blue
 bromothymol blue – blue
 phenolphthalein – bright pink
eg) vinegar, lemon juice
eg) ammonia, Tums





Bases
usually soluble in H2O
conduct electricity
neutralize acids
taste bitter
feel slippery
 the pH scale was devised to indicate how acidic or basic a substance
is
stomach
acid
strong acids
0
coffee
1
5
antacid
7
acids
11
bases
Neutral
Naming and Writing Formulas for Bases
 use rules for ionic compounds
 many of them contain the hydroxide ion (OH)
eg) NaOH
sodium hydroxide
KOH
potassium hydroxide
Ba(OH)2
barium hydroxide
NaHCO3
sodium hydrogen carbonate
calcium hydroxide
Ca(OH)2
calcium carbonate
CaCO3
Naming Acids
 acids always contain hydrogen (usually as the first element)
 acids are always aqueous (dissolved in water) (aq)
eg) HCl(aq), H2SO4(aq), HNO3(s)
not an acid
Rules
hydrogen ________ide becomes hydro_______ic acid
hydrogen ________ate becomes _______ic acid
hydrogen ________ite becomes _______ous acid
Try These:
Write the acid name for the following:
1. HF(aq)  hydrogen fluoride  hydrofluoric acid
2. H2SO3(aq)  hydrogen sulphite  sulphurous acid
3. H3BO3(aq)  hydrogen borate  boric acid
4. HCl(g)  hydrogen chloride  not an acid!!!!!!!!
KOH
14
strong bases
Writing Formulas
 use the naming acids rules in the opposite direction
 come up with the “ionic” name then write the formula, balance the
charges and add “(aq)” to the end
Try These:
Write the ionic name and formula for the following:
1. hydrosulphuric acid
hydrogen sulphide
2. carbonic acid
hydrogen carbonate
3. chlorous acid
hydrogen chlorite
H2S(aq)
H2CO3(aq)
HClO2(aq)
***Note the special cases of sulphur and phosphorus…they add “ur” and
“or” respectively to the name
Your Assignment: pgs 7-8
M. States and Solubility
 acids – always (aq)
 elements – can be (s), (l) or (g)…look up state on periodic table
 molecular compounds – can be (s), (l), or (g)…the question usually
tells you (or use common sense!)
 ionic compounds – either (s) or (aq)…look up on the solubility chart
Examples
1. LiCl (aq)
2. AgCl(s)
3. NaNO3(aq)
4. Ba(OH)2(aq)
5. BaSO4(s)
6. K2S(aq )
Your Assignment: pg 9
N. Chemical Reactions
 chemical reactions can cause both physical and chemical changes
and always involve the formation of a new substance
Evidence
1. temperature change
2. colour change
3. solid (precipitate) produced
4. gas produced
Reactants
Products
balancing
1 Zn(s) + 2 HCl(aq)  1 ZnCl2(aq) + 1 H2(g)
states
states
 energy changes that occur with chemical reactions can be:
1. endothermic = energy is absorbed (enters)
eg) cooking food
reactants + energy  products
2. exothermic = energy is released (exits)
eg) combustion
reactants  products + energy
O. Law of Conservation of Matter
 Law of Conservation of Matter states that matter 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
4. Ba(OH)2
7. 2 CaCl2
2. BaBr2
5. NH4CH3COO
8. 8 PbI2
3. (NH4)3P
6. 3 (NH4)2S
9. 4 Zn(CH3COO)2
 when chemicals react they follow the Law of Conservation of Matter
 there must be equal numbers of each element on both sides of the
reaction
 coefficients are used to increase the number of atoms in a compound
(balancing)
Examples
1.
2 Mg(s) + 1 O2(g)  2 MgO(s)
 Mg
O
1x2=2
2
Mg
O
1 2
1x2=2 
2 H2O(l)  2 H2(g) + 1 O2(g)
2.

3.
Cu 1
Ag 1 2
NO3 1 x 2 = 2
Cl
Na
Br
2
1 2
1x2=2
2
2x2=2 
Cu 1
Ag 1 x 2 = 2 
NO3 2
Cl
Na
Br
1 2
1x2=2 
2
2 KI(aq) + 1 Pb(NO3)2(aq)  1 PbI2(s) + 2 KNO3(aq)
 K
I
NO3
Pb
6.
H
O
1 Cl2(g) + 2 NaBr(aq)  1 Br2(l) + 2 NaCl(aq)

5.
2x2=2
1 2
1 Cu(s) + 2 AgNO3(aq)  2 Ag(s) + 1 Cu(NO3)2(g)

