Download CHEMISTRY Periodic Table of the Elements

CHEMISTRY
Periodic Table of the Elements
1
2
3
4
5
6
7
8
9
10
11
12
13
14
15
16
17
18
19
20
21
22
23
24
25
26
27
28
29
30
31
32
33
34
35
36
37
38
39
40
41
42
43
44
45
46
47
48
49
50
51
52
53
54
55
56
57
72
73
74
75
76
77
78
79
80
81
82
83
84
85
86
87
88
89
104
105
106
107
108
109
110
111
112
113
114
115
116
117
118
Elements of the Periodic Table You Need to Know
atomic
#
symbol
element name
atomic symbol
#
1
19
2
20
3
24
4
25
5
26
6
27
7
28
8
29
9
30
10
35
11
46
12
47
13
50
14
51
15
53
16
79
17
80
18
82
element name
Chemistry Terminology
Place the following terms next to the appropriate definition.
anion
atom
atomic number
cation
electron
isotope
mass number
metal
metalloid
molecule
neutral
neutron
non-metal
period
proton
group
The negative particle in an atom
Having no charge
Two or more atoms bonded together
A negatively charged ion
The neutral particle in an atom
Tells the number of protons
A positively charged ion
The positive particle in an atom
Is made of protons, electrons and neutrons together
An element located near the “staircase” of the periodic table with some
metallic and some non-metallic properties
An element that is generally a non-conductor of electricity and is brittle
An element that is a good conductor of electricity, malleable and ductile
Atoms of an element that have the same number of protons and electrons, but
differ in the number of neutrons
A column of elements in the periodic table
A row of elements in the periodic table
Bohr diagrams
Complete the following Bohr diagrams.
1.
28
Al
35
2.
Cl
p=
n=
4. 55Fe
5. 48Ti
79
51
V+2
32 -2
S
11.
9.
29
Si
31 +3
P
p=
n=
12.
80
p=
n=
14.
p=
n=
p=
n=
p=
n=
p=
n=
13.
p=
n=
p=
n=
p=
n=
10.
Mn
6. 84Kr
8. 52Cr
Se
55
p=
n=
p=
n=
7.
3.
56
Fe+3
p=
n=
Br-1
p=
n=
15.
57
Ni+2
p=
n=
Electron Configurations
Fill in the following chart using  and  symbols for electrons for the given elements.
1s
1. Li
2. N
3. Ne
4. Zn
5. Ti
6. Ca
7. Mn
8. V
9. Si
10. Mg
11. Ar
12. C
13. Ni
14. Sc
15. O
16. K
17. F
2s
2p
3s
3p
4s
3d
Electron Configurations
atomic
#
1
2
3
4
5
6
7
8
9
10
name
symbol
total
e-
valence
e-
closest
noble
gas
full electron configuration
abbreviated electron configuration
Lewis dot diagram
atomic
#
15
17
19
20
22
25
26
29
30
35
36
name
symbol
total
e-
valence
e-
closest
noble
gas
full electron configuration
abbreviated electron configuration
Lewis dot diagram
Properties of Elements Lab
Use careful observations and research to complete the following table. Make sure you record
your sources on the last page.
Atomic
#
1
2
3
6
7
8
10
11
12
13
14
Symbol
Element
Name
Mass of 1
a.m.u.
(2 decimals)
Properties
(colour, clarity, state
and 1 other property)
Two Common Uses
15
16
17
18
19
20
24
25
26
27
28
29
30
32
33
34
35
36
47
48
50
51
52
53
78
79
80
82
83
References:
Classification of Elements
Metals, Non-metals, Metalloids
1
2
H
He
3
4
5
6
Li
Be
B
C
11
12
13
14
Na
Mg
Al
Si
19
20
21
22
23
24
25
26
27
28
29
30
31
K
Ca
Sc
Ti
V
Cr
Mn
Fe
Co
Ni
Cu
Zn
37
38
39
40
41
42
43
44
45
46
47
48
Rb
Sr
Y
Zr
Nb
Mo
Tc
Ru
Rh
Pd
Ag
55
56
57
72
73
74
75
76
77
78
Cs
Ba
La
Hf
Ta
W
Re
Os
Ir
87
88
89
104
105
106
107
108
109
Fr
Ra
Ac
58
59
60
61
62
63
Ce
Pr
Nd
Pm
Sm
Eu
90
91
92
93
94
Th
Pa
U
Np
Pu
7
8
9
10
O
F
Ne
15
16
17
18
P
S
Cl
Ar
32
33
34
35
36
Ga
Ge
As
Se
Br
Kr
49
50
51
52
53
54
Cd
In
Sn
Sb
Te
I
Xe
79
80
81
82
83
84
85
86
Pt
Au
Hg
Tl
Pb
Bi
Po
At
Rn
110
111
112
113
114
115
116
117
118
64
65
66
67
68
69
70
71
Gd
Tb
Dy
Ho
Er
Tm
Yb
Lu
95
96
97
98
99
100
101
101
103
Am
Cm
Bk
Cf
Es
Fm
Md
No
Lr
N
metals
non-metals
semi-metals
transition metals
Chemical Groups (Families)
1
2
H
He
3
4
5
6
7
8
9
10
Li
Be
B
C
11
12
13
14
15
O
F
Ne
16
17
Na
Mg
Al
Si
18
P
S
Cl
Ar
19
20
21
22
23
24
25
26
27
28
29
30
31
K
Ca
Sc
Ti
V
Cr
Mn
Fe
Co
Ni
Cu
Zn
Ga
32
33
34
35
36
Ge
As
Se
Br
37
38
39
40
41
42
43
44
45
46
47
48
Kr
49
50
51
52
53
Rb
Sr
Y
Zr
Nb
Mo
Tc
Ru
Rh
Pd
Ag
54
Cd
In
Sn
Sb
Te
I
55
56
57
72
73
74
75
76
77
78
Xe
79
80
81
82
83
84
85
Cs
Ba
La
Hf
Ta
W
Re
Os
Ir
86
Pt
Au
Hg
Tl
Pb
Bi
Po
At
Rn
87
88
89
104
105
106
107
108
109
Fr
Ra
Ac
110
111
112
113
114
115
116
117
118
N
alkali metals
alkaline earth
metals
halogens
noble gases
58
59
60
61
62
63
64
65
66
67
68
69
70
71
Ce
Pr
Nd
Pm
Sm
Eu
Gd
Tb
Dy
Ho
Er
Tm
Yb
Lu
90
91
92
93
94
95
96
97
98
99
100
101
101
103
Th
Pa
U
Np
Pu
Am
Cm
Bk
Cf
Es
Fm
Md
No
Lr
Common Charges
1
2
H
He
3
4
5
6
7
8
9
10
Li
Be
B
C
11
12
13
14
15
O
F
Ne
16
17
18
Na
Mg
Al
Si
19
20
21
22
23
24
25
26
27
28
29
30
31
32
P
S
Cl
Ar
33
34
35
K
Ca
Sc
Ti
V
Cr
Mn
Fe
Co
Ni
Cu
Zn
Ga
36
Ge
As
Se
Br
Kr
37
38
39
40
41
42
43
44
45
46
47
48
Rb
Sr
Y
Zr
Nb
Mo
Tc
Ru
Rh
Pd
Ag
Cd
49
50
51
52
53
54
In
Sn
Sb
Te
I
55
56
57
72
73
74
75
76
77
78
79
Xe
80
81
82
83
84
85
86
Cs
Ba
La
Hf
Ta
W
Re
Os
Ir
Pt
87
88
89
104
105
106
107
108
109
110
Au
Hg
Tl
Pb
Bi
Po
At
Rn
111
112
113
114
115
116
117
118
Fr
Ra
Ac
58
59
60
61
62
63
64
65
66
67
68
69
70
71
Ce
Pr
Nd
Pm
Sm
Eu
90
91
92
93
94
95
Gd
