To study of the chemistry

1. 1-1
Keys to the Study of Chemistry
Chapter 1

2. 1-2
Chapter 1 : Keys to the Study of Chemistry
1.1 Some Fundamental Definitions
1.2 Chemical Arts and the Origins of Modern Chemistry
1.3 The Scientific Approach: Developing a Model
1.4 Chemical Problem Solving
1.5 Measurement in Scientific Study
1.6 Uncertainty in Measurement: Significant Figures

3. 1-3
CHEMISTRY
is the study of matter,
its properties,
the changes that matter undergoes,
and
the energy associated with these changes.

4. 1-4
Definitions
Chemical Properties
Those which the substance shows
as it interacts with, or transforms
into, other substances (such as
flammability, corrosiveness)
Matter anything that has mass and volume -the “stuff” of the
universe: books, planets, trees, professors, students
Composition the types and amounts of simpler substances that
make up a sample of matter
Properties the characteristics that give each substance a unique
identity
Physical Properties
Those which the substance
shows by itself without
interacting with another
substance (such as color,
melting point, boiling point,
density)

5. 1-5

6. 1-6

7. 1-7
Sample Problem 1.1 Distinguishing Between Physical and
Chemical Change
PROBLEM: Decide whether each of the following processes is primarily a
physical or a chemical change, and explain briefly.
PLAN: Does the substance change composition or just change form?
SOLUTION:
(a) Frost forms as the temperature drops on a humid winter night.
(b) A cornstalk grows from a seed that is watered and fertilized.
(c) Dynamite explodes to form a mixture of gases.
(d) Perspiration evaporates when you relax after jogging.
(e) A silver fork tarnishes in air.
(a) physical change (b) chemical change (c) chemical change
(d) physical change (e) chemical change

8. 1-8
Figure 1.2
The Physical States of Matter

9. 1-9
Figure 1.3
Energy is the capacity to do work.
energy due to the position of the
object or energy from a chemical
reaction
Potential Energy
Kinetic Energy
energy due to the motion of the
object
Potential and kinetic energy
can be interconverted.

10. 1-10
Figure 1.3(continued) Energy is the capacity to do work.
Potential Energy
energy due to the position of the
object or energy from a chemical
reaction
Kinetic Energy
energy due to the motion of the
object
Potential and kinetic energy
can be interconverted.

11. 1-11
Scientific Approach: Developing a Model
Observations :
Hypothesis:
Experiment:
Model (Theory):
Further Experiment:
Natural phenomena and measured events;
universally consistent ones can be stated as a
natural law
Tentative proposal that explains observations
Procedure to test hypothesis; measures one
variable at a time
Set of conceptual assumptions that explains
data from accumulated experiments; predicts
related phenomena
Tests predictions based on model
revised if
experiments do
not support it
altered if
predictions do
not support it

12. 1-12
Sample Problem 1.2 Converting Units of Length
PROBLEM: What is the price of a piece of copper wire 325 centimeters (cm)
long that sells for $0.15/ft?
PLAN: Know length (in cm) of wire and cost per length (in ft)
Need to convert cm to inches and inches to ft followed by finding
the cost for the length in ft.
SOLUTION:
length (cm) of wire
length (ft) of wire
length (in) of wire
Price ($) of wire
2.54 cm = 1 in
12 in = 1 ft
1 ft = $0.15
Length (in) = length (cm) x conversion factor
= 325 cm x in
2.54 cm
= 128 in
Length (ft) = length (in) x conversion factor
= 128 in x ft
12 in
= 10.7 ft
Price ($) = length (ft) x conversion factor
= 10.7 ft x $0.15
ft
= $1.60

13. 1-13
Table 1. 2 SI - Base Units
Physical Quantity Unit Name Abbreviation
mass
meter
kg
length
kilogram
m
time second s
temperature kelvin K
electric current ampere A
amount of substance mole mol
luminous intensity candela cd