4.
H
O
1x2=2
1 2
2
1
K
I
NO3
Pb
1 2
2
2x2=2 
1
1 CH4(g) + 2 O2(g)  1 CO2(g) + 2 H2O(g)

C
H
O
1
4
2x2=4
C
H
O
Your Assignment: pg 10
1
2x2=4 
2+1=3
2+2=4
P. Identifying Chemical Reactions
1. Hydrocarbon Combustion
 a compound containing hydrogen and carbon (a hydrocarbon)
burns/combusts in the presence of O2(g)
 always forms CO2(g) and H2O(g)
C?H? + O2(g)  CO2(g) + H2O(g)
eg) 1 CH4(g) + 2 O2(g)  1 CO2(g) + 2 H2O(g)
2. Composition, Combination or Formation
 will always have more reactants than products
 are usually exothermic
i. Combination of Elements
 elements combine to form a compound
element + element  compound
eg) 2 Mg(s) + 1 O2(g)  2 MgO(s)
ii. Combination to Produce an Acid (AP)
 combination of a non-metal oxide (acidic oxide) with water
non-metal oxide + water  acid
eg) 1 CO2(s) + 1 H2O()  1 H2CO3(aq)
iii. Combination to Produce a Base (AP)
 combination of a metal oxide (basic oxide) with water
metal oxide + water  base
eg) 1 Li2O(s) + 1 H2O()  2 LiOH(aq)
iv. Combination to Produce a Salt (AP)
 combination of a metal oxide (basic oxide) with a non-metal oxide
(acidic oxide)
metal oxide + non-metal oxide  a salt
eg) 1 MgO(s) + 1 CO2(g)  1 MgCO3(s)
3. Decomposition
 will always have more products than reactants
 are usually endothermic
i. Decomposition to Produce Elements
 compound decomposes into its elements
compound  element + element
eg) 2 H2O(l)  2 H2(g) + 1 O2(g)
ii. Decomposition of Acids (AP)
 acid decomposes into non-metal oxide (acidic oxide) and water
acid  non-metal oxide + water
eg) 1 H2SO4(aq)  1 SO3(g) + 1 H2O(l)
iii. Decomposition of Bases (AP)
 base decomposes into metal oxide (basic oxide) and water
base  metal oxide + water
eg) 1 Ca(OH)2(s)  1 CaO(s) + 1 H2O(l)
iv. Decomposition of Salts (AP)
 a salt decomposes into non-metal oxide (acidic oxide) and metal
oxide (basic oxide)
a salt  non-metal oxide + metal oxide
eg) 1 Fe2(SO4)3(s)  3 SO3(g) + 1 Fe2O3(s)
4. Single Replacement (Displacement)
 an element reacts with an ionic compound to form a different
element and a different ionic compound
element + compound  element + compound
eg) Cu(s) + 2 AgNO3(aq)  2 Ag(s) + Cu(NO3)2(g)
Cl2(g) + 2 NaBr(aq)  Br2(l) + 2 NaCl(aq)
5. Double Replacement
 two ionic compounds react to form two different ionic compounds
compound + compound  compound + compound
i. Precipitation Reactions
 one of the products formed is insoluble
compound + compound  insoluble compound + compound
eg) KI(aq) + Pb(NO3)2(aq)  PbI2(s) + KNO3(aq)
ii. Gas Formation (AP)
 sometimes one of the products in a double replacement undergoes
spontaneous decomposition, forming a gas
 you must write down the products of this decomposition
 memorize these!
H2CO3(aq)  H2O() + CO2(g)
NH4OH(aq)  H2O() + NH3(g)
H2SO3(aq)  H2O() + SO2(g)
eg) CaCO3(s) + HNO3(aq)  Ca(NO3)2(aq) + H2O() + CO2(g)
iii. Neutralization Reactions
 an acid reacts with a base to produce water and a salt
acid + base  water + a salt
eg) 1 HCl(aq) + 1 NaOH(aq)  1 H2O(l) + 1 NaCl(aq)
1 ZnO(s) + 2 HNO3(aq)  1 H2O(l) + 1 Zn(NO3)2(aq)
2 LiOH(aq) + 1 CO2(g)  1 H2O(l) + 1 Li2CO3(aq)
Your Assignment: pgs 11-13
Q.





