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Chemistry 20
Unit 0: Chemistry Review
1
Chemistry 20
Unit 0: Chemistry Review
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

MSDS is a material safety data sheet
2
Chemistry 20
Unit 0: Chemistry Review
The Atom

proton  p+  positive charge

neutron  n0  zero charge

electron  e-  negative charge

electrons are found in a cloud region around the nucleus

nucleus contains the proton and neutrons which make up most of
mass of the atom

mass number =  protons +  neutrons
 neutrons = mass number -  protons

isotope = atoms with the same  of protons but a different  of neutrons (different mass
numbers)
Examples:
carbon-12 
carbon-14 
3
Chemistry 20
Unit 0: Chemistry Review
Periodic Table

arranged in groups (columns) and periods (rows)

group number = number of outer level (valence) e-

period number = number of energy levels occupied by eExample:
Na
Group Number = 1
Period Number = 3
This information is helpful for drawing Energy Level Diagrams
Energy Level Diagrams

atoms are electrically neutral which means that the  of protons =  of electrons

maximum number of e-:
3rd level = 8 e2nd level = 8 e1st level = 2 e-
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Chemistry 20
Unit 0: Chemistry Review
Examples:
-
Lesson #R1: Review of Atomic Structure
Draw the energy level diagrams for each of the following elements.
1. hydrogen atom
6. aluminum atom
2. carbon atom
7. phosphorus atom
3. helium atom
8. chlorine atom
4. oxygen atom
9. argon atom
5. sodium atom
10. calcium atom
5
Chemistry 20
Unit 0: Chemistry Review
11. What is the relationship between group number and number of valence (outermost)
electrons?
12. What is the relationship between period number and the number of energy levels
occupied by electrons?
Ions

ions are particles or groups of particles that have a net charge (either positive of negative)

neutral atoms are unstable if their valence level is not full

atoms will strive to satisfy the octet rule in order to become stable…in other words, they
strive to have a full valence level and do so by giving away or taking e-
Metals
- metals  give away e- and become positive ions

Na+, Ca2+, Fe3+

cations: positively charged ions have fewer electrons than protons

naming: metal name gets ion added to it
 sodium turn to sodium ion
Example:
8 e-
Chemistry 20
Unit 0: Chemistry Review
Na
a
Na+
+
+
electron
e-
Chemistry 20
Unit 0: Chemistry Review
Non-Metal
- non-metals  take e- and become negative ions

Cl-, P3-, O2-

anions: negatively charged ions have more electrons than protons

naming: non-metal name has –ine dropped and –ide added
 chlorine turn to chlroide
Example:
A chloride ion, Cl-
Chemistry 20
Unit 0: Chemistry Review
Cl
+
electron
+
e-
Cl-
Lesson #R2: Review of Ionic Structure
Draw the energy level diagrams for each of the following ion.
1. lithium ion
6. nitride ion
2. fluoride ion
7. sodium ion
3. aluminum ion
8. sulphide ion
4. chloride ion
9. calcium ion
5. magnesium ion
10. oxide ion
Chemistry 20
Unit 0: Chemistry Review
11. What is the relationship between the electron configuration of an ion of one of the representative
elements and the electron configuration of the nearest noble gas?
12. What problem arises when trying to predict the charge on an ion in Group 14?
Lesson #R3: Review of Atoms vs. Ions
Complete the following chart:
Name
Example: calcium ion
Symbol
# of Protons
# of Electrons
Net Charge
Ca2+
20
18
2+
1. oxygen atom
2. fluoride ion
3.
C
4.
Cl
5.
12
2+
6.
16
2
7.
18
1+
8.
10
0
9.
Ba2+
10. helium atom
11.
H+
12.
7
13.
Fe3+
14.
Sn4+
10
Chemistry 20
Unit 0: Chemistry Review
15. sodium ion
Al3+
16.
17.
29
18.
2+
54
1
54
1+
19. gold atom
20.
Elements