14. 1-14
Common Decimal Prefixes Used with SI Units
Prefix Prefix
Symbol
Number Word Exponential
Notation
tera T 1,000,000,000,000 trillion 1012
giga G 1,000,000,000 billion 109
mega M 1,000,000 million 106
kilo k 1,000 thousand 103
hecto h 100 hundred 102
deka da 10 ten 101
----- ---- 1 one 100
deci d 0.1 tenth 10-1
centi c 0.01 hundredth 10-2
milli m 0.001 thousandth 10-3
micro  0.000001 millionth 10-6
nano n 0.000000001 billionth 10-9
pico p 0.000000000001 trillionth 10-12
femto f 0.000000000000001 quadrillionth 10-15
Table 1.3

15. 1-15
Common SI-English Equivalent Quantities
English Equivalent
Mass
Length 1 kilometer (km) 1000 (103) m 0.6214 mi
1 meter (m) 100 (102) cm 1.094 yd
1000 (103) mm 39.37 in
1 centimeter (cm) 0.01 (10-2 ) m 0.3937 in
Volume 1,000,000 (106) cm3 35.31 ft3
1 cubic meter (m3)
1 cubic decimeter
(dm3)
1000 cm3 0.2642 gal
1.057 qt
1 cubic
centimeter (cm3)
0.001 dm3 0.03381 fluid ounce
1 kilogram (kg) 1000 g 2,205 lb
1 gram (g) 1000 mg 0.03527 oz
Table 1.4
Quantity SI Unit SI Equivalent
English to
SI Equivalent
1 mi = 1.609 km
1 yd = 0.9144 m
1 ft = 0.3048 m
1 in = 2.54 cm
(exactly)
1 ft3 = 0.02832 m3
1 gal = 3.785 dm3
1 qt = 0.9464 dm3
1 qt = 946.4 cm3
1 fluid ounce = 29.57 cm3
1 lb = 0.4536 kg
1 lb = 453.6 g
1 oz = 28.35 g

16. 1-16

17. 1-17
Sample Problem 1.3 Determining the Volume of a Solid by
Displacement of Water
PROBLEM: The volume of an irregularly shaped solid can be determined
from the volume of water it displaces. A graduated cylinder
contains 19.9 mL water. When a small piece of galena, an ore
of lead, is submerged in the water, the volume increases to
24.5 mL. What is the volume of the piece of galena in cm3
and in L?
PLAN: The volume of galena is equal to the change in the water
volume before and after submerging the solid.
volume (mL) before and after addition
volume (mL) of galena
volume (cm3) of
galena
subtract
volume (L) of
galena
1 mL = 1 cm3 1 mL = 10-3 L
SOLUTION:
(24.5 - 19.9)mL = volume of galena
4.6 mL x
mL
1 cm3
= 4.6 cm3
4.6 mL x
mL
10-3 L = 4.6x10-3 L

18. 1-18
Sample Problem 1.4 Converting Units of Mass
PROBLEM: International computer communications will soon be carried by
optical fibers in cables laid along the ocean floor. If one strand of
optical fiber weighs 1.19 x 10 -3 lbs/m, what is the total mass (in kg)
of a cable made of six strands of optical fiber, each long enough to
link New York and Paris (8.84 x 103km)?
PLAN: The sequence of steps may vary but essentially you have to find
the length of the entire cable and convert it to mass.
length (km) of fiber
mass (lb) of fiber
length (m) of fiber
1 km = 103 m
mass (kg) of cable
mass (lb) of cable
6 fibers = 1 cable
1 m = 1.19x10-3 lb
2.205 lb = 1 kg
SOLUTION:
8.84 x 103 km x
km
103 m
8.84 x 106 m x
m
1.19 x 10 -3lbs
1.05 x 104 lb x
cable
6 fibers
= 1.05 x 104 lb
= 8.84 x 106 m
2.205 lb
1 kg
x
6.30x 104 lb
cable
6.30x 104 lb
=
cable
2.86x104 kg
cable
=