Tb
Dy
Ho
Er
Tm
Yb
Lu
96
97
98
99
100
101
101
103
Th
Pa
U
Np
Pu
Am
Cm
Bk
Cf
Es
Fm
Md
No
Lr
N
+1
+2
+3
-1
-2
-3
-4
Ionic Charges and Chemical Families
Alkali Metals
element name
symbol
valence
electrons
electron dot
diagram of
atom
electron dot diagram
of stable ion
ionic charge
Alkaline Earth Metals
element name
symbol
valence
electrons
electron dot
diagram of
atom
electron dot diagram
of stable ion
ionic charge
Halogens
element name
symbol
valence
electrons
electron dot
diagram of
atom
electron dot diagram
of stable ion
ionic charge
electron dot diagram
of stable ion
ionic charge
Noble Gases
element name
symbol
valence
electrons
electron dot
diagram of
atom
Metallic Elements
element name
symbol
valence
electrons
electron dot
diagram of
atom
electron dot diagram
of stable ion
ionic charge
aluminum
tin
lead
Non-Metallic Elements
element name
symbol
valence
electrons
electron dot
diagram of
atom
electron dot diagram
of stable ion
ionic charge
nitrogen
oxygen
phosphorous
Conclusions:
1. The alkali metals always develop a charge of
.
2. The alkaline earth metals always develop a charge of
.
3. The halogens always develop a charge of
.
4. The noble gases always develop a charge of
.
5. Metals develop positive charges and become
.
6. Non-metals will develop a negative charge and become
.
Australian – atoms and their electrons
1.
How are atoms grouped on the periodic table?
2.
All matter is made up of
3.
All atoms are made up of
4.
Protons have a
5.
Most of the atoms is
6.
Atoma means
7.
Dalton said that molecules are made of
8.
JJ Thompson discovered the
.
9.
Thompson’s model was called
.
.
,
and
.
charge
.
.
.
10. Rutherford shot the nuclei of helium atoms at a thin sheet of
11. Rutherford thought the positive
spinning around it.
.
was in the centre of the atom with electrons
12. Bohr said the electrons circled in
.
13. Atoms are positioned in the periodic table according to the arrangement of
14. The smallest atom is
.
.
15. Deuterium is called
.
16.
has a low density.
17. Beryllium is in
on the periodic table.
18. Fluorine has
outer electrons.
19.
is very corrosive and the only element to attack glass.
20.
is satisfied because it has a full outer shell of electrons.
21. The elements on the right side of the periodic table are called
22. Write the symbols for 3 noble gases.
23. Chlorine has
layers of electrons.
24. Which halogen is the most attractive?
25.
is below carbon on the periodic table.
26. Who are the bouncers at the under 18 disco?
27. Who is an undesirable element?
28. Who’s always trying to get into under 18 parties?
.
29. Which elements are police?
30. How many electrons does oxygen want?
31. Sharing electrons is called a
.
32. The hardest substance in the world is
.
33. CH4 is called
.
34. NH3 is called
.
35. H2O is called
.
36. CO2 is called
.
37. Who is the element of sodium’s dream?
38. Sodium and chlorine form an
bond.
39. How can sodium and chlorine split?
Atom bond: the atom with the golden electrons
1.
Who is the atom with the blonde bomb shell?
2.
What is the only atom that can attack silica – SiO2 (glass)?
3.
How many bonds can halogens form?
4.
How many bonds can carbon form?
5.
Atoms all want
6.
The smaller an atom is, the
it holds its electrons.
7.
Lop-sided charges are called
?
8.
The strongest of the dipole-dipole bonds is the
covalent bond.
9.
KCl is called
10. Water is a
electrons?
, but it is only 1/10th the strength of a
.
solvent.
11. Which element is the police officer?
12. What is O3 called?
13. CCl4 is called
.
14. What is project D?
15. This molecule may be unzipped by breaking
.
Elements and the Periodic Table – Simple Ions
Complete the following table. Note that the name of the NON-METALLIC ion ends in “ide”
while the name for METALLIC ions uses the normal name of the element.
ion name
symbol
number of
protons
number of
electrons
number of
electrons
transferred
fluoride
F-1
9
10
0 1 2 3 4
lost
gained
53
54
0 1 2 3 4
lost
gained
0 1 2 3 4
lost
gained
0 1 2 3 4
lost
gained
0 1 2 3 4
lost
gained
0 1 2 3 4
lost
gained
Sr+2
0 1 2 3 4
lost
gained
H+1
0 1 2 3 4
lost
gained
8
0 1 2 3 4
lost
gained
12
0 1 2 3 4
lost
gained
10
0 1 2 3 4
lost
gained
36
0 1 2 3 4
lost
gained
0 1 2 3 4
lost
gained
0 1 2 3 4
lost
gained
0 1 2 3 4
lost
gained
0 1 2 3 4
lost
gained
0 1 2 3 4
lost
gained
0 1 2 3 4
lost
gained
0 1 2 3 4
lost
gained
0 1 2 3 4
lost
gained
0 1 2 3 4
lost
gained
16
potassium
Ca+2
35
aluminum
34
H-1
lithium
Rb+1
17
beryllium
7
C-4
P-3
18
are electrons
lost or gained?
same number of
e- as this noble
gas
neon
neon
Flame Colours Lab
name:
Introduction:
When atoms are “excited”, electrons can jump up one or more energy levels. Energy
is needed for this to happen. When these electrons fall back down towards their
original energy level, they release energy, some of which we see as visible light. The
colours of light emitted when the electrons fall back down can be used to identify the
atom.
