1. Name _________
THE TRIPLE AND FOUR
BEAM BALANCES
What masses are shown on each of the following balances?
Triple Beam Balance
0 100 200 300 ~00 500 600 700 800 QOO 1000
Answer:
0 2. 3 4 5 6 7 8 q 10
0 100 zoo 2100 400 500 600 700 800
t-trf'TT'II'Ir"!'h"rrnnnmTmlTmlrTTTTTTm;rrrniTTTTlrrnTTITnTITrnTTTTTTTT""TTTrrTITTTTTTTTTTT1rrrrr,.,.,-r--1 Answer:
0 2. 4 5 6 7 8 10
Four Beam Balance
Answer: 0 10
I
0 0.! 0.2. 0.3 0.4- 0.5 1.0
0
Answer:
0 0.! 0.2 0. 3 0.4 0.5 0.6 0.7 o .B o.C! J.O
Answer:
0 0. I 0.2.. 0.3 0.4- o.5 0.6 0.7 o.s o.q 1.0
Chemistry IF8766 3 ©Instructional Fair, Inc.

2. LA MEASURING LIQUID VOLUME Name _______
Lot What volume is indicated on each of the graduated cylinders below? The unit of vol
is ml.
----60 - - - - 3C
----4
----50 - - - - 2:
----3
a) b) c)
-----80
____ .;:
----75 - - - - -=-
----5 ----70
10 -----65
d) e) f)
-----40 -----4
-----30 -----3 - - - - 5:
-----20 -----2.
----10 - - - - <:.
g) h) i)
Cher Chemistry IF8766 4 ©lnstructiono "

3. Name _________
READING THERMOMETERS
What temperature is indicated on each of the thermometers below?
80 10
5
70 0
-5
60 -10
a) b) c)
IO 30
~0
0 10
0
-10 -10
d) e) f)
20 5 100
10 0 qq
0 -5 qB
-10 -10 97
-20 -15 q6
g) h) i)
Chemistry IF8766 5 ©Instructional Fair, Inc.

4. u ... ..,."'
-····-· . -·-·,.- .
(FACTOR LABEL METHOD)
_.,.~
. , ........ ,, '-" ------~
Lo t
Using this method, it is possible to solve many problems by using the re lationsh ip o~
to another. For example, 12 inches= one foot. Since these two numbers represe "·
same value, the fractions 12 in/1 ft and 1 ft/12 in are both equal to one. When yo ~
multiply another number by the number one, you do not change its value. Howe .
may change its unit.
Example 1: Convert 2 miles to inches.
2 miles x 5.280 ft x 12 inches = 126,720 in
1 mile 1 ft (Using significant figures,
2 mi = 100,000 in.)
Example 2: How many seconds are in 4 days?
4 days x 24 hrs x 60 min x 60 sec = 345,600 sec
1 day 1 hr 1 min (Using significant figures,
4 days =300,000 sec.)
Solve the following problems. Write the answers in significant figures.
1. 3 hrs = sec
2. 0.035 mg = cg
3. 5.5 kg = lbs
4. 2.5 yds = in
5. 1.3 yrs = hr ( l yr = 365 days)
6 . 3 moles = molecules (1 mole= 6.02 x 1023 molecules)
7. 2.5 x 1024 molecules = moles
8. 5 moles = _ _ _ liters (1 mole= 22.41iters)
9. lOO.Iiters = _ _ _ moles
10. 50. liters = _ _ _ molecules
ll . 5.0 x l 024 molecules = liters
12. 7.5xl03 ml = liters
Chemistry IF8766 6 ©lnstructiono
Cher

5. METRICS AND MEASUREMENT Name _________
In the chemistry classroom and lab, the metric system of measurement is used, so it is
important to be able to convert from one unit to another.
mega kilo hecto deca Basic Unit deci centi milli micro
(M) (k) (h) (da) gram (g) (d) (c) (m) (~)
1,000,000 1000 100 10 liter (L) .1 .01 .001 .000001
1os 103 102 10 1 meter (m) 1Q-1 1Q•2 1Q-3 1Q·6
Factor Label Method
1. Write the given number and unit.
2. Set up a conversion factor (fraction used to convert one unit to another).
a . Place the given unit as denominator of conversion factor.
b. Place desired unit as numerator.
c. Place a "1" in front of the larger unit.
d. Determine the number of smaller units needed to make "1" of the larger unit.
3. Cancel units. Solve the problem.
Example 1: 55 mm =__ m Example 2: 88 km =__ m
55 a>m I l m
1000_o:mi
= 0.055 m 88 J<m 11000 m
1jgli
= 88,000 m
Example 3: 7000 em =__ hm Example 4: 8 dal =__ dl
7000 em I 1m I l hm = 0.7 hm wx 10 dl = 800 dl
100 em 100 J'f'( 1.deft 1):
The factor label method can be used to solve virtually any problem including changes in
units. It is especially useful in making complex conversions dealing with concentrations
and derived units.
Convert the following.
1. 35 ml = dl 6. 4,500 mg = g
2. 950g = kg 7. 25cm = mm
3. 275mm = em 8. 0.005 kg = dog
4. 1,000 L = kl 9. 0.075 m = em
5. 1,000 ml = L 10. 15 g = mg
Chemistry IF8766 7 ©Instructional Fair, Inc.