metals exist as single atoms
Example: Li(s), Cu(s), Hg(l)

nonmetals and hydrogen do not exist as single atoms – 7 starting at 7 that make a 7 or flag pole!
H2
N2
O2
F2
P4
S8 Cl2
Br2
I2
Practice:
1. Cu(s) =
4. fluorine gas =
2. O2(g) =
5. barium =
3. Al(s) =
6. nitrogen gas =
Ionic Compounds
Cl
Cl
Chemistry 20
Unit 0: Chemistry Review

metals + non-metals or polyatomic ions

monovalent (K+, Be2+) or polyatomic metals (Fe3+, Fe2+)

charges on the ions are the result of taking or giving e-

to go from formula to name: name of first ion, then
brackets for charge if multivalent in roman numerals,
then name for the second ion
- first element (
) second element-ide
Examples:
AlCl3 =
Fe2O3 =
Practice:
1. Zn3P2 =
2. NaNO3 =
3. NiF3 =
4. MnO2 =
5. Cr2(SO4)3 =

to go from name to formula: write the symbol for each ion, then add subscripts to balance charges

remember to reduce to lowest whole numbers
Examples:
calcium sulphide =
iron (II) hydroxide =
Chemistry 20
Unit 0: Chemistry Review
Practice:
1. lithium bromide =
4. ammonium sulphate =
2. sodium phosphate =
5. calcium phosphate =
3. magnesium nitride =
Hydrated Compounds

ionic compounds containing water in their structure

water is represented by “ x H2O” in the formula where x is the number of water molecules

prefixes:
1 = mono
6 = hexa
2 = di
7 = hepta
3 = tri
8 = octa
4 = tetra
9 = nona
5 = penta
10 = deca
Chemistry 20
Unit 0: Chemistry Review

to go from formula to name: give the ionic name for the first part of the compound, then name the “
x H2O” part as prefix + “hydrate”
Examples:
NaF3H2O =
CuSO45H2O =

to go from name to formula: first part is the same as before…look up the symbol for each ion then
balance the charges using subscripts, then for the hydrate part…add “ x H2O” where x is the
number given in the prefix
Examples:
iron (III) nitrate nonahydrate =
sodium chlorate tetrahydrate =
nickel (II) sulphite heptahydrate =
Lesson #R4: Review of Elements and Ionic Nomenclature
Complete the following chart:
Formula
1.
IUPAC Name
CdO(s)
2.
sodium fluoride
3.
chlorine gas
4.
AlP 5 H2O(s)
copper (II) oxide
5.
6.
Mg(OH)2(s)
7.
Na2CO3 3 H2O(s)
magnesium sulphate nonahydrate
8.
9.
N2(g)
10.
lithium chloride
11.
sodium chlorate
12.
13.
K3PO4(s)
calcium metal
Chemistry 20
Unit 0: Chemistry Review
nickel (III) bromide
14.
15.
MnO2(s)
16.
ammonium sulphite
17.
zinc sulphide
18.
NaHSO3(s)
ammonium sulphate
19.
20.
Au(s)
copper (II) chloride
21.
22.
SnF2(s)
23.
phosphorus
24.
sodium hypochlorite
25.
KMnO4(s)
26.
SrF2(s)
27.
RbCl(s)
28.
Li2O(s)
29.
iron (III) sulphide
30.
zinc chloride
31.
aluminum sulphide
32.
CoCl2(s)
33.
Au(NO3)3(s)
34.
Cu2O(s)
35.
lead (IV) acetate trihydrate
36.
chromium (II) oxide
37.
magnesium iodide
38.
KC6H5COO(s)
39.
Na2S2O3(s)
40.
NH4HCO3(s)
41.
ammonium sulphide
42.
barium sulphite
43.
magnesium hydroxide
Chemistry 20
Unit 0: Chemistry Review
44.
FeSO4 4 H2O(s)
45.
LiCl 2 H2O(s)
sodium phosphate decahydrate
46.
47.
TiO2(s)
48.
bismuth (V) sulphate
49.
tin (IV) sulphide
50.
NaOH(s)
Molecular Compounds