Purpose: To use the flame test to distinguish between different cation salts.
Materials:
evaporating dishes
cation salts (in methanol)
lighter
safety glasses
Procedure:
1. Make qualitative observations of the original salts and solutions.
2. Light a control of pure methanol.
3. Light each of the solutions and observe the flame colour.
4. Allow the flames to burn out. The dishes will be very hot at this point.
5. If required, allow the dishes to cool, then relight the solutions using more methanol.
6. Test the unknowns and decide which cations they contain.
Observations:
cation
none
sodium
barium
calcium
strontium
cobalt
lithium
potassium
copper
symbol
qualitative observations
of salt and solutions
qualitative observations
of flame
qualitative observations
of salt and solutions
qualitative observations
of flame
cation
unknown #1
unknown #2
unknown #3
unknown #4
Questions:
1.
What is a cation?
2.
Which cations were easiest to distinguish? Why?
3.
Which cations were difficult to distinguish? Why?
4.
Some of these cations are used in fireworks. Which ones would you choose to
make a fireworks display?
Bonus:
The control sample that was burned contained only alcohol. Why would
methanol create a flame that had more than one colour?
Bill Nye – Chemical Reactions
1.
Everything is made up of ___________________________.
2.
Metal rusting, candles burning and food being digested are examples of:
3.
What gas makes iron rust?
4.
When electrons recombine with other electrons, _________________ is given off.
5.
CH2H4O2 (vinegar)+ NaHCO3 (baking soda) react to give off____________ gas.
6.
Two poisons that we can’t live without are _________________________ and
____________________________.
7.
Another name for NaCl is ________________________.
8.
Pyrotechnics is another name for the manufacture of:
9.
How many elements make up everything we know and love?
10. What prize is awarded to people who have done great works for humanity?
11. Alfred Nobel became very rich because he invented _______________________.
12. Vinegar is an ________________________, while baking soda is a
_______________________________.
13. All of the elements are listed on the ___________________________________.
14. Ca is found in____________________, salt contains________________________, and
K is found in _________________________.
15. Why are sodium and potassium grouped together?
Reactions of Metals with Water
Part A:
Review positive tests for OXYGEN, HYDROGEN, CARBON DIOXIDE, ACIDS
and BASES. Draw a series of labeled diagrams with brief descriptions.
Part B: Calcium and Magnesium
For each metal, add a small piece carefully to a test tube approximately ¼ full of
water along with 3 drops of phenolphthalein indicator.
Record at least 4 qualitative observations for calcium with water.
Record at least 2 qualitative observations for magnesium with water.
Questions:
1.
2.
3.
4.
What gas is produced when calcium is added to water? How do you know?
Is the remaining liquid acidic or basic? How do you know?
Which of the metals is more active in its reaction with water?
Write full electron configurations for calcium and magnesium. How many electrons do
each of them have in their outer shell?
5. Which of these 2 metals has their outer shell electrons closer to the nucleus?
Part C: LITHIUM, SODIUM and POTASSIUM
For each metal, add a small piece carefully to a test tube approximately ¼ full of
water along with 3 drops of phenolphthalein indicator.
Make qualitative observations for each of the metals. Be sure to include:
a) the shape of the metal
b) movement of the metal
c) relationship between the metal and the surface of the water
d) any materials produced in the reaction
Questions:
1.
2.
3.
4.
What does the phenolphthalein indicate? What does this mean?
Based on part B, what other substance do you think was produced? Why?
Rate these 3 metals in order of their reactivity with the water.
Write full electron configurations for Li, Na and K. How many electrons do each of them
have in their outer shell?
5. Which of these 3 metals has their electrons closest to the nucleus? How do you think this
is related to your answer in question 3?
6. Make some predictions about the behaviour of:
a) cesium (ie more or less reactive than Li, Na and K. Why?)
b) barium (ie more or less reactive than Ca and Mg. Why?)
Conclusions
1. What happens to the reactivity of the alkali metals as you move down the periodic table?
What does this have to do with outer shell electrons?
2. What happens to the reactivity of the alkaline earth metals as you move down the
periodic table? What does this have to do with outer shell electrons?