6. L Name _________
SCIENTIFIC NOTATION
a
Scientists very often deal with very small and very large numbers, which can lead to a lot o·
confusion when counting zeros! We have learned to express these numbers as powers of
Scientific notation takes the form of M X 1on where 1 ~ M < 10 and n" represents the numb
of decimal places to be moved. Positive n indicates the standard form is a large number.
Negative n indicates a number between zero and one.
Example 1: Convert 1,500,000 to scientific notation.
We move the decimal point so that there is only
one digit to its left, a total of 6 places.
1,500,000 = 1.5 X 106
Example 2: Convert 0.000025 to scientific notation.
For this, we move the decimal point 5 places to the
right.
0.000025 = 2.5 X 10-5
(Note that when a number starts out less than one,
the exponent is always negative.)
Convert the following to scientific notation.
1. 0.005 = 6. 0.25 =
2. 5,050 = 7. 0.025 =
3. 0.0008 = 8. 0.0025 =
4. 1,000 = 9. 500 =
5. 1,000,000 = 10. 5,000 =
Convert the following to standard notation.
1. 1.5 X 103 = 6. 3.35 X 1Q·l =
2. 1.5 X 10·3 = 7. 1.2 X 10·4 =
3. 3.75 X 10·2 = 8. 1 X 1Q4 =
4. 3.75 X 102 = 9. 1 X 10·1 =
5. 2.2 X lOS = 10. 4x100 =
Chemistry IF8766 8 ©Instructional Fair, I -

7. SIGNIFICANT FIGURES Name -----------------
A measurement can only be as accurate and precise as the instrument that produced it.
>f
A scientist must be able to express the accuracy of a number, not just its numerical value.
10.
We can determine the accuracy of a number by the number of significant figures it
er contains.
1) All digits 1-9 inclusive are significant.
Example: 129 has 3 significant figures.
2) Zeros between significant digits are always significant.
Example: 5,007 has 4 significant figures.
3) Trailing zeros in a number are significant only if the
number contains a decimal point.
Example: 100.0 has 4 significant figures.
100 has 1 significant figure.
4) Zeros in the beginning of a number whose only function
is to place the decimal point are not significant.
Example: 0.0025 has 2 significant figures.
5) Zeros following a decimal significant figure are
significant.
Example: 0.000470 has 3 significant figures.
0.47000 has 5 significant figures.
Determine the number of significant figures in the following numbers.
1. 0.02 6. 5,000.
2. 0.020 7. 6,051.00
3. 501 8. 0.0005
4. 501 .0 9. 0.1020
5. 5,000 10. 10,001
Determine the location of the last significant place value by placing a bar over the digit.
(Example: 1. 700)
1• 8040 6, 901 100
2. 0.0300 7. 4.7 X 10-s
3. 699.5 8. 10,800,000.
4. 2.000 X 102 9. 3.01 X 1021
5. 0.90100 10. 0.000410
Chemistry IF8766 9 ©Instructional Fair, Inc.

8. PERCENTAGE ERROR Name _________
Percentage error is a way for scientists to express how far off a laboratory value is from the
commonly accepted value. cf "' 1 "" y _-_ __!.:~ "' u.)(r_.
1 1
The formula is: ~ (c
% error = ' AcceptedValu~erimental Value ,1~ x 100
- Accepted Value ~
-+
absolut~alue
Determine the percentage error in the following problems.
1' Experimental Value = 1.24 g
Accepted Value = 1.30 g
Answer:
2. Experimental Value = 1.24 x 1Q-2 g
Accepted Value = 9.98 x 10-3 g
Answer:
3. Experimental Value = 252 ml
Accepted Value = 225 ml
Answer:
4. Experimental Value = 22.2 L
Accepted Value = 22.4 L
Answer:
5. Experimental Value = 125.2 mg
Accepted Value= 124.8 mg
Answer:
C hemistry IF8766 11 ©Instructional Fair, Inc.