non-metals only

e- are shared therefore no ions are formed

no charges involved

use prefixes in naming

to go from formula to name: name of first element
(including prefix if necessary), then name for
second element with “-ide” ending (including prefix)
- ___ first element ___ second element-ide
Example:
N2O =
CO2 =
P4O10 =

to go from name to formula: write the symbol for each elelemt, then use the prefixes to determine th
subscripts
Example:
carbon monoxide =
carbon tetrachloride =

remember the memorizers??????
NH3 = ammonia
H2O = water
Chemistry 20
Unit 0: Chemistry Review
H2S = hydrogen sulphide
HF, HCl, HBr, HI = mo prefixes
CH4 = methane
CH3OH = methanol
C2H6 = ethane
C2H5OH = ethanol
C6H12O6 = glucose
C12H22O11 = sucrose
O3 = ozone
H2O2 = hydrogen perioxide
Acids

always have aqueous (aq) as the state and always have hydrogen
Rules
hydrogen _________ide becomes hydro_______ic acid
hydrogen _________ate becomes __________ic acid
hydrogen _________ite
becomes __________ous acid
Examples:
Change each of the following to the appropriate acid name and give the formula:
hydrogen iodide =
hydrogen phosphate =
hydrogen nitrite =
hydrogen sulphite =
Chemistry 20
Unit 0: Chemistry Review
Lesson #R5: Review of Molecular Nomenclature and Acids
Complete the following chart:
Formula
1.
NO3(g)
ammonia
2.
3.
IUPAC Name
H2S(g)
4.
oxygen difluoride
5.
methane
6.
CH3OH(l)
7.
HBr(aq)
sulphurous acid
8.
9.
CS2(l)
hydrosulphuric acid
10.
11.
SO2(g)
dinitrogen tetraoxide
12.
13.
HNO2(aq)
carbon monoxide
14.
15.
C12H22O11(s)
16.
hypochlorous acid
17.
diarsenic trioxide
18.
ethanol
19.
H2CO3(aq)
perchloric acid
20.
21.
P4O10(s)
sulphur trioxide
22.
23.
CF4(l)
silicon dioxide
24.
25.
CH3COOH(aq)
Chemistry 20
Unit 0: Chemistry Review
Lesson #R6: Review of Nomenclature
Complete the following chart:
Class
1.
Formula
IUPAC Name
H3PO4(aq)
2.
chlorous acid
3.
magnesium
4.
Al2(SO4)3(s)
magnesium chloride
5.
6.
NH4NO2(s)
phosphorus trihydride
7.
8.
KNO3(s)
sodium nitrate hexahydrate
9.
10.
HNO2(aq)
11.
Al(OH)3(s)
sodium sulphate octahydrate
12.
13.
(NH4)2SO4(s)
14.
PbF4(s)
hydrogen peroxide
15.
16.
PbO(s)
hydrofluoric acid
17.
18.
KClO(s)
bromine
19.
20.
N2O3(g)
21.
K2CO3 2 H2O(s)
nitric acid
22.
23.
HF(g)
sodium hydroxide
24.
25.
NaHSO3(s)
Chemistry 20
Unit 0: Chemistry Review
Class
Formula
magnesium sulphate octahydrate
26.
27.
Ca(OH)2(s)
gold (I) chloride
28.
29.
IUPAC Name
CaO(s)
30.
copper (II) sulphate pentahydrate
31.
sulphur
32.
Ca(HCO3)2(s)
33.
KBr(s)
titanium (IV) oxide
34.
35.
PCl5(g)
sodium chlorate
36.
37.
N2H4(l)
38.
hydrogen chloride
39.
chloric acid
40.
lithium thiosulphate
41.
B2H6(g)
42.
nitrogen trichloride
43.
sodium hydrogen sulphite
44.
Al(s)
45.
HBr(aq)
46.
silicon
47.
ammonium phosphate
48.
xenon
49.
SF2(s)
50.
Na2SiO3(s)
Chemistry 20
Unit 0: Chemistry Review
States