19. 1-19
Some Interesting Quantities
Length Volume Mass
Figure 1. 10

20. 1-20
Substance Physical State Density (g/cm3)
Densities of Some Common Substances
Table 1.5
Hydrogen Gas 0.000089
Oxygen Gas 0.0014
Grain alcohol Liquid 0. 789
Water Liquid 1.000
Table salt Solid 2.16
Aluminum Solid 2.70
Lead Solid 11.3
Gold Solid 19.3

21. 1-21
Sample Problem 1.5 Calculating Density from Mass and Length
PROBLEM: Lithium (Li) is a soft, gray solid that has the lowest density
of any metal. If a slab of Li weighs 1.49 x 103 mg and has
sides that measure 20.9 mm by 11.1 mm by 12.0 mm, what
is the density of Li in g/cm3 ?
PLAN: Density is expressed in g/cm3 so we need the mass in grams
and the volume in cm3.
mass (mg) of Li
lengths (mm) of sides
mass (g) of Li
density (g/cm3) of Li
103mg = 1g
10 mm = 1 cm
lengths (cm) of sides
volume (cm3)
multiply lengths
SOLUTION:
20.9 mm x
mg
10-3g = 1.49 g
1.49x103 mg x
10 mm
cm
= 2.09 cm
Similarly the other sides will be 1.11
cm and 1.20 cm, respectively.
2.09 x 1.11 x 1.20 = 2.76 cm3
density of Li =
1.49 g
2.76cm3
= 0.540 g/cm3

22. 1-22
Figure 1.11
Some
Interesting
Temperatures

23. 1-23
Figure 1.12
The Freezing and Boiling Points of Water

24. 1-24
Temperature Scales and Interconversions
Kelvin ( K ) - The “absolute temperature scale” begins at
absolute zero and only has positive values.
Celsius ( o
C ) - The temperature scale used by science,
formally called centigrade and most commonly used scale around
the world. Water freezes at 0o
C and boils at 100o
C.
Fahrenheit ( o
F ) - Commonly used scale in the US for weather
reports. Water freezes at 32o
F and boils at 212o
F.
T (in K) = T (in oC) + 273.15
T (in oC) = T (in K) - 273.15
T (in oF) = 9/5 T (in oC) + 32
T (in oC) = [ T (in oF) - 32 ] 5/9

25. 1-25
Sample Problem 1.6 Converting Units of Temperature
PROBLEM: A child has a body temperature of 38.70C.
PLAN: We have to convert 0C to 0F to find out if the child has a fever
and we use the 0C to kelvin relationship to find the temperature
in kelvins.
(a) If normal body temperature is 98.60F, does the child have a fever?
(b) What is the child’s temperature in kelvins?
SOLUTION:
(a) Converting from 0C to 0F
9
5
(38.70C) + 32 = 101.70F
(b) Converting from 0C to K 38.70C + 273.15 = 311.8K

26. 1-26
The Number of Significant Figures in a Measurement
Depends Upon the Measuring Device
Figure 1.14

27. 1-27

28. 1-28
Rules for Determining Which Digits are Significant
All digits are significant
•Make sure that the measured quantity has a decimal point.
•Start at the left of the number and move right until you reach
the first nonzero digit.
•Count that digit and every digit to its right as significant.
Zeros that end a number and lie either after or before the
decimal point are significant; thus 1.030 mL has four
significant figures, and 5300. L has four significant figures
also. Numbers such as 5300 L are assumed to only have 2
significant figures. A terminal decimal point is often used to
clarify the situation, but scientific notation is the best!
except zeros that are used only to
position the decimal point.