9. L TEMPERATURE AND Name _________
_(
ITS MEASUREMENT
Temperature (which measures average kinetic energy of the molecules) can be
measured using three common scales: Celsius, Kelvin and Fahrenheit. We use the
following formulas to convert from one scale to another. Celsius is the scale most
desirable for laboratory work. Kelvin represents the absolute scale. Fahrenheit is the old
English scale which is never used in lab.
oc = K - 273 K = oc + 273
oF= 9/s°C + 32 oc = 5 /9(°F- 32)
Complete the following chart. All measurements are good to 1o C or better.
450 K
294 K
225 K
Chemistry IF8766 12 ©Instructional Fe

10. Name ________
AlTER-SUBSTANCES
VS. MIXTURES
- matter can be classified as either a substance (element or compound) or a mixture
~eterogeneous or homogeneous).
MaHer
Substance Mixtures
can write chemical variable ratio
formula, homogeneous
Homogeneous Heterogeneous
Element Compound
solutions colloids and
one type two or more different suspensions
atom atoms chemically
bonded
~! ossify each of the following as to whether it is a substance or a mixture. If it is a
substance, write Element or Compound in the substance column. If it is a mixture, write
eterogeneous or Homogeneous in the mixture column.
Type of Matter Substance Mixture
-
1. chlorine
2. water
3. soil
4. sugar water
5. oxygen
6. carbon dioxide
7. rocky road ice cream
8. alcohol
9. pure air
10. iron
Chemistry IF8766 17 ©Instructional Fair, Inc.

11. l PHYSICAL VS. CHEMICAL Name -----------------
PROPERTIES
A physical property is observed with the senses and can be determined without destroying
the object. For example, color, shape, mass, length and odor are all examples of physical
properties.
A chemical property indicates how a substance reacts with something else. The original
substance is fundamentally changed in observing a chemical property. For example, the
ability of iron to rust is a chemical property. The iron has reacted with oxygen, and the
original iron metal is changed. It now exists as iron oxide, a different substance .
Classify the following properties as either chemical or physical by putting a check in the
appropriate column.
Physical Chemical
Property Property
1. blue color
2. density
3. flammability
4. solubility
5. reacts with acid to form H2
6. supports combustion
7. sour taste
8. melting point
9. reacts with water to form a gas
10. reacts with a base to form water
11. hardness
12. boiling point
13. can neutralize a base
14. luster
15. odor
~~ em1stry IF8766 18 ©Instructional Fair, Inc

12. PHYSICAL VS. Name _________
CHEMICAL CHANGES
lg In a physical change, the original substance still exists, it has only changed in form. In a
al chemical change, a new substance is produced. Energy changes always accompany
chemical changes.
Classify the following as being a physical or chemical change.
e
l. Sodium hydroxide dissolves in water. _ _ _ _ _ _ __
2. Hydrochloric acid reacts with potassium hydroxide to produce a salt, water and
heat. ___________
3. A pellet of sodium is sliced in two. __________
4. Water is heated and changed to steam. _ _ _ _ _ _ _ __
5. Potassium chlorate decomposes to potassium chloride and oxygen gas.
6. Iron rusts. _ _ _ _ _ _ _ __
7. When placed in H20, a sodium pellet catches on fire as hydrogen gas is liberated and
sodium hydroxide forms.
8. Evaporation
9. Ice melting
l 0. Milk sours.
l l . Sugar dissolves in water.
12. Wood rotting
13. Pancakes cooking on a griddle
14. Grass growing in a lawn
15. A tire is inflated with air.
16. Food is digested in the stomach.
17. Water is absorbed by a paper towel.
, Inc. Chemistry IF8766 19 ©Instructional Fair, Inc .

13. ELEMENT SYMBOLS Name -----------------
An element symbol can stand for one atom of the element or one mole of atoms of the
element. (One mole = 6.02 x 1023 atoms of an element.)
Write the symbol for the following elements.
1' oxygen 11 ' plutonium
2. hydrogen 12. americium
3. chlorine 13. radium
4. mercury 14. germanium
5. fluorine 15. zinc
6. barium 16. arsenic
7. helium 17. lead
8. uranium 18. iron
9. radon 19. calcium
10. sulfur 20. cobalt
Write the name of the element that corresponds to each of the following symbols.
21 . Kr 31. Cu _____________
22. K 32. Ag
23 . c 33 . p
24. Ne 34. Mn
25. Si 35.
26. Zr 36. Au
27 . Sn 37. Mg
28. Pt 38. Ni
29. No 39. Br
30. AI 40. Hg
Chemistry IF8766 26 ©Instructional Fair, Inc.