acids – always (aq)

elements – can be (s), (l) or (g) … see periodic table

molecular compounds – can be (s), (l) or (g)

ionic compounds – if not in a solution always (s) or if in a solution either (s) or (aq)… look up on
solubility chart
High
solubility
(aq)
all
NO3–
ClO3–
ClO4–
all
Low
solubility
(s)
none
none
NH4+
Ions
CH3COO–
Ag+
Hg+
most
Cl–
Br–
I–
most
SO42–
S2–
OH–
most
group 1
group 2
NH4+
Ag+
Pb2+
Cu+
Hg+
Tl+
Ag+
Pb2+
Ca2+
Ba2+
Sr2+
Ra2+
most
group 1
NH4+
Sr2+
Ba2+
Tl+
most
Examples:
1. NaCH3COO(
2. BaSO4(
3. KOH(
6. CaCO3(
)
7. FeSO4(
)
)
4. Pb(NO3)4(
)
5. Hg(CH3COO)2(
)
endothermic vs. exothermic

reaction types:
)
9. Pb(SO4)2(
)
10. Ca3(PO4)2(
1. hydrocarbon combustion
C?H? + O2(g)
CO2(g) + H2O(g)
Example:
CH4(g) + 2 O2(g)
)
8. (NH4)2S(
Chemical Reactions

)
CO2(g) + 2 H2O(g)
)
CO32–
PO43–
SO32–
group 1
NH4+
most
Chemistry 20
Unit 0: Chemistry Review
2. formation (simple composition)
element + element
compound
Example:
2 Mg(s) + O2(g)
2 MgO(s)
3. simple decomposition
compound
element + element
Example:
2 H2O(l)
2 H2(g) + O2(g)
4. single replacement
element + compound
element + compound
Example:
Cu(s) + 2 AgNO3(aq)
2 Ag(s) + Cu(NO3)2(aq)
5. double replacement
compound + compound
compound + compound
Example:
Pb(NO3)2(aq) + 2 KI(aq)
Balancing Reactions

law of conservation of matter says that matter
cannot be created or destroyed, it can only
change forms

we must balance chemical equations to
conserve matter
2 KNO3(aq) + PbI2(aq)
Chemistry 20
Unit 0: Chemistry Review
Examples:
___ CH4(g) + ___ O2(g)
___ CO2(g) + ___ H2O(g)
___ C2H4(g) + ___ O2(g)
___ CO2(g) + ___ H2O(g)
Lesson #R7: Review of Chemical Reactions
A. For each of the following reactions, identify the reaction type and balance the reaction.
________________1. _____Al(s) + _____O2(g) 
_____ Al2O3(s)
________________2. _____HCl(aq) + _____Ca(OH)2(s) 
________________3. _____CH4(g)
+ _____ O2(g) 
_____CaCl2(aq)
+ _____HOH(l)
_____CO2(g) + _____H2O(g)
________________4. _____Zn(s) + _____Pb(CH3COO)2(aq)  _____Pb(s) + _____Zn(CH3COO)2(aq)
________________5. _____SO3(g) + _____H2O(g) 
________________6. _____HgO(l) 
_____Hg(l)
________________7. _____CaCO3(s) 
_____H2SO4(aq)
+ _____O2(g)
_____CaO(s) + _____CO2(g)
________________8. _____NaI(aq) + _____Pb(NO3)2(aq) 
________________9. _____Cl2(g) + _____NaI(aq)

