Download Exercise 1.8 - All Chemistry

Matter and Scientific Measurement
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Unit 1: Matter and Scientific Measurement
(Link to Prentice Hall Text: Chapters 2, 3 and 4)
Name:___________________________________________________________________________
Assignments
Page Number, Problem Numbers
Assignment 1
Memorize the conversion between
symbols and names for elements 1-38,
47, 53-56, 78-80, 82, 85-88. Use the
Reference Periodic Table (Table S) to
establish the conversions.
Assignment 2: Identifying Mixtures;
Physical/Chemical Changes
47: 30, 32, 34, 35
Assignment 3: Separating Mixtures
47: 38, 41
Assignment 4: Calculations with
Significant Digits
78: 41, 44
Assignment 5: Calculations with Density 79: 61
A. Matter, Change and The Governing Law of Chemistry
Matter and Scientific Measurement
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Chemistry is the study of how “stuff” (matter) changes. Chemists are particularly interested in how matter
changes at the particle level, on a scale that is much too small to see. Because the scale on which chemists seek to
understand change cannot be seen, chemists often perform experiments on an observable scale and then use their
evidence to make inferences about change at the particle level.
Chemists point to two types of change: physical and chemical.
When a change alters a material without changing its composition, or identity, it is described as a physical change.
Physical changes include cutting, grinding, bending, melting, freezing, condensing, boiling, splitting, cracking,
dissolving, crushing, etc.
When a change fundamentally changes the identity of a material, it is called a chemical change. Chemical changes
include burning, rotting, rusting, decomposing, fermenting, exploding, corroding, etc. Chemical changes are often
accompanied by a color change, an odor change, flames, a formation of a gas or a formation of a solid from two
liquid solutions.
Exercise 1.1
Classify the following as being either a chemical or a physical change:
1. Sodium chloride dissolves in water. ______
2. Hydrochloric acid reacts with sodium hydroxide to produce a salt, water and heat. ______
3. A pellet of sodium is sliced in half. ______
4. Water is heated and changed to steam. ______
5. Food is digested. ______
6. Starch molecules are formed from smaller glucose molecules. ______
7. Ice melts. ______
8. Plant leaves lose water through evaporation. ______
9. A red blood cell placed in distilled water swells and bursts. ______
10. The roots of a plant absorb water. ______
11. Iron rusts. ______
12. A person cools by sweating. ______
13. A match burns. ______
Exercise 1.2
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Practice with the Law of Conservation of Matter
The Law of Conservation of Matter states: ____________________________________________
_______________________________________________________________________________
_______________________________________________________________________________
The Law of Conservation of Energy states: ____________________________________________
_______________________________________________________________________________
_______________________________________________________________________________
Put a (P) in the blank if the statement is probable. Put a (I) in the blank if the statement is improbable.
1.
If a fire is hot enough, it could convert garbage into pure energy. ________
2.
When wood burns in a fireplace, the ashes have the same mass as the wood before it was burned.
________
3.
When you exercise and burn calories, fat is converted to energy and the mass disappears. ________
4.
When a tree grows, it converts the sunlight into the mass that becomes the bark, the trunk, the branches
and the leaves. ________
5.
When you were a baby, developing in utero, you grew because energy in your mother’s body was
converted into flesh. ________
B. Four Properties of Matter
Physical:_______________________________________________________
Chemical:______________________________________________________
Extensive: ________________________________________________________
Intensive: ________________________________________________________
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Exercise 1.4
Classify each of the following substances as a pure substance (P) or as a
mixture (M). If the substance is pure, classify it as an element (E) or a
compound (C). If the substance is a mixture, classify it as a
heterogeneous (HT) or homogeneous (HO) mixture.