14. Name _________
ATOMIC STRUCTURE
..., atom is made up of protons and neutrons (both found in the nucleus) and electrons
~t he surrounding electron cloud). The atomic number is equal to the number of protons.
--e mass number is equal to the number of protons plus neutrons. In a neutral atom, the
"mber of protons equals the number of electrons. The charge on an ion indicates an
~o alance between protons and electrons. Too many electrons produces a negative
: arge, too few, a positive charge .
-~ s structure can be written as part of a chemical symbol.
Example: mass Always write the magnitude of the charge fol-
number h lowed by the sign: N3+
+ c arge
3
7 7 protons Can you find the other mis-
71
8 neutrons ( 15 - 7) takes on this page?
atomic
number 4 electrons
:::omplete the following chart.
1
Element/ Atomic Mass
Atomic Mass Protons Neutrons Electrons
Ion Number Number
H
H+
12c
6
7Lj+
3
3scl-
17
39K
19
24Mg2+
12
As 3-
Ag
Ag+1
s-2
u
:: :: -emistry IF8766 27 ©Instructional Fair, Inc.

15. Name _________
ISOTOPES AND AVERAGE E
ATOMIC MASS (I
Elements come in a variety of isotopes, meaning they are made up of atoms with the -.
-
same atomic number but different atomic masses. These atoms differ in the number
of neutrons.
The average atomic mass is the weighted average of all the isotopes of an element.
Example: A sample of cesium is 75% 133Cs, 20% 132 Cs and
5% 134Cs. What is its average atomic mass?
Answer: .75 x 133 = 99.75
.20 X 132 = 26.4
.05 X 134 = 6.7
Total = 132.85 amu =average atomic mass
Determine the average atomic mass of the following mixtures of isotopes.
1. 80% 1271, 17% 1261, 3% 1281
197 198
2. 50% Au, 50% Au
56
3. 15% 55 Fe, 85% Fe
4. 99% 1H, 0.8% 2 H, 0.2% 3 H
14 15 N, 16 N
5. 95% N, 3% 2%
6. 98% 12C, 2% 14C
Chemistry IF8766 28 ©Instructional Fair, Inc.

16. Name _________
ALENCE ELECTRONS
-- e valence electrons are the electrons in the outermost principal energy level. They are
.·;ays "s" or "sand p" electrons. Since the total number of electrons possible ins and p
_::>l evels is eight, there can be no more than eight valence electrons.
: e ermine the number of valence electrons in the atoms below.
Example: carbon
Electron configuration is 1s2 I 2s2 2p 2 1.
Carbon has 4 valence electrons.
. fluorine 11. lithium
2. phosphorus 12. zinc
3. calcium 13. carbon
4. nitrogen 14. iodine
5. iron 15. oxygen
6. argon 16. barium
7. potassium 17. aluminum
8. helium 18. hydrogen
9. magnesium 19. xenon
10. sulfur 20. copper
dr,lnc. Chemistry IF8766 31 ©Instructional Fair, Inc.

17. Name _________
PERIODIC TABLE WORKSHEET
1. Where are the most active metals located? _ _ _ _ _ _ _ _ _ _ _ _ _ __
2. Where are the most active nonmetals located? _ _ _ _ _ _ _ _ _ _ _ _ __
3. As you go from left to right across a period, the atomic size Cdecreases 1 increases).
Why? _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ ___
4. As you travel down a group, the atomic size (decreases I increases). Why?
5. A negative ion is ( larger I smaller) than its parent atom.
6. A positive ion is (larger I smaller) than its parent atom.
7. As you go from left to right across a period, the first ionization energy generally
(decreases I increases). Why? _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ __
8. As you go down a group, the first ionization energy generally (decreases 1 increases).
Why? _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ __
9. Where is the highest electronegativity found? _ _ _ _ _ _ _ _ _ _ _ _ __
10. Where is the lowest electro negativity found? _ _ _ _ _ _ _ _ _ _ _ _ __
11. Elements of Group 1 are called _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ __
12. Elements of Group 2 are called - - - - - - - - - - - - - - - - - - -
13. Elements of Group 3-12 are called _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ __
14. As you go from left to right across the periodic table, the elements go from
(metals I nonmetals) to (metals I nonmetals).
15. Group 17 elements are called _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ __
16. The most active element in Group 17 is _ _ _ _ _ _ _ _ _ _ _ _ _ _ __
17. Group 18 elements are called _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ __
18. What sublevels are filling across the Transition Elements? _ _ _ _ _ _ _ _ _ __
19. Elements within a group have a similar number of _ _ _ _ _ _ _ _ _ _ _ __
20. Elements across a series have the same number of _ _ _ _ _ _ _ _ _ _ __
21. A colored ion generally indicates a _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ __
22. As you go down a group, the elements generally become (more I less) metallic.
23. The majority of elements in the periodic table are (metals I nonmetals).
24. Elements in the periodic table are arranged according to their _ _ _ _ _ _ __
25. An element with both metallic and nonmetallic properties is called a _ _ _ __
Chemistry IF8766 36 ©Instructional Fair, Inc.