_____PbI2(s) + _____NaNO3(aq)
_____I2(s)
________________10. _____Al2(SO4)3(aq) + _____Ca(OH)2(aq) 
+ _____NaCl(aq)
_____Al(OH)3(s) + _____CaSO4(s)
________________11. ___Al2(SO4)3(aq) + ___Ca(HCO3)2(aq)  ___Al(OH)3(s) + ___CaSO4(s) + ___CO2(g)
________________12. _____C8H18(l)
+ _____ O2(g) 
_____CO2(g) + _____H2O(g)
Chemistry 20
Unit 0: Chemistry Review
________________13. _____H2O(l) 
________________14. _____Ba(s)
_____H2(g) + _____O2(g)
+ _____HOH(l) 
_____H2(g)
________________15. _____H2SO4(aq) + _____Ca3(PO4)2(s) 
+ _____Ba(OH)2(aq)
_____H3PO4(aq) + _____CaSO4(s)
B. For each of the following word equations, write out the balanced chemical reaction including all
states and identify the reaction type.
________________1. water  hydrogen + oxygen
________________2. nitrogen + hydrogen  ammonia gas
________________3. sulphuric acid + sodium hydroxide  water + sodium sulphate
________________4. aluminum + copper (II) nitrate  copper + aluminum nitrate
________________5. chlorine + potassium bromide  bromine + potassium chloride
________________6. sodium hydroxide + aluminum sulphate  aluminum hydroxide + sodium sulphate
________________7. phosphorus + oxygen  solid tetraphosphorus decaoxide
________________8. lead (II) nitrate + sodium iodide  lead (II) iodide + sodium nitrate
________________9. methanol + oxygen  carbon dioxide + water vapour
Chemistry 20
Unit 0: Chemistry Review
________________10. nitrogen dioxide gas + water  nitric acid + nitrogen monoxide gas
Predicting Reactions

SR and DR reactions always happen in solutions so for ionic compounds check solubility table

composition and decomposition do NOT happen in solutions so ionic compounds are (s)
Example:
1. potassium iodide solution is added to lead(II) nitrate solution
2. copper metal is added to a solution of silver nitrate
3. chlorine gas is bubbled through a solution of phosphide
Lesson #R8: Review of Predicting Chemical Reactions
For each of the following reactions:
1.
Write the correct equation including states for each element and compound.
2.
Balance the equation.
3.
State the reaction type.
1. nitrogen triiodide decomposes explosively into its elements.
2. gallium metal reacts with hydrochloric acid.
3. In a charcoal barbeque, some of the carbon undergoes incomplete combustion to produce deadly
carbon monoxide gas.
Chemistry 20
Unit 0: Chemistry Review
4. Solutions of calcium nitrate and potassium phosphate are mixed.
5. chlorine gas is bubbled through an aluminum iodide solution.
6. iron reacts with silver nitrate. The iron (III) compound is formed.
7. acetylene (C2H2(g)) burns in a welding torch.
8. copper ore (copper (II) oxide) is decomposed to produce copper metal.
9. titanium (IV) chloride solution reacts with a sodium phosphate solution.
10. sulphuric acid is neutralized by sodium hydroxide.
Significant Digits

any digit from 1-9 is significant

trailing zeros are significant
Examples:
6.3800
12 000

“sandwich” zeros are significant
Examples:
2.04
1005.002
Chemistry 20
Unit 0: Chemistry Review