29. 1-29
Sample Problem 1.7 Determining the Number of Significant Figures
PROBLEM: For each of the following quantities, underline the zeros that are
significant figures(sf), and determine the number of significant
figures in each quantity. For (d) to (f) express each in
exponential notation first.
PLAN: Determine the number of sf by counting digits and paying attention
to the placement of zeros.
SOLUTION:
(b) 0.1044 g
(a) 0.0030 L (c) 53.069 mL
(e) 57,600. s
(d) 0.00004715 m (f) 0.0000007160 cm3
(b) 0.1044 g
(a) 0.0030 L (c) 53.069 mL
(e) 57,600. s
(d) 0.00004715 m (f) 0.0000007160 cm3
2sf 4sf 5sf
(d) 4.715x10-5 m 4sf (e) 5.7600x104 s 5sf (f) 7.160x10-7 cm3 4sf

30. 1-30
Rules for Significant Figures in Answers
1. For addition and subtraction. The answer has the
same number of decimal places as there are in the
measurement with the fewest decimal places.
106.78 mL = 106.8 mL
Example: subtracting two volumes
863.0879 mL = 863.1 mL
865.9 mL
- 2.8121 mL
Example: adding two volumes 83.5 mL
+ 23.28 mL

31. 1-31
2. For multiplication and division. The number with the least
certainty limits the certainty of the result. Therefore, the answer
contains the same number of significant figures as there are in the
measurement with the fewest significant figures.
Rules for Significant Figures in Answers
Multiply the following numbers:
9.2 cm x 6.8 cm x 0.3744 cm = 23.4225 cm3 = 23 cm3

32. 1-32
Issues Concerning Significant Figures
graduated cylinder < buret ≤ pipet
numbers with no uncertainty
1000 mg = 1 g
60 min = 1 hr
These have as many significant digits as the calculation requires.
be sure to correlate with the problem
FIX function on some calculators
Electronic Calculators
Choice of Measuring Device
Exact Numbers

33. 1-33
Rules for Rounding Off Numbers
1. If the digit removed is more than 5, the preceding number
increases by 1.
5.379 rounds to 5.38 if three significant figures are retained
and to 5.4 if two significant figures are retained.
2. If the digit removed is less than 5, the preceding number is
unchanged.
0.2413 rounds to 0.241 if three significant figures are retained
and to 0.24 if two significant figures are retained.
3.If the digit removed is 5, the preceding number increases by
1 if it is odd and remains unchanged if it is even.
17.75 rounds to 17.8, but 17.65 rounds to 17.6.
If the 5 is followed only by zeros, rule 3 is followed; if the 5 is
followed by nonzeros, rule 1 is followed:
17.6500 rounds to 17.6, but 17.6513 rounds to 17.7
4. Be sure to carry two or more additional significant figures
through a multistep calculation and round off only the final
answer.

34. 1-34
Sample Problem 1.8 Significant Figures and Rounding
PROBLEM: Perform the following calculations and round the answer to the
correct number of significant figures.
PLAN: In (a) we subtract before we divide; for (b) we are using an exact
number.
SOLUTION:
7.085 cm2
16.3521 cm2 - 1.448 cm2
(a)
11.55 cm3
4.80x104 mg
(b)
1 g
1000 mg
7.085 cm2
16.3521 cm2 - 1.448 cm2
(a) =
7.085 cm2
14.904 cm2
= 2.104 cm2
11.55 cm3
4.80x104 mg
(b)
1 g
1000 mg
=
48.0 g
11.55 cm3
= 4.16 g/ cm3

35. 1-35
Precision and Accuracy
Errors in Scientific Measurements
Random Error
In the absence of systematic error, some values that are higher and
some that are lower than the actual value
Precision
Refers to reproducibility or how close the measurements are to each
other
Accuracy
Refers to how close a measurement is to the actual value
Systematic error
Values that are either all higher or all lower than the actual value

36. 1-36
Figure 1.16
precise and accurate
precise but not accurate
Precision and Accuracy in the Laboratory

37. 1-37
systematic error
random error
Precision and Accuracy in the Laboratory
Figure 1.16 continued