18. Name ________
PERIODIC TABLE PUZZLE
1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18
r--
QJ
I
F G H
B A
c E J
iS).
! •
IIIDIIIIIIIIIIII
Place the letter of each of the above elements next to its description below. ~.--.
1. An alkali metal
2. An alkaline earth metal
3. An inactive gas _ _
4. An active nonmetal
5. A semi-metal
6. An inner transition element
7. Its most common oxidation state is -2.
8. A metal with more than one oxidation state
9. Metal with an oxidation number of +3
10. Has oxidation numbers of + 1 and -1
c. Chemistry IF8766 37 ©Instructional Fair, Inc.

19. Name _________
IONIC BONDING
Ionic bonding occurs when a metal transfers one or more electrons to a nonmetal in an
effort to attain a stable octet of electrons. For example, the transfer of an electron from
sodium to chlorine can be depicted by a Lewis dot diagram .
••
Calcium would need two chlorine atoms to get rid of its two valence electrons.
.. .~··
:CI· + ·Ca
.. ~
+ .cl:..
Show the transfer of electrons in the following combinations.
l. K + F
2. Mg +I
3. Be + S
4. Na + 0
5. AI + Br
Chemistry IF8766 38 ©Instructional Fair,

20. COVALENT BONDING Name _________
ovalent bonding occurs when two or more nonmetals share electrons, attempting to
attain a stable octet of electrons at least part of the time . For example:
XX Note that hydrogen
H· + xCI~ is content with 2,
XX
not 8, electrons.
Show how covalent bonding occurs in each of the following pairs of atoms. Atoms may
share one, two or three pairs of electrons.
,I c Chemistry IF8766 39 ©Instructional Fair, Inc.

21. Name ________
TYPES OF CHEMICAL BONDS
c
Classify the following compounds as ionic (metal + nonmetal), covalent (nonmetal +
nonmetal) or both (compound containing a polyatomic ion).
1. CaCI 2 11. MgO
13. HCI
4. BaS04 14. Kl
15. NaOH
6. NaF
17. AIP0 4
18. FeCI 3
10. LiBr
Chemistry IF8766 40 ©Instructional Fair, lr

22. Name _ _ _ _ _ _ __
WRITING FORMULAS
(CRIS$-CROSS METHOD)
Write the formulas of the compounds produced from the listed ions.
Cl· co3-2 OH· so4-2 PO4-3 NO·
3
Na+
NH+
4
K+
Ca+2
Mg+2
Zn+2
Fe+3
Al+3
Co+3
Fe+2
H+
Chemistry IF8766 44 ©Instructional Fair, Inc.

23. Name ________
AMING IONIC COMPOUNDS
.ame the following compounds using the Stock Naming System.
1' CaC03
2. KCI
3. FeS04
4. LiBr
5. MgCI2
6. FeCI 3
7. Zn/P04 2
)
8. NH 4 3
N0
9. AI(OH) 3
·o. CuC 2 H30 2
' 1' PbS03
2. NaCI03
3. CaC 20 4
14. Fe20 3
15. (NH4)p04
16. NaHS04
17. Hg 2CI 2
18. Mg(N02) 2
19. CuS0 4
20. NaHC03
21' NiBr3
22. Be(N0 3) 2
23. ZnS04
24. AuCI 3
25. KMn0 4
·.Inc. Chemistry IF8766 45 ©Instructional Fair, Inc.

24. Name ________
NAMING MOLECULAR COMPOUNDS N
Name the following covalent compounds.
1. C02
2. co
3. S0 2
4. S0 3
5. N20
6. NO
7. N203
8. N02
9. N204
10. N20s
11. PCI 3
12. PCI 5
13. NH 3
14. SCI 6
15. P20s
16. CCI4
17. Si02
18. CS2
19. OF2
20. PBr3
Chemistry IF8766 46 ©Instructional Fair, Inc.

25. NAMING ACIDS Name
Name the following acids.
1. HN03
2. HCI
3. H2S04
4. H2S03
5. HC2 H30 2
6. HBr
7. HN02
8. Hp04
9. H2S
10. H2C03
Write the formulas of the following acids.
11. sulfuric acid
12. nitric acid
13. hydrochloric acid
14. acetic acid
15. hydrofluoric acid
16. phosphorous acid
17. carbonic acid
18. nitrous acid
19. phosphoric acid
20. hydrosulfuric acid
, Inc. Chemistry IF8766 47 ©Instructional Fair, Inc.