leading zeros are not significant
Example:
0.0065

counted objects and constants are not included in sig digs
/
: multiply or divide then round answer to the lowest number of significant digits
+/
: add or subtract then round answer to the lowest number of decimal places
Lesson #R9:Review of Significant Digits, Scientific Notation and SI Units
State the number of significant digits in each of the following measured values:
1. 18.56 g
_________________
4. 1.00 W
2. 1500C
_________________
5. 0.05730 mol _________________
3. 0.0062 L
_________________
6. 8.0 × 10-2 mL _________________
7. 14.08 cm
_________________
9. 0.100 km
_________________
10. 62 km/h
_________________
8. 1.58 × 108 m _________________
_________________
Convert the following numbers into scientific notation. The number in brackets indicates the number
of significant digits the answer is to be rounded to.
1. 1000
_________________ (1)
4. 0.00001098
_________________ (3)
2. 492.32
_________________ (3)
5. 6 995 000
_________________ (3)
3. 0.0573
_________________ (2)
6. 62.49
_________________ (2)
Using the SI Prefixes table on your data sheet, perform the following conversions. Maintain the same
number of significant digits in each conversion.
1. 0.520 km = _________________m
6. 200 ML = _________________L
2. 100 mL = _________________ L
7. 45 g = _________________kg
3. 152.5 cm = _________________m
8. 10.8 mol = _________________ mmol
4. 3300 mg = _________________g
9. 0.450 L = _________________ mL
5. 650 kg = _________________g
10. 1500 m =_________________ km
Perform the following calculations. Round your answer to the correct number of significant digits,
using scientific notation where necessary. Include units.
1. 16.56 mL – 6.3 mL = _________________
2. 21.4 g ÷ 0.825 mol = _________________
Chemistry 20
Unit 0: Chemistry Review
3. 480 km + 24.07 km = _________________
4. 0.550 mol × 40.00 g/mol = _________________
5. 18.4 g/mL × 5.5 mL = _________________
6. 22.99 g/mol + 35.45 g/mol = _________________
7. 18.5C  4.5C = _________________
8. 6.0 g ÷ 24.30 g/mol = _________________
9. 19.55 mL  17.55 mL = _________________
10. 15 600 g ÷ 2000 mol = _________________
The Mole
it
is a number = 6.02x1023 “molecules”
1. Molar Mass