Periodic Trends
1. Select the element which has the:
a) highest electronegativity
b) smallest atomic radius
c) smallest electronegativity of the alkali metals
d) largest first ionization energy of period 3
e) smallest first ionization energy of the noble gases
f) largest atomic radius of period 5
g) greatest electronegativity of the halogens
2. Which of the following has the lowest ionization energy?
chromium
iron
calcium
arsenic
copper
3. Which of the following has the largest electronegativity?
neon
xenon
silver
antimony
tin
phosphorous
aluminum
4. Which of the following is the largest?
lithium
neon
potassium
5. Which of the following has the smallest electronegativity?
lead
platinum
bromine
germanium
nickel
6. Which of the following has the smallest ionic radius?
K+
Ca2+
Ar
Cl-
S2-
7. Which of the following has the largest first ionization energy?
Ba
Hg
Au
P
Si
Chemical Bonding Between Atoms
Molecular
Substance
Molecular
formula
trichloromethane
CHCl3
ammonia
NH3
water
H2O
bromine
Br2
hydrogen
H2
chlorine
Cl2
hydrogen
chloride
HCl
Lewis dot
diagram of the
atoms
Lewis dot
diagram of the
molecule
Structural
diagram
Molecular
Substance
Molecular
formula
methane
CH4
arsenious chloride
AsCl3
phosphorous
bromide
PBr2
ethane
C2H6
dinitrogen
tetrahydride
N2H4
hydrogen peroxide
H2O2
silicon tetrahydride
SiH4
Lewis dot
diagram of the
atoms
Lewis dot
diagram of the
molecule
Structural
diagram
Molecular
Substance
Molecular
formula
methanal
H2CO
ethyne
C2H2
nitrogen
N2
carbon dioxide
CO2
hydrogen cyanide
HCN
bromoiodomethane
CH2BrI
tetrafluoroethene
C2F4
Lewis dot
diagram of the
atoms
Lewis dot
diagram of the
molecule
Structural
diagram
Legal Grade 10 Nomenclature Cheat Sheet
Compound Ions
nitrate
NO3-1
carbonate
CO3-2
fluorate
FO3-1
sulphate
SO4-2
chlorate
ClO3-1
phosphate
PO4-3
bromate
BrO3-1
hydrogen carbonate
HCO3-1
iodate
IO3-1
hydrogen sulphate
HSO4-1
hydroxide
OH-1
monohydrogen phosphate
HPO4-2
ammonium
NH4+1
dihydrogen phosphate
H2PO4-1
Latin “ous / ic” Cations
element
higher charge
lower charge
iron
+3
ferric
+2
ferrous
copper
+2
cupric
+1
cuprous
tin
+4
stannic
+2
stannous
antimony
+5
stibbic
stibbous
+3
antimonic
antimonous
gold
+3
auric
+1
aurous
mercury
+2
mercuric
+1
mercurous
lead
+4
plumbic
+2
plumbous
phosphorous
+5
phosphoric
+3
phosphorous
Prefixes
mono
di
tri
tetra penta
hexa hepta
octa
nona deca
Naming Binary Compounds
Rules:
1. Determine which element is the cation. The cations are usually found toward the left of
the periodic table and are written first in the name and formula of a molecule.
2. Determine which element is the anion. The anions are non-metals (or hydrogen), located
on the right side of the periodic table, and are written second in the name and formula of
a molecule.
3. Write the cation first using the name of the element.
4. Write the anion second, dropping the usual ending and replacing it with “ide”.
element
fluorine
chlorine
bromine
iodine
hydrogen
anion
fluoride
chloride
bromide
iodide
hydride
element
oxygen
sulphur
nitrogen
phosphorous
carbon
eg. KCl = potassium chloride
Write the correct chemical name for each of the following:
1. MgO
11. K2S
2. LiF
12. Na2O
3. NaBr
13. K2O
4. CaO
14. Ca3P2
5. AlN
15. H2O
6. NaI
16. CaBr2
7. Al2S3
17. MgS
8. Ag3P
18. ZnBr2
9. BaCl2
19. B2O3
10. AlCl3
20. Ba2C
anion
oxide
sulphide
nitride
phosphide
carbide
Writing Chemical Formulae
Rules:
1. Write the chemical symbol for the cation first, followed by the symbol of the anion.
2. Write the charge of each ion above each symbol.
3. Cross the charges, ignoring the signs.
4. Reduce the numbers if there is a common factor.
5. If the number beside an element is 1, do not write it.
(The total positive charge will now equal the total negative charge in the molecule.)
Example: silicon oxide
Rule 1
Si
O
Rule 2
Si+4
O-2
Rule 3
Si2O4
Rule 4
SiO2
Write the correct chemical formula for each of the following:
1. sodium nitride
11. calcium phosphide
2. sodium oxide
12. sodium fluoride
3. calcium chloride
13. boron nitride
4. magnesium sulphide
14. calcium hydride
5. silicon oxide
15. hydrogen oxide
6. aluminum carbide
16. aluminum nitride
7. boron fluoride
17. potassium carbide
8. potassium nitride
18. zinc iodide
9. cesium oxide
19. barium bromide
10. aluminum bromide
20. silver selenide
Multiple Valences
Latin method – “ous/ic”
Many cations have more than one possible charge. The latin method is the oldest method
used to deal with this program, and while it can’t be used for many molecules, it is still used
in industry.
Rules:
1. Determine the charge on the cation.
2. Select the proper name for the cation.
a) The “ous” ending refers to the lower cation charge.
b) The “ic” ending refers to the higher cation charge.
3. Write the name of the anion as before, using the “ide” ending
element
iron
copper
tin
antimony
higher charge
+3
ferric
+2
cupric
+4
stannic
stibbic
+5
antimonic
lower charge
ferrous
cuprous
stannous
stibbous
+3
antimonous
+2
+1
+2
element
gold
mercury
lead
phosphorous
higher charge
+3
auric
+2
mercuric
+4
plumbic
+5
phosphoric
Write the correct “ous/ic” name for each of the following:
1. FeCl2
6. Au2S3
2. Cu2O
7. Sb2O5
3. Hg3N
8. SnBr5
4. PbO2
9. AuCl
5. CuF2
10. SbF5
Write the correct formula for each of the following:
1. ferrous chloride
6. stannous phosphide
2. plumbic oxide
7. aurous fluoride
3. ferric nitride
8. mercuric nitride
4. cuprous sulphide
9. stibbic bromide
5. stibbous oxide
10. stannic carbide
+1
+1
+2
+3
lower charge
aurous
mercurous
plumbous
phosphorous
Multiple Valences
Prefix Method
This method is commonly used only for naming binary compounds composed of two non-metals.
Rules:
1. A prefix is used to indicate the number of atoms in the molecule.
number of
atoms
1
2
3
4
5
prefix
mono
di
tri
tetra
penta
number of
atoms
6
7
8
9
10
prefix
hexa
hepta
octa
nona
deca
2. Place the appropriate prefix in front of the cation (mono is dropped in the first element).
3. Place the appropriate prefix in front of the anion, using the “ide” suffix as before.
Exceptions:
1. Peroxides – contain O2-2 ion
Peroxides have an extra oxygen atom. Write the formula for the ordinary oxide and add
one additional oxygen atom. Peroxides are NOT reduced.
barium oxide – BaO
barium peroxide – BaO2
hydrogen oxide – H2O
hydrogen peroxide – H2O2
2. Diatomic Elements
The following gaseous elements consist of two atoms joined together. They do not occur
naturally as a single atom. (mnemonic – HOFBrINCl or Hey NO halogens)
H2 N2 O2 F2 Cl2 Br2 I2
1. sulphur dioxide
2. carbon disulphide
3. nitrogen trichloride
4. phosphorous pentabromide
5. diiodine pentasulphide
6. selenium tetrachloride
7. bromine heptafluoride
8. nitrogen monoxide
9. selenium trioxide
10. dinitrogen trisulphide
Multiple Valences
IUPAC (Roman Numeral) Method
The IUPAC (International Union of Pure and Applied Chemists) method is a standardized
nomenclature system that always works. The Roman Numerals are NOT used when there is only one
possible positive valence (ie Columns I, II, III, Ag, Zn and Cd).