26. Name _________
WRITING FORMULAS FROM NAMES
Write the formulas of the following compounds.
l . ammonium phosphate
2. iron (II) oxide
3. iron (Ill) oxide
4. carbon monoxide
5. calcium chloride
6. potassium nitrate
7. magnesium hydroxide
8. aluminum sulfate
9. copper (II) sulfate
10. lead (IV) chromate
11 . diphosphorus pentoxide
12. potassium permanganate
13. sodium hydrogen carbonate
14. zinc nitrate
15. aluminum sulfite
Chemistry IF8766 48 ©Instructional Fair, Inc.

27. Name _________
GRAM FORMULA MASS
etermine the gram formula mass (the mass of one mole) of each compound below.
1. KMn04
2. KCI
3. Na 2 4
S0
4. )
Ca(N0 3 2
(S0 )
5. AI 2 4 3
6. )
(NH 4 3P04
7. CuS04•5H 20
8. Mg3(P04)2
9. Zn(C 2H30 2 2•2H 2
) 0
10. Zn 3(P04 2•4H 2
) 0
11. H2 3
C0
12. Hg2 2
Cr 0 7
13. Ba(CI0) 2
)
14. Fe2(S0 3 3
15. NH 4 2H3 2
C 0
nc . Chemistry IF8766 49 ©Instructional Fair, Inc.

28. Name ________
MOLES AND MASS T
Determine the number of moles in each of the quantities below.
1. 25 g of NaCI
2. 125 g of H2S04
-
3. 100. g of KMn0 4
4. 74 g of KCI
-
5. 35 g of CuS0 4 •5H 20
Determine the number of grams in each of the quantities below.
1. 2.5 moles of NaCI
2. 0.50 moles of H2SO 4
1.70 moles of KMnO 4
-
3.
4. 0.25 moles of KCI
5. 3.2 moles of CuS04 •5H 20
Chemistry IF8766 50 ©Instructional Fair, Inc .

29. Name _________
HE MOLE AND VOLUME
: ·gases at STP (273 K and 1 atm pressure), one mole occupies a volume of 22.4 L. What
: me will the following quantities of gases occupy at STP?
1.00 mole of H2
2 3.20 moles of 0 2
0.750 mole of N2
1.75 moles of C02
o. 0.50 mole of NH 3
6. 5.0 g of H2
7. 100. g of 0 2
8. 28.0 g of N2
9. 60. g of C02
10. 10. g of NH 3
::::hemistry IF8766 51 ©Instructional Fair. Inc.

30. Name _________
THE MOLE AND
AVOGADRO'S NUMBER
One mole of a substance contains Avogadro's Number (6.02 x 1023 ) of molecules.
How many molecules are in the quantities below?
1. 2.0 moles
2. 1.5 moles
3. 0.75 mole
4. 15 moles
5. 0.35 mole
How many moles are in the number of molecules below?
1. 6.02 X 1023
2. 1.204 X 1024
3. 1.5 X 1020
4. 3.4 X 1026
5. 7.5x10 19
Chemistry IF8766 52 ©Instructional Fair, Inc.

31. Name _________
MIXED MOLE PROBLEMS
Solve the following problems.
m 1' How many grams ore there In 1.5 X 1025 molecules of co,?
2. What volume would the C0 2 in Problem 1 occupy at STP?
3. A sample of NH 3 gas occupies 75.0 liters at STP. How many molecules is this?
4. What is the mass of the sample of NH 3 in Problem 3?
5. How many atoms are there in 1.3 x 1Q22 molecules of N02 ?
6. A 5.0 g sample of 0 2 is in a container at STP. What volume is the container?
7. How many molecules of 0 2 are in the container in Problem 6? How many atoms
of oxygen?
Jlr, Inc. Chemistry IF8766 53 ©Instructional Fair, Inc.

32. Name ________
PERCENTAGE COMPOSITION
Determine the percentage composition of each of the compounds below.
1. KMn0 4
K = ----
Mn = ____
0 = ----
2. HCI
H = ----
CI = _ _ __
3. Mg(N0) 2
Mg = - - - -
N = ----
0 = ----
4. (NH 4) 3 P04
N = ----
H = ----
p = ----
0 = ----
5. AI 2 (SO4) 3
AI= _ _ __
s= ----
0 = ----
Solve the following problems.
6. How many grams of oxygen can be produced from the decomposition of 100. g
of KCI0 ? _ _ _ _ __
3
7. How much iron can be recovered from 25.0 g of Fe 20 3? _ _ _ _ __
8. How much silver can be produced from 125 g of Ag 2S? _ _ _ _ __
Chemistry IF8766 54 ©Instructional Fair, I

33. Name _________
DETERMINING
EMPIRICAL FORMULAS
What is the empirical formula (lowest whole number ratio) of the compounds below?
1. 75% carbon , 25% hydrogen
2. 52.7% potassium, 47 .3% chlorine
3. 22.1% aluminum , 25.4% phosphorus, 52.5% oxygen
4. 13% magnesium, 87% bromine
5. 32.4% sodium, 22.5% sulfur, 45.1% oxygen
.g 6. 25.3% copper, 12.9% sulfur, 25.7% oxygen, 36.1% water
Jl Fair, lr Chemistry IF8766 55 ©Instructional Fair, Inc.