Significant Digits
any digit from 1 – 9 is significant
trailing zeros are significant eg) 6.3800, 12 000
“sandwich” zeros are significant eg) 2.04, 1005.002
leading zeros are not significant eg) 0.0032
counted objects and constants are not included in sig digs
 / : multiply of divide then round answer to the lowest number of sig
digs
 +/ : add or subtract then round answer to the lowest number of
decimal places
Your Assignment: pg 14
R. The Mole
 the mass of a single atom is so small that we cannot easily measure it
 the mole, n, is a measurement that is used so that we can actually
measure the mass of elements and compounds…IT IS JUST A
NUMBER!!!
1. Avogadro’s Number
 1 mole = Avogadro’s number (NA) = 6.02 x 1023 atoms, molecules
etc. that’s 602 000 000 000 000 000 000 000 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
n = # atoms
NA
Example
A diamond contains 5.0 x 1025 atoms of carbon. How many moles of carbon
are in this diamond?
n = 5.0  1025 atoms
6.02  1023 atoms/mol
= 83.056….mol
= 83 mol
Your Assignment: pg 15 #1-3
2. Atomic Molar Mass
 12C, carbon 12, has been assigned by chemists (1961) as the standard
for measuring atomic masses
 12C is assigned a mass of exactly 12 atomic mass units (amu)
 comparing masses of atoms involves the use of a mass spectrometer
 an electron beam and magnetic field are used to deflect ions…amount
of deflection indicates the mass (higher mass = lower deflection)
 once you know the mass of each isotope, they can then be averaged
(weighted) to come up with the average atomic mass for the element
 the atomic masses given on the periodic table are an average of all the
naturally occurring isotopes of each element
Example 1
It is known that natural carbon is composed of 98.89% 12C at 12 amu and
1.11% 13C at 13.003355 amu. The amount of 14C is so small it is not used
in these calculations. Calculate the average atomic mass of natural carbon.
atomic mass = 98.89% of 12 amu + 1.11% of 13.003355 amu
= (0.9889)(12 amu) + (0.0111)(13. 003355 amu)
= 12.01113724 amu
= 12.0 amu
Example 2
When a sample of natural copper is vaporized and injected into a mass
spectrometer, the sample is found to be 69.09% 63Cu with a mass of 62.93
amu and 30.91% 65Cu with a mass of 64.93 amu. Calculate the average
atomic mass of copper.
atomic mass = 69.09% of 62.93 amu + 30.91% of 64.93 amu
= (0.6909)(62.93 amu) + (0.3091)(64.93 amu)
= 63.5482 amu
= 63.55 amu
Your Assignment: pg 15 #4-5
 molar mass is the average atomic mass of one mole of a substance
 measured in g/mol
 the average atomic molar masses are given on the periodic table
Examples:
Calculate the molar mass of the following substances:
1. Na(s)
1 x 22.99 g/mol = 22.99 g/mol
2. O2(g)
2 x 16.00 g/mol = 32.00 g/mol
3. Al2O3(s)
2 x 26.98 g/mol = 53.96 g/mol
3 x 16.00 g/mol = 48.00 g/mol
101.96 g/mol
4. Ca(OH)2(s) 1 x 40.08 g/mol = 40.08 g/mol
2 x 16.00 g/mol = 32.00 g/mol
2 x 1.0079 g/mol = 2.0158 g/mol
74.10 g/mol
3. Mole Calculations
 now we can use number of moles and molar mass in a formula:
n=m
M
where:
m = nM
n = number of moles in mol
m = mass in g
M = molar mass in g/mol
Example 1:
How many moles are in 200 g of table salt (NaCl)?
m = 200 g
n=m
M = 1 x 22.99 g/mol = 22.99 g/mol
M
1 x 35.453 g/mol = 35.453 g/mol
=
200 g
58.443 g/mol
58.443 g/mol
= 3.42 mol
Example 2:
How many grams are in 62.9 mol of lead (II) nitrate? Pb(NO3)2
n = 62.9 mol
m = nM
M = 1 x 207.19 g/mol = 207.2 g/mol
= (62.9 mol)(331.2 g/mol)
2 x 14.01 g/mol = 28.02 g/mol
= 2.08  104 g
6 x 16.00 g/mol = 96.00 g/mol
331.2 g/mol
Your Assignment: pg 15-17 #4-14
S. Gravimetric Stoichiometry
 gravimetric = mass measurements
 when you balance a reaction, you are balancing the moles of each
reactant and product…this gives you the mole ratio
eg) 2 Mg(s) + 1 O2(g)  2 MgO(s)
mole ratio is 2:1:2
 certain assumptions must be made:
1.
the reaction is spontaneous
2.
the reaction is fast
3.
the reaction goes to completion
4.
the reaction is stoichiometric ie) whole number ratio of
moles of reactants and products
Steps
1.
2.
3.
4.
Write a balanced equation including the states. Write the
information given.
Find the moles of the given species using n=m/M.
Find the moles of the wanted species using mole ratio
(wanted/given).
Calculate mass of the wanted species using m=nM.
Example 1
Calculate the moles of aluminum formed when 2.50 moles of aluminum
oxide decomposes into aluminum and oxygen.
2 Al2O3(s)

3 O2(g)
+ 4 Al(s)
x mol
n = 2.50 mol
n = 2.50 mol  4
2
= 5.00 mol
iron (III)
ion atomic #
26
mass #
55.85
Your Assignment: pg 18
Example 2
Calculate the mass of sulfur needed to react with 0.500 moles of zinc.
8 Zn(s)
+
1 S8(s) 
8 ZnS(s)
xg
n = 0.500 mol
n = 0.500 mol  1
8
= 0.0625 mol
M = 256.56 g/mol
m = nM
= (0.0625 mol)(256.56 g/mol)
= 16.035 g
= 16.0 g
Your Assignment: pg 19
Example 3
Iron is produced by the reaction of iron(III) oxide with carbon monoxide to
produce iron and carbon dioxide. What mass of iron(III) oxide is required to
produce 1000 g of iron?
Fe2O3(s)
+
3CO(g)
 2Fe(s)
+
3CO2(g)
xg
M = 159.70 g/mol
m = 1000 g
n = 17.90… x 1/2
M = 55.85 g/mol
= 8.95… mol
n = m/M
m = nM
= 1000 g/55.85g/mol
= (8.95… mol)(159.70 g/mol)
= 17.90… mol
= 1429.72… g
= 1430 g
Your Assignment: pgs 20-21
Chemistry Review: pgs 22-24