Rules:
1. Determine the charge on the anion (there is only one possibility).
2. Determine the total negative charge by multiplying the anion charge by the number of anions
present.
3. The total positive charge equals the total negative charge in a neutral molecule.
4. Divide the total negative charge by the number of cations present to determine the charge on
each cation.
5. Write down the name of the cation.
6. Write the charge on the cation using Roman Numerals in brackets after the cation.
7. Write down the name of the anion using the “ide” ending.
eg. Fe2O3
1.
2.
3.
4.
charge on O = -2
total negative charge = –2 x 3 = -6
total positive charge = +6
charge on iron = +6 2 = +3
5. name of molecule = iron (III) oxide
Write the correct IUPAC name for each of the following.
1. FeCl2
6. Au2S3
2. Cu2O
7. Sb2O5
3. Hg3N
8. SnBr4
4. PbO2
9. AuCl3
5. CuF2
10. CrBr3
Write the correct formula for each of the following.
1. mercury (I) oxide
6. tin (II) phosphide
2. lead (IV) chloride
7. gold (I) fluoride
3. iron (III) nitride
8. mercury (II) nitride
4. copper (I) sulphide
9. antimony (V) bromide
5. antimony (III) oxide
10. tin (IV) carbide
Compound Ions
Many ions consist of more than one element. These ions all have special names which you
will not need to memorize. A chart of the compound ions will be provided to you for all tests
and quizzes.
The charge given in the chart is the charge on the compound ion as a unit.
Compound molecules are named using the IUPAC system, the only difference being that if
more than one of the compound ions is needed to form a neutral molecule, brackets are placed
around the ion.
nitrate
NO3-1
carbonate
CO3-2
fluorate
FO3-1
sulphate
SO4-2
chlorate
ClO3-1
phosphate
PO4-3
bromate
BrO3-1
hydrogen carbonate
HCO3-1
iodate
IO3-1
hydrogen sulphate
HSO4-1
hydroxide
OH-1
monohydrogen phosphate
HPO4-2
ammonium
NH4+1
dihydrogen phosphate
H2PO4-1
eg. iron (III) sulphate =
Fe+3
SO4-2
Fe2(SO4)3
Complete the following table.
1. silver carbonate
11. Fe(NO3)3
2. calcium nitrate
12. AuClO3
3. lead (II) bromate
13. Mn(HCO3)2
4. ammonium chloride
14. Sr(FO3)2
5. manganese (IV) iodate
15. Ti(BrO3)4
6. potassium phosphate
16. Co3(PO4)4
7. lithium hydrogen carbonate
17. (NH4)2SO4
8. copper (II) sulphate
18. Ni(OH)3
9. zinc dihydrogen phosphate
19. Sb(IO3)5
10. aluminum hydroxide
20. Sn(CO3)2
Simple Nomenclature
ions
molecule
molecular name
1. silicon oxide
21. MgCl2
2. boron fluoride
22. SiC
3. aluminum carbide
23. Al2S3
4. potassium nitride
24. SiH4
5. cesium oxide
25. H2S
6. aluminum bromide
26. Ag3P
7. calcium phosphide
27. H2O
8. sodium fluoride
28. MgO
9. boron nitride
29. CaH2
10. nitrogen hydride
30. NaBr
11. hydrogen oxide
31. KF
12. calcium nitride
32. C3N4
13. aluminum nitride
33. H2S
14. calcium oxide
34. B2S3
15. potassium sulphide
35. BaO
16. zinc oxide
36. ZnO
17. silver nitride
37. SrS
18. lithium fluoride
38. BeS
19. magnesium iodide
39. SiCl4
20. hydrogen arsenide
40. AlF3
Binary Nomenclature
1. ferric sulphide
21. MgCl2
2. calcium chloride
22. SiC
3. tin (IV) carbide
23. Al2S3
4. carbon dioxide
24. SiH4
5. aluminum bromide
25. H2S
6. rubidium nitride
26. Ag3P
7. cuprous phosphide
27. H2O
8. stibbic fluoride
28. MgO
9. antimony (V) fluoride
29. CaH2
10. cesium oxide
30. NaBr
11. mercury (II) iodide
31. KF
12. plumbic chloride
32. C3N4
13. gold (I) nitride
33. H2S
14. zinc sulphide
34. B2S3
15. silver bromide
35. BaO
16. stannous oxide
36. ZnO
17. copper (II) phosphide
37. SrS
18. beryllium iodide
38. BeS
19. mercuric carbide
39. SiCl4
20. table salt
40. AlF3
Compound Ion Nomenclature
1. copper (II) nitrate
21. K2CO3
2. ferrous sulphate
22. Na2SO4
3. potassium chlorate
23. Zn3(PO4)2
4. zinc carbonate
24. Hg2SO4
5. silver phosphate
25. Ba(NO3)2
6. sodium sulphate
26. Fe(HSO4)3
7. barium hydroxide
27. Pb3(PO4)4
8. ammonium phosphate
28. Hg(NO3)2
9.
29. FeSO4
plumbous hydrogen carbonate
10. cuprous nitrate
30. Sb(HCO3)5
11. mercury (II) hydrogen sulphate
31. MgSO4
12. zinc sulphate
32. Ag3PO4
13. auric phosphate
33. NH4NO3
14. aluminum nitrate
34. Sn(OH)4
15. ammonium hydroxide
35. BPO4
16. boron carbonate
36. Be(OH)2
17. lead (IV) hydrogen carbonate
37. AuHSO4
18. ammonium sulphide
38. Cu3(PO4)2
19. mercuric phosphate
39. AgHCO3
20. stibbous carbonate
40. Li2SO4
Nomenclature: A Little Bit of Everything
1. oxygen gas
21. NaClO3
2. magnesium chloride
22. Sb2O5
3. tin (IV) carbonate
23. Sn(NO3)4
4. carbon monoxide
24. Na2O2
5. aluminum hydrogen sulphate
25. HgBr
6. copper (II) phosphate
26. Zn2C
7. cupric phosphide
27. H2
8. stibbous nitrate
28. Al2O3
9. argon gas
29. ZnCO3
10. ammonium hydroxide
30. Mg(HSO4)2
11. mercury (II) sulphate
31. Pb3N4
12. diantimony trioxide
32. AlF3
13. gold (I) hydrogen carbonate
33. He
14. beryllium sulphide
34. SO3
15. silver bromide
35. NaCl
16. boron oxide
36. N2
17. plumbous bromide
37. Pb(CO3)2
18. barium carbonate
38. CO2
19. mercuric carbide
39. (NH4)2O
20. chlorine gas
40. AuBr3
Nomenclature Rummy Rules
Group size: 3-5 players (6 if absolutely necessary)
Periodic Tables or “Official Cheat Sheets” are permitted during the game.