34. Name ________
DETERMINING MOLECULAR
FORMULAS (TRUE FORMULAS)
Solve the problems below.
1. The empirical formula of a compound is N02 . Its molecular mass is 92 g/mol.
What is its molecular formula?
2. The empirical formula of a compound is CH 2 • Its molecular mass is 70 g/mol.
What is its molecular formula?
J
3. A compound is found to be 40.0% carbon, 6.7% hydrogen and 53.5% oxygen.
Its molecular mass is 60. g/mol. What is its molecular formula?
4. A compound is 64.9% carbon, 13.5% hydrogen and 21.6% oxygen. Its molecular
mass is 74 g/mol. What is its molecular formula?
5. A compound is 54.5% carbon , 9.1% hydrogen and 36.4% oxygen. Its molecular
mass is 88 g/mol. What is its molecular formula?
Chemistry IF8766 56 ©Instructional Fa ir

35. BALANCING CHEMICAL EQUATIONS N a m e - - - - -
Rewrite and balance the equations below.
1. N + H --+ NH _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ __
2 2 3
2. KCI03 --+ KCI + 0 2- - - - - - - - - - - - - - - - - - - - - - -
3. NaCI + F
2
--+ NaF + Cl _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ __
2
4. H2 + 0 2 --+ H2 - - - - - - - - - - - - - - - - - - - - - - - -
0
5. AgN0 3 + MgCI2 --+ AgCI + Mg(N0)2_ _ _ _ _ _ _ _ _ _ _ _ _ __
7. CH 4 + 0 2 --+ C02 + H20
8. C3Ha + 02 --+co 2 + H20
9. CaH1a + 02 --+ C0 2 + H20
10. FeCI 3 + NaOH --+ Fe(OH)3 + NaCI
11 0 p + 02 --+ P20s
12. No+ H20 --+ NaOH + H2
13. Ag 20 --+ Ag + 02
14. Sa + 02 --+ S03
15. C0 2 + Hp --+ C6H,206 + 02
16. K + MgBr2 --+ KBr + Mg
17. HCI + CaC0 3 --+ CaCI 2 + H2 + C0 2_ _ _ _ _ _ _ _ _ _ _ _ _ _ __
0
Chemistry IF8766 58 ©Instructional Fair, Inc

36. Name _________
WORD EQUATIONS
Write the word equations below as chemical equations and balance.
1. zinc + lead (II) nitrate yield zinc nitrate + lead
2. aluminum bromide + chlorine yield aluminum chloride + bromine
3. sodium phosphate + calcium chloride yield calcium phosphate +
sodium chloride
4. potassium chlorate when heated yields potassium chloride + oxygen gas
5. aluminum + hydrochloric acid yield aluminum chloride + hydrogen gas
6. calcium hydroxide + phosphoric acid yield calcium phosphate + water
7. copper + sulfuric acid yield copper (II) sulfate + water + sulfur dioxide
8. hydrogen + nitrogen monoxide yield water + nitrogen
1C Chemistry IF8766 59 ©Instructional Fair, Inc.
---' . . -:?.o~

37. Name _________
CLASSIFICATION OF
CHEMICAL REACTIONS
Classify the reactions below as synthesis, decomposition, single replacement (cationic or
anionic) or double replacement.
3. Zn + H2S0 4 -- ZnS04 + H2
5. 2Hg0 -- 2Hg + 0 2
6. 2KBr + Cl 2 -- 2KCI + Br2
8. AgN0 3 + NaCI -- AgCI + NaN03
Chemistry IF8766 60 ©Instructional Fair, Inc.

38. Name _________
PREDICTING PRODUCTS
OF CHEMICAL REACTIONS
Predict the products of the reactions below. Then, write the balanced equation and
classify the reaction.
1. magnesium bromide + chlorine
2. aluminum + iron (Ill) oxide
3. silver nitrate + zinc chloride
4. hydrogen peroxide (catalyzed by manganese dioxide)
5. zinc + hydrochloric acid
6. sulfuric acid + sodium hydroxide
7. sodium + hydrogen
8. acetic acid + copper
Chemistry IF8766 61 ©Instructional Fair, Inc.