1. Sodium (Na)_________
2. Water (H2O) _________
3. Soil _________
4. Coffee_________
5. Oxygen (O2)_________
6. Alcohol (CH3CH2OH)_________
7. Carbon Dioxide (CO2)_________
8. Cake Batter_________
9. Air (Mixture of N2, O2, CO2 and Ar)_________
10. Soap_________
11. Iron (Fe) _________
12. Salt Water_________
13. Ice Cream_________
14. Nitrogen (N2) _________
15. Eggs_________
16. Blood_________
17. Table Salt (NaCl) _________
18. Methane (CH4) _________
19. Gold (Au) _________
20. Palladium (Pd) _________
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C. Separation of Mixtures
Miscible means that the components of a mixture dissolve each other. Immiscible liquids do not dissolve each
other.
(1) Filtration
a. Solid in a Liquid
b. Solid in a Gas
c. Liquid in a Liquid
(2) Distillation
a. Solid in a Liquid
b. Liquid in a Liquid
(3) Chromatography
a. Separating Liquids
Matter and Scientific Measurement
D. Units
Quantity
Symbol
Unit Name
Unit Symbol
Length
l
meter
m
Mass
m
gram
g
Time
t
second
s
Temperature
T
Kelvin
K
Amount of
substance
n
mole
mol
electric current
I
Ampere
A
luminous intensity
Iv
Candela
cd
Quantity
Symbol
Unit Name
Unit
Derivation
Area
A
square meters
m2
A=lXw
Volume
V
cubic meters
m3
V=lXwXh
Density
d
grams per mL
g/mL
d = m/V
Molar mass
MM
grams per mole
g/mol
MM = m/n
Concentration
c
moles per liter
M
M = mol/V
Molar volume
Vm
liters per mole
L/mol
Vm = V/mol
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Matter and Scientific Measurement
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E. Significant Digits in Measurements
Significant figures – These are all the digits you know for sure, plus one place that is an estimate.
Uncertainty – Limit of precision of the reading (based on your ability to estimate the final digit). See
examples below.
Rules for zeros: All zeros count except placeholder zeros – these are the ones that disappear when you
write the number in scientific notation. Examples:
93,000,000 = 9.3 x 107 2 sf’s
0.000372 = 3.72 x 10-4 3 sf’s
0.0200 = 2.00 x 10-2
3 sf’s
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Exercise 1.5
For each of the following, write the scale reading, then the number of significant figures in the reading.
Reading
1.
1.
2.
2.
3.
3.
4.
4.
5.
5.
6.
6.
7.
7.
8.
8. -
SF’s
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For each of the volume devices below record the scale reading and indicate the uncertainty in the measurement.
scale
reading ________
uncertainty
________
scale
reading ________
scale
reading ________
uncertainty
________
10
9
uncertainty
________
11
scale
reading ________
uncertainty
________
scale
reading ________
uncertainty
________
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F. The Atlantic Pacific Rule For Determining Significant Digits
(1)
(2)
Pacific – "P" is for decimal point is present. If a decimal point is present, count significant digits starting with the
first non-zero digit on the left.
Examples:
(a) 0.004703 has 4 significant digits.
(b) 18.00 also has 4 significant digits.
Atlantic – "A" is for decimal point is absent. If there is no decimal point, start counting significant digits with the
first non-zero digit on the right.
Examples:
(a) 140,000 has 2 significant digits.
(b) 20060 has 4 significant digits.
Imagine a map of the United States. If the decimal is absent count from the Atlantic side. If the decimal point is
present, count from the Pacific side. In both cases, start counting with the first non-zero digit.
Exercise 1.6
How many significant digits are in each of the following numbers?
a)
b)
c)
d)
e)
f)
g)
h)
i)
45.67
4095
30,000,000
45.00043
20
1.0
4
34.8700
0.0034500600
Calculations shouldn't have more precision than the least precise measurement. This leads to 2 Rules for
Calculations:
(A)
(B)
For addition and subtraction: The answer should not have more places past the decimal than the number with the
least places past the decimal.