sum of the individual atomic masses for each element in a compound
Examples:
1. CO2 =
2. Al(OH)3 =
3. Cu(ClO3)2 =
Practice:
1.
NaOH
2.
Ca3(PO4)2
3.
ammonium hydroxide
Chemistry 20
Unit 0: Chemistry Review
2. Mole/Mass Calculations
n =
m
M
m = nM
where: n = number of moles in mol
m = mass in g
M = molar mass in
g
mol
Examples:
1. How many moles are in 8.06 g of MgO?
2. What is the mass of 0.677 mol of potassium sulphide?
Chemistry 20
Unit 0: Chemistry Review
Lesson #R10: Review of Molar Mass and Mole Calculations
Complete the following chart, showing all calculations, formulas, substitutions, units and significant
digits.
Name and
Formula
Molar Mass
Mass
Moles
1. NaCl(s)
0.20 mol
2. sodium
hydroxide
5.48 g
3. (NH4)3PO4(s)
0.600 mol
4. sodium
carbonate
octahydrate
50 g
Chemistry 20
Unit 0: Chemistry Review
Name and
Formula
Molar Mass
Mass
Moles
5. Ca(NO3)2(s)
8.45 g
6. potassium
dichromate
5.65 g
7. Na2CO3(s)
0.850 mol
8. sulphur
trioxide
1.45 mol
Chemistry 20
Unit 0: Chemistry Review
Science 10 Review
1.Define the following terms:
a) proton
b) neutron
c) electron
d) atom
e) ion
f) valence electron
g) octet rule
h) monovalent
i) multivalent
j) endothermic
k) exothermic
l) law of conservation of matter
m) mole
n) molar mass
o) mass number
p) isotope
2. Draw the energy level diagrams for the following:
a) fluorine atom
d) nitride ion
b) carbon atom
e) argon atom
c) lithium ion
f) magnesium ion
3. Where on the periodic table would you find nonmetals? What kind of charge do all nonmetals have?
4. Where on the periodic table would you find metals? What kind of charge do all metals have?
5. Perform the following unit conversions:
a) 500 kg = _________________ g
b) 25.5 mL = _________________ L
c) 102.6 mmol = __________________mol
d) 58.2 MJ = _____________________ J
e)
f)
g)
h)
600 mg = ____________________ g
9.85 GL = ____________________ L
6.85 cm = ____________________ m
680 nm = _____________________ m
6. Calculate the number of moles in 6.55 g of NaHCO3(s).
7. Calculate the mass of 8.98 mol of AgNO3(s).
8. What is the mass of 0.155 mol of potassium phosphate?
Chemistry 20
Unit 0: Chemistry Review
9. How many moles are in 0.558 kg of dinitrogen dioxide?
10. Complete the following chart:
Class
Formula
1.
SrCl2
2.
H2S(aq)
3.
Na2O
4.
H2O
5.
CaS2O3  7H2O
6.
Fe(IO3)3
7.
P2O4
8.
S8(s)
9.
Ni(OH)2
10.
H3PO4(aq)
11.
NaCl
12.
N2(g)
13.
Sb2(SO3)5
14.
Ca(s)
IUPAC Name
15.
sodium chloride
16.
copper (II) sulphate pentahydrate
17.
ammonium sulphide
18.
bismuth (III) sulphate
19.
sodium sulphate decahydrate
20.
water
21.
copper (I) oxide
22.
calcium
23.
hydroiodic acid
24.
radon gas
25.
ethanol
26.
sucrose
Chemistry 20
Unit 0: Chemistry Review
27.
nitrogen gas
28.
carbonic acid
29.
dinitrogen monoxide
30.
nitrous acid
11. After the chemical formula for each compound, state the solubility with either (aq) for soluble or (s)
for low solubility in water.
1. K2S (
)
2. NH4CH3COO (
)
3. Fe(OH)3 (
)
4. HgBr (
)
5. BaSO4 (
)
6. CaCl2 (
)
7. CuI2 (
)
8. Ca(CH3COO)2 (
)
9. FeSO4 (
)
10.
Co(NO3)2 (
)
11.
12.
13.
14.
15.
16.
17.
18.
19.
20.
Zn3(PO4)2 (
)
PbI2 (
)
ZnSO4 (
)
Cu(NO3)2 (
)
AgCl (
)
CdSO4 (
)
NH4Cl (
)
CuS (
)
PbCl2 (
)
Na3PO4 (
)
12. Balance the following reactions and give the reaction type:
a)
_____Al2S3(s) 
_____Al(s) + _____S8(s)
b) _____N2(g) + _____O2(g) 
_____ NO2(g)
c) _____Na(s) + _____Pb(CH3COO)2(aq) 
d) _____Ba(s)
e) _____CH4(g)
+ _____HOH(l) 
+ _____ O2(g) 
_____H2(g)
+ _____Ba(OH)2(aq)
_____CO2(g) + _____H2O(g)
f) _____CaSO4(s) + _____AgNO3(aq) 
g) _____CH3OH(l)
_____Pb(s) + _____NaCH3COO(aq)
+ _____ O2(g) 
_____Ag2SO4(s) + _____Ca(NO3)2(aq)
_____CO2(g) + _____H2O(g)
h) _____Na2SO4(aq) + _____FeCl3(aq) 
_____NaCl(aq) + _____Fe2(SO4)3(aq)
Chemistry 20
Unit 0: Chemistry Review
i) _____Cr2O3(s) 
_____Cr(s) + _____O2(g)
j) _____V(s) + _____S8(g)

_____ V2S5(s)
13. For each of the following word problems, give the reaction type and write out a balanced
chemical reaction including all states of matter.
a) nitrogen triiodide decomposes explosively into its elements.
b) In a charcoal barbeque, some of the carbon undergoes incomplete combustion to produce deadly
carbon monoxide gas.
c) Solutions of calcium nitrate and potassium phosphate are mixed.
d) The main fuel used to propel rockets into outer space is liquid hydrogen combining with liquid
oxygen to produce water vapour.
e) chlorine gas is bubbled through an aluminum iodide solution.