39. Name _________
STOICHIOMETRY:
MOLE-MOLE PROBLEMS
1. N2 + 3H 2 ~ 2NH 3
How many moles of hydrogen are needed to completely react with two moles
of nitrogen?
2. 2KCI03 ~ 2KCI + 30 2
How many moles of oxygen are produced by the decomposition of six moles
of potassium chlorate?
3. Zn + 2HCI ~ ZnCI 2 + H2
How many moles of hydrogen are produced from the reaction of three moles
of zinc with an excess of hydrochloric acid?
4. C 3 H8 + 50 2 ~ 3C0 2 + 4H 20
How many moles of oxygen are necessary to react completely with four moles of
propane (C 3 H8)?
5. Kl0 4 + AI(N0) 3 ~ 3KN0 3 + AIP0 4
How many moles of potassium nitrate are produced when two moles of potassium
phosphate react with two moles of aluminum nitrate?
Chemistry IF8766 62 ©Instructional Fair, Inc.

40. Name _________
STOICHIOMETRY:
VOLUME-VOLUME PROBLEMS
1. N2 + 3H 2 ---+ 2NH 3
What volume of hydrogen is necessary to react with five liters of nitrogen to produce
ammonia? (Assume constant temperature and pressure.)
2. What volume of ammonia is produced in the reaction in Problem 1?
3. C 3 H8 + 502 ---+ 3C0 2 + 4H 20
If 20 liters of oxygen are consumed in the above reaction, how many liters of carbon
dioxide are produced?
4. 2H 20 ---+ 2H 2 + 0 2
If 30 ml of hydrogen are produced in the above reaction, how many milliliters of
oxygen are produced?
5. 2CO + 0 2 ---+ 2C0 2
How many liters of carbon dioxide are produced if 751iters of carbon monoxide
are burned in oxygen? How many liters of oxygen are necessary?
1
C Chemistry IF8766 63 ©Instructional Fair, Inc.

41. Name _________
STOICHIOMETRY:
MASS-MASS PROBLEMS
1. 2KCI03 --+ 2KCI + 30 2
How many grams of potassium chloride are produced if 25 g of potassium chlorate
decompose?
2. N2 + 3H 2 --+ 2NH 3
How many grams of hydrogen are necessary to react completely with 50.0 g of
nitrogen in the above reaction?
3. How many grams of ammonia are produced in the reaction in Problem 2?
4. 2AgN0 3 + BaCI 2 --+ 2AgCI + Ba(N0) 2
How many grams of silver chloride are produced from 5.0 g of silver nitrate reacting
with an excess of barium chloride?
5. How much barium chloride is necessary to react with the silver nitrate in Problem 4?
Chemistry IF8766 64 ©Instructional Fair, Inc.

42. Name _________
STOICHIOMETRY:
MIXED PROBLEMS
1. N2 + 3H 2 -- 2NH 3
What volume of NH 3 at STP is produced if 25.0 g of N2 is reacted with an excess
of H2 ?
2. 2KCI03 -- 2KCI + 30 2
If 5.0 g of KCI03 is decomposed, what volume of 0 2 is produced at STP?
3. How many grams of KCI are produced in Problem 2?
4. Zn + 2HCI -- ZnCI 2 + H2
What volume of hydrogen at STP is produced when 2.5 g of zinc react with an
excess of hydrochloric acid?
1g 5. H2S04 + 2NaOH -- H20 + Na 2S04
How many molecules of water are produced if 2.0 g of sodium sulfate are
produced in the above reaction?
1? 6. 2AICI 3 -- 2AI + 3CI 2
If 10.0 g of aluminum chloride are decomposed, how many molecules of Cl 2
are p roduced?
ir, Inc . Chemistry IF8766 65 ©Instructional Fair, Inc.

43. Name _________
STOICHIOMETRY:
LIMITING REAGENT
l . N2 + 3H 2 ~ 2NH 3
How many grams of NH 3 can be produced from the reaction of 28 g of N2 and
25 g of H)
2. How much of the excess reagent in Problem l is left over?
3. Mg + 2HCI ~ MgCI 2 + H2
What volume of hydrogen at STP is produced from the reaction of 50.0 g of Mg and
the equivalent of 75 g of HCI?
4. How much of the excess reagent in Problem 3 is left over?
5. 3AgN0 3 + Nap04 ~ Agp0 4 + 3NaN0 3
Silver nitrate and sodium phosphate are reacted in equal amounts of 200. g each.
How many grams of silver phosphate are produced?
6. How much of the excess reagent in Problem 5 is left?
Chemistry IF8766 66 ©Instructional Fair, Inc