Example:
1.2 + 12.348 = 13.5
Not 13.548
For multiplication and division: The answer should not have more significant figures than the number with the least
amount of significant figures.
Example:
502 x 3.6 = 1800
Not 1807.2
These last 2 rules can be called the Many-Places rule. For multiplication/division, how many significant figures is
important. For plus/minus, number of places is important.
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I. How many significant digits are in the following numbers?
1.
2.
3.
4.
5.
3.4069
0.56
0.00890
25,000
14.987
II. Write each of the above numbers with one significant digit.
1.
2.
3.
4.
5.
______
______
______
______
______
III. Multiply or divide, according to the problem given below. Make sure that your answer contains the
correct amount of significant digits.
1.
2.
3.
4.
5.
5.67 cm × 6 cm =
0.004090 mm × 12.4 mm =
34 m × 1 m =
30,000 m ÷ 9.008 =
5.67miles ÷ 8.07 =
IV. Add or subtract, according to the problem given below. Make sure that your answer contains the
correct amount of significant digits.
1.
2.
3.
4.
5.
4.56 cm + 6.0 cm =
0.00089 m + 3.4 m =
4.5 m – 0.897 m =
25,000 m/s – 349.00 m/s =
Tricky! 67.8 °C -0 °C =
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H. Factor-Label Method of Dimensional Analysis
Name
Symbol
Size
n
1 × 10-9 _ = 1 n_
micro
μ
1 × 10-6_ = 1 μ _
milli
m
1 × 10-3_ = 1 m_
centi
c
1 × 10-2_ = 1 c_
kilo
k
1 × 103_ = 1 k_
Mega
M
1 × 106_ = 1 M_
nano
Start with a number fact, such as 4.1 cm or 0.075 mL. Examine the units of the desired answer. Multiply
your fact with the factor
what you want
. The starting units cancel out and you end up with the
what you have
desired units. Some conversions require more than one factor. For example, we do not typically convert
directly from kg to μg. So, the best approach is to convert from kg to g (the base unit) then from g to μg.
 though we write factors with x signs, we multiply by the numerators and divide by the
Remember, even
denominators.
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Exercise 1.7
Use the factor-label method to make the following conversions. Remember to use the appropriate
number of significant figures in your answer.
1.
74 cm x
=
meters
2.
8.32 x 10-2 kg x
=
3.
55.5 mL x
4.
0.00527 cal x
5.
9.52 x 10-4 m x
6.
41.0 mL x
=
liters
7.
6.0 x 10-1 g x
=
mg
8.
8.34 x 10-9 cg x
=
g
9.
5.0 x 103 mm x
=
m
10.
1 day x
x
11.
5 x 104 mm x
x
12.
9.1 x 10-13 kg x
x
13.
1 year x
x
14.
4.22 cL x
x
=
mL
15.
1 mile x
x
=
inches
grams
cm3
=
=
kilocalories
=
micrometers
x
=
=
seconds
km
=
=
ng
hours (approximately)
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Exercise 1.8
Use the factor-label method to make the following conversions. Remember to use the appropriate
number of significant figures in your answer.
1.
How many nickels could you trade for 250 yen? $1 = 150 yen.
2.
Your school club sold 600 tickets to a chili supper. The chili recipe for 10 persons requires 2
teaspoons of chili powder? How many teaspoons of chili powder will you need altogether?
3.
How many cups of chili powder will you need? Three teaspoons (tsp) equal one tablespoon
(TBS) and 16 tablespoons equal 1 cup.
4.
How many seconds in a year? (Assume 365 day in a year.)
5.
Chloroform is a liquid once used for anesthetic purposes. What is the volume of 5.0 g of
chloroform? The density of chloroform 1.49 g/mL.
6.
How many inches long is a football field? (A football field is 100 yards.)
7.
How many m3 is 4.6 cm3? Express your answer in scientific notation.
8.
How many mg is 59.0 kg? Express your answer in scientific notation.