1.
Dealer deals 7 cards to each player.
2.
Place the remaining cards face down in a stack with the top card turned face up beside it as the
start of the discard pile.
3.
Each turn begins with picking up a card from either the stack or the discard pile.
4.
The person to the left of the dealer starts the game by picking up the face down card from the
stack or the face up card on the discard pile.
5.
If this player can lay down any COMPLETE MOLECULES, they can do so at this time. They
must correctly name each molecule as they lay it down (Jokers may be used for any ion.)
eg. CaCl2 = 1 card of Ca+2 and 2 cards of Cl-1
Ne = 1 card of Ne0
O2 = 2 cards of O-2
6.
This player ends their turn by placing a card from their hand onto the discard pile.
7.
The play then passes to the next person on the left.
8.
The round is completed when a player gets rid of all of their cards. They do NOT have to have
a discard.
9.
Once the round is scored (see below), the next person to the left becomes the new dealer.
DECOMPOSITION!
If a player lays down an incorrect molecule (ex they call out calcium chloride and place down 1
Ca+2 card and 1 Cl-1 card) or if they name the molecule incorrectly (don’t forget about multiple
valences), another player may yell DECOMPOSITION as the mistake is made. If there was
indeed a mistake, the player who made the mistake must pick up all of the molecules they put
down during that turn. Their turn is now over and they do NOT get to discard. If the molecule
was correct, then the player who yelled incorrectly yelled decomposition loses 10 points.
SCORING
 winner of the round: 10 bonus points
 incorrect decomposition: –10 points
 cards in your hand: joker –10, noble gas –5, all other cards are worth the negative of their
valence, ions with more than one charge are counted as the highest value against you
 cards on the table: all cards count as the positive of the valence used
Keep running total of all rounds played during the class. Make a grand total of all the rounds when
the time is up. Rank players 1st, 2nd, etc. and hand in your results.
name:
Nomenclature Rummy Practice
Cl-1
O-2
Ca+2
Hg+1 +2
H+1
Fe+2 +3
O-2
H+1
PO4-3
CO3-2
HSO4-1
Sb+3 +5
H+1
Al+3
Cl-1
Zn+2
NO3-1
Sn+2 +4
H+1
Cl-1
Cl-1
Build as many molecules as you can from the above nomenclature rummy cards. You can use
every card as many times as you like. You must write the formula and name for each molecule.
Fill in the top 8 rows first. Any molecules you make after this are BONUS .
Comparing the Masses of Reactants and Products
Purpose: To test the law of conservation of mass.
Materials:
 electronic balance
 250 mL beaker
 small test tube
 copper (II) sulphate solution
 sodium hydroxide solution
Procedure:
1. Pour 20 mL of sodium hydroxide solution into the beaker.
2. Pour copper (II) sulphate solution into the small test tube until the test tube is about half
full.
3. Place the test tube in the beaker, being careful not to mix the solutions.
4. Determine the mass of the beaker and its contents.
5. Record observations and measurements.
6. Carefully pour the copper (II) sulphate solution into the beaker, leaving the empty test tube
in the beaker.
7. Predict the mass of the beaker and its contents.
8. Record observations about the appearance of the contents and the measurement.
9. Flush contents of the beaker down the sink with water.
10. Clean the beaker and the test tube and return them to your teacher.
Discussion Questions:
1. When the solutions were mixed, did a chemical change occur? What is your evidence for
this?
2. How did the mass of the reactants and glassware before the reaction compare with the mass
after the reaction?
3. If you noted a change in mass, offer an explanation to account for it.
4. Do your results support the law of conservation of mass? If not, what sources of error
might have affected your observations and measurements?
5. The products of the reaction are copper (II) hydroxide and sodium sulphate. Write a word
equation to describe the reaction.
Counting Atoms Worksheet
Na2CO3
Type of atom
# of atoms
Total
K2CrO4
Type of atom
# of atoms
# of atoms
Total
# of atoms
4 Al2(CO3)3
Type of atom
# of atoms
Total
# of atoms
Total
10 (NH4)2S
Type of atom
3 BaCl2
Type of atom
Total
Total
Pb(NO3)2
Type of atom
# of atoms
Total
Total
NH4C2H3O2
Type of atom
Ca3(PO4)2
Type of atom
2 (NH4)2Cr2O7
Type of atom
# of atoms
Total
# of atoms
25 CH3COOH
Type of atom
Total
# of atoms
Balancing Equations
Chemical equations are used to show a chemical process or reaction. They are often written first as a word
equation and then translated to chemical formulae using your nomenclature skills.
One of the key concepts of a chemical reaction is the fact that both sides of the equation MUST balance.
This means that not only must both sides have the same total number of atoms, but that the amount of a
single type of atom on both sides must also balance.
Example: If you started with only two oxygen atoms, you must still have two oxygen atoms when
you finish. You cannot have less or more. So, the total number of oxygen atoms on the left side of
the equation (reactants - before the arrow) and the total number of oxygen atoms on the right side of
the equation (products - after the arrow) must be the same.
To make the number of atoms the same, the coefficient (number in front of each molecule) can be changed,
adding entire molecules to either side until the overall equation is balanced.
You cannot change the chemical formula – only leading numbers.
For this simulation, we will be using M&M’s to represent the different atoms. Do NOT start eating the
“atoms” until after you have completed the work. Once you are finished all the questions, you will be
permitted to eat your molecules.
Legend (you may substitute colours if necessary)
Hydrogen – yellow
Oxygen – red
Nitrogen – green
Chlorine – blue
Carbon – brown
Calcium – orange
Example:
word equation: oxygen + hydrogen  water
translates to
chemical equation:
O2 + H2  H2O
To model this, we place an oxygen molecule (2 reds) in the place of the first reactant and a hydrogen
molecule (2 yellows) in the place of the second reactant. In the space for the first product, place a water
molecule (2 yellows and a red). Since there isn’t a second product, leave that space blank.
a) Are there equal numbers of reds on both sides? NO
 Add an extra molecule that includes a red to the side that is missing some.
b) (Add one more water on the product side – two yellows and a red together.)
c) Are there equal numbers of yellows on both sides? NO
 Add an extra molecule that includes a yellow to the side that is missing some.
(Add one more hydrogen molecule – 2 yellows – to the reactant side)
d) Are there equal numbers of all colours on both sides? YES
e) Count the number of each molecule and place it in front of the molecules in the equation.
O2 + 2 H2  2 H2O
Balancing Equations
Reactant 1
+
Reactant 2

Product 1
+
Product 2
Practice
In each of the questions below, you are given the word equation, an unbalanced chemical equation
and a chemical equation with blanks. Use the M&M’s to help determine the number that belongs
in each of the blanks. Remember, the number represents the number of MOLECULES, not the
number of atoms (M&M’s).
1. methane + oxygen  carbon dioxide + water
CH4 + O2  CO2 + H2O
O2 
CH4 +
CO2 +
H2O
2. calcium + water  calcium hydroxide + hydrogen
Ca + H2O  Ca(OH)2 + H2
H2O 
Ca +
Ca(OH)2 +
H2
3. nitrogen + hydrogen  ammonia
N2 + H2  NH3
H2 
N2 +
NH3
4. calcium chloride + water  calcium hydroxide + hydrochloric acid
CaCl2 + H2O  Ca(OH)2 + HCl
CaCl2 +
H2O 
Ca(OH)2 +
HCl
5. nitric acid  hydrogen + nitrogen dioxide + oxygen
HNO3  H2 + NO2 + O2
HNO3 
H2 +
NO2 +
O2
Bonus:
6. oxygen + glucose  carbon dioxide + water
O2 + C6H12O6  CO2 + H2O
O2 +
C6H12O6 
CO2 +
H2O
Balancing Chemical Equations
H2 
1.
N2 +
NH3
2.
KClO3 
3.
NaCl +
4.
H2 +
5.
Pb(OH)2 +
6.
AlBr3 +
7.
CH4 +
8.
C3H8 +
O2 
9.
FeCl3 +
NaOH 
10.
P +
11.
Na +
12.
Ag2O 
13.
S8 +
14.
CO2 +
15.
K +
16.
HCl +
17.
HNO3 +
18.
H2O +
19.
NaBr +
20.
H2SO4 +
KCl +
O2
F2 
NaF +
O2 
Cl2
H2O
HCl 
H2O +
PbCl2
K2SO4 
KBr +
Al2(SO4)3
O2 
O2 
CO2 +
H2O
CO2 +
H2O
Fe(OH)3 +
NaCl
P2O5
H2 O 
NaOH +
Ag +
O2 
O2
SO3
H2O 
C6H12O6 +
MgBr2 
KBr +
CaCO3 
O2
Mg
CaCl2 +
NaHCO3 
O2 
H2
H2O +
NaNO3 +
H2O +
H2O2
CaF2 
NaNO2 
NaF +
HNO2 +
CaBr2
Na2SO4
CO2
CO2
More Balancing Equations
HCl 
1.
Ca(OH)2 +
2.
SO2 +
3.
Al +
4.
P4O10 +
5.
AgNO3 +
6.
Pb(NO3)2 
7.
Ag2O 
8.
Br2 +
9.
C5H12O +
10.
N2 +
11.
AsCl3 +
H2S 
12.
FeCl3 +
(NH4)2S 
13.
Na2SO4 +
14.
Cl2 +
15.
PCl3 +
16.
Mg3N2 +
17.
SiCl4 +
H2O 
HCl +
18.
SbCl5 +
H2O 
Sb2O5 +
19.
Al4C3 +
H2O 
Al(OH)3 +
20.
KOH +
SO2 
21.
C10H16 +
22.
Al +
23.
Al2(CO3)3 +
24.
K2SO4 +
25.
Cl2 +
26.
Pb(NO3)2 +
O2 
CaCl2 +
H2O
SO3
H2SO4 
Al2(SO4)3 +
H2O 
H2
H3PO4
CaCl2 
AgCl +
Ca(NO3)2
PbO +
NO2 +
O2
Ag +
O2
KI 
I2 +
O2 
O2 
KBr
CO2 +
H2O
N2O3
As2S3 +
Fe2S3 +
C 
H2O 
Na2S +
HCl +
H2O 
NH4Cl
CO2
HOCl
H3PO3 +
H2O 
HCl
Mg(OH)2 +
HCl
CH4
H2O
HCl +
H2SO4 
NH3
SiO2
K2SO3 +
Cl2 
C
Al2(SO4)3 +
H2O 
BaCl2 
NaBr 
HCl
Al(OH)3 +
BaSO4 +
Br2 +
NaBr 
H2O +
CO2
KCl
NaCl
PbBr2 +
NaNO3
SO2
Types of Chemical Reactions
Purpose: To study four different types of chemical reactions
Introduction: It is not easy to classify all chemical reactions precisely. Nevertheless, most
reactions can be classified into one of four major categories.
Synthesis Reactions – are reactions in which atoms and molecules join together directly to
produce larger molecules. An example of a synthesis reaction is the combustion of sulphur to form
sulphur dioxide:
S(s) + O2 (g)  SO2 (g)
Decomposition Reactions – are reactions that are just the opposite of synthesis reactions. One
larger molecule is broken up into several smaller elements or molecules. An example is the
decomposition of carbonic acid:
H2CO3 (aq)  H2O (g) + CO2 (g)
Single Displacement Reactions – are reactions that involve a change of partners between an
element and a molecule. An example is the liberation of bromine from calcium bromide by
reacting it with chlorine:
CaBr2 (aq) + Cl2 (g)  CaCl2 (aq) + Br2 (g)
Double Displacement Reactions – are reactions that involve the joint exchange of partners
between two molecules. An example of this is the precipitation of silver chloride when solutions
of silver nitrate and sodium chloride are mixed.
AgNO3 (aq) + NaCl (aq)  AgCl (s) + NaNO3 (aq)
In this experiment, you will study six different chemical reactions. You will identify some of the
products, and then classify each reaction as one of the four types. Assign each member in your
group the task of preparing a particular reaction and then take the time to observe each reaction
within the group.
Materials:
Bunsen burner, crucible tongs, 5 test tubes, medicine dropper, spatula, retort stand, test tube clamp,
10 mL graduated cylinder, test tube rack, wood splints, sodium sulphate solution, strontium
chloride solution, mossy zinc pieces, hydrochloric acid, manganese dioxide, aluminum pieces,
copper (II) chloride solution, copper wire, hydrogen peroxide, copper (II) sulphate pentahydrate
Procedure:
Lab A
Place a piece of aluminum in a test tube and cover it with a solution of copper (II) chloride. After
about 30 minutes, remove the aluminum and examine it.
1. Do you observe any change in the appearance of the aluminum?
2. If so, what is the charge?
3. Do you observe any change in the appearance of the copper (II) chloride solution?
4. If so, what is the change?
Lab B
Place 5 mL of sodium sulphate solution in a test tube. Add a medicine dropper full of strontium
chloride. Make observations and then leave the test tube aside for 20-30 minutes?
5. What do you observe right away? after 20-30 minutes?
6. If one of the products is sodium chloride (which is soluble in water), what is the other
product?
Lab C
Using crucible tongs, hold a 5 cm piece of copper wire in a burner flame. Remove the wire and
examine it. Dip it in strong hydrochloric acid that your teacher has to clean the wire and then
return it to the place you got it.
7. What do you observe?
8. With what did the copper combine?
Lab D
Place a small piece of mossy zinc in a test tube. Add 5 mL of dilute hydrochloric acid. Hold your
thumb over the mouth of the test tube.
9. What do you observe?
Insert a BURNING splint into the mouth of the test tube.
10. What do you observe?
Lab E
Place 3 mL of hydrogen peroxide in a test tube. Add a pinch of manganese dioxide using a
spatula. Hold your thumb over the mouth of the test tube.
11. What do you observe?
Insert a GLOWING splint into the mouth of the test tube.
12. What do you observe?
Water is one of the products of this reaction.
13. What is the other product of this reaction.
14. The manganese dioxide used in this experiment is a CATALYST. What is a catalyst? Is
the catalyst a reactant in the reaction?
Lab F
Place about 5-6 crystals of copper (II) sulphate pentahydrate in a clean, dry, heatable test tube.
Clamp the test tube to a retort stand and heat until no further changes occur.
15. What do you observe in the upper part of the test tube?
16. What do you observe in the lower part of the test tube?
17. What is the substance remaining in the test tube after heating?
Allow the test tube to cool and add 3 drops of water.
18. What do you observe?
Concluding Questions
1. Write word equations to describe the chemical reactions in each of the labs.
2. What TYPE of chemical reaction is illustrated in each of these reactions?
3. Write balanced chemical equations for each of the six reactions.
Balancing Equations Again?
Balance each skeletal equation provided. Identify each reaction as synthesis or decomposition.
H2O (l) 
1.
CaO(s) +
Ca(OH)2 (aq)
2.
FeCl3 (s) 
3.
P2O5 (s) +
4.
NH4NO3 (s) 
H2O (l) +
N2O(g)
5.
NaHCO3 (s) 
H2O (l) +
CO2 (g) +
6.
NO(g) +
7.
Ag2CO3 (s) 
Ag2O (s) +
CO2 (g)
8.
Ca(OH)2 (s) 
H2O (l) +
Ca (s) +
9.
CaCO3 (s) 
CaO (s) +
CO2 (g)
synthesis or decomposition?
10.
PbCO3 (s) 
PbO (s) +
CO2 (g)
synthesis or decomposition?
11.
Be Cl2 (s) +
O2 (g) 
Be(ClO3)2 (s)
synthesis or decomposition?
12.
CuO (s) 
O2 (g) +
Cu (s)
synthesis or decomposition?
13.
Mg (s) +
Cl2 (g) 
MgCl2 (s)
synthesis or decomposition?
Fe (s) +
H2O (l) 
O2 (g) 
synthesis or decomposition?
Cl2 (g)
synthesis or decomposition?
H3PO4 (aq)
synthesis or decomposition?
synthesis or decomposition?
Na2CO3 (s)
synthesis or decomposition?
NO2 (g)
synthesis or decomposition?
synthesis or decomposition?
O2 (g)
synthesis or decomposition?
Writing Displacement Reactions
Complete each equation by writing the correct products. State whether the reaction is a single or
double displacement reaction.
1. CaO (s) + Mg (s) 
single or double displacement
2. NaOH (aq) + HCl (aq) 
single or double displacement
3. LiCl (aq) + Na (s) 
single or double displacement
4. H2O (l) + Mg (s) 
single or double displacement
5. KOH (aq) + HBr (aq) 
single or double displacement
6. LiOH (aq) + HF (aq) 
single or double displacement
7. Ag2SO4 (s) + Cu (s) 
single or double displacement?
Writing and Balancing Chemical Equations
1.
solid aluminum + oxygen gas  solid aluminum oxide
2.
solid iron + solid sulphur  solid iron (III) sulphide
3.
solid calcium oxide + water  aqueous calcium hydroxide
4.
liquid hydrogen nitrate  water vapour + nitrogen dioxide gas + oxygen gas
5.
aqueous copper (II) nitrate + aqueous potassium hydroxide
 aqueous copper (II) hydroxide + aqueous potassium nitrate
6.
aqueous hydrogen chloride + aqueous sodium sulphate
 aqueous sodium chloride + water + sulphur trioxide gas
7.
aqueous sodium hydroxide + aqueous ammonium sulphate
 aqueous ammonium hydroxide + aqueous sodium sulphate
8.
solid cobalt + aqueous lead (IV) chlorate  aqueous cobalt (II) chlorate + solid lead
9.
aqueous iron (III) bromide + aqueous magnesium carbonate
 aqueous magnesium bromide + aqueous iron (III) carbonate
10.
solid tin (IV) hydroxide  solid tin (IV) oxide + water vapour
11.
solid sodium + water  hydrogen gas + aqueous sodium hydroxide
12.
aqueous aluminum phosphate + aqueous ammonium carbonate
 aqueous aluminum carbonate + aqueous ammonium phosphate