CHM 013 Chapter 1-1
CHM 013
Chapter 1 INTRODUCTION
Chemistry - study of properties of materials and changes that they undergo.
We study chemistry because:
1. It provides important understanding of our world and how it works. It has considerable
impact on society (health care, food, clothing, conservation of natural resources,
environmental issues, etc.).
2. It is an essential part of your curriculum. It serves biology, engineering, agriculture, geology,
physics, etc. Chemistry is the central science.
Major Areas of Chemistry:
1. Organic chemistry - study of matter that is composed principally of carbon and hydrogen.
2. Inorganic chemistry - study of matter that consists of all the elements other than carbon
and hydrogen and their combinations.
3. Analytical chemistry - involves the analysis of matter to determine its composition and the
quantity of each kind of matter that is present.
4. Physical chemistry - study of the physical principles that underlie the structure of matter,
energy changes, and chemical transformations.
5. Biochemistry - study of life in the molecular level and the processes associated with life,
such as reproduction, growth and respiration.
Scientific method - a systematic approach to the discovery of new information. It is an
organized approach to solving scientific problems.
1. Observation.
2. Look for patterns, try to explain them and develop a hypothesis - a tentative explanation of
certain facts that provides a basis for further experimentations.
3. Test hypothesis through experimentation (collect data and analyze), refine it.
4. Bring hypotheses together into a theory - a well established hypothesis that explains general
principles of certain phenomena with considerable evidences or facts to support it.
5. Bring all information together into a scientific law - a concise statement or equation that
summarizes tested hypotheses.
A law addresses how matter behaves while a theory addresses why it behaves that way.
Figure 1.1 The scientific method
Classification and Physical States of Matter
Matter - has physical material of the universe, has mass, and occupies space. It is classified by
state (solid, liquid or gas) or by composition (element, compound or mixture).
States of matter - solids, liquids and gases.
Gas - limited to those substances that exists in the gaseous state at room temperature.
Vapor - gaseous state of a substance that is generally a solid or liquid at room temperature.
Properties described on the macroscopic level:
1. Gas - no fixed volume or shape, conforms to shape of container, compressible.
2. Liquid - volume independent of container, no fixed shape, incompressible.
3. Solid - volume and shape independent of container, rigid, incompressible.
Properties described on the molecular level:
1. Gas - molecules far apart, move at high speeds, collide often.
2. Liquid - molecules closer than gas, move rapidly but can slide over each other.
3. Solid - molecules packed closely in definite arrangements.
Composition of Matter
1. Pure substances - matter with fixed composition and distinct proportions.
a. Elements - cannot be decomposed into simpler substances and cannot be broken down
by chemical reaction, made up of only one kind of atom.
b. Compounds - consist of two or more elements and can be broken down by a chemical
reaction. They have different properties than their component elements. For example,
water is liquid which can be broken down to hydrogen and oxygen which are both gases
at the same temperature and pressure.
Law of Constant Proportions - a compound always consists of the same number and
combination of elements (e.g. water is always 11% H and 89% O).
2.
Mixtures - combination of two or more pure substances where each substance retains its
own identity.
a. Heterogeneous - do not have uniform composition, properties and appearance (e.g.
sand and alcohol).
b. Homogeneous - uniform throughout (e.g. coffee in water).
CHM 013 Chapter 1-2
Figure 1.2 Classification scheme of matter
Physical and Chemical Properties of Matter
Each substance has a unique set of physical and chemical properties.
1. Physical properties – quality or condition of a substance that can be observed or measured
without changing its composition (e.g. color, density, odor, melting point, malleability,
ductility, etc.).
2. Chemical properties - describe how substances react or change to form different
substances (e.g. resistance to corrosion, flammability, etc.).
Properties may also be categorized as:
1. Intensive properties - do not depend on amount of substance present and can be used to
establish the identity of the substance (e.g. temperature, melting point, density, etc.).
2. Extensive properties – properties that depend on quantity of substance present (e.g. mass,
volume, etc.). USEFUL MNEMONIC: EXAM - EXtensive property depend on the AMount of
substance.
Intensive properties give an idea of the composition/identity of a substance whereas extensive
properties given an indication of the quantity of substance present.
Physical and Chemical Changes
1. Physical change - substance changes physical appearance without altering its identity (e.g.
change in shape or color, changes of state).
Figure 1.3 Changes of state
2.
Chemical change - substances transform into chemically different substances (e.g. burning
of wood, rusting of iron nail, growth of plants).
Separation of Mixtures
Separation of Heteregeneous Mixtures:
1. Filtration - removal of a solid from liquid hetereogeneous mixture through a medium from
which only the liquid can pass.
2. Decantation - separation of liquid from heavier solid by carefully pouring the liquid from the
container in order to leave the solid in the bottom of the original container.
Separation of Homogeneous Mixtures:
1. Distillation - a process exploiting the differences between the boiling points of the
components of the mixture. The component that boils off first is the distillate while the one
that remains in the original container is the residue.
2. Solvent extraction (liquid-liquid extraction) - is a method to separate compounds based on
their relative solubilities in two different immiscible liquids usually water and an organic
solvent.
3. Adsorption - a process where a solid is used to remove a soluble substance from an
aqueous solution.
Measurement: Units and Handling Numbers
CHM 013 Chapter 1-3
Units of Measurement
A measured quantity must have both a number and a unit. The units most often used for
scientific measurement are those of the metric system.
Scientific Notation
Some numbers are very large or very small and are thus difficult to express such as:
Avogadro’s number = 602 000 000 000 000 000 000 000 particles/mol
An electron’s mass = 0.000 000 000 000 000 000 000 000 000 91 kg
Also, it's not clear how many significant figures (SF) there are in some measurements. To handle
such numbers, we use a system called scientific notation wherein numbers can be expressed in
the form N x 10n where N (digit term) - a number from 1 to 9 such that there can only be one
number to the left of the decimal point (e.g. 6.02XX) and n (exponent) - a positive or a negative
integer (whole number).
To express a number in scientific notation, count the number of places you must move the
decimal point to get N.
 Move the decimal point to the right if the number is less than 1 = negative exponent.
 Move the decimal point to the left if the number is greater than 1 = positive exponent.
EXERCISE 1.1
Express the following numbers in scientific notation:
1. 555,000
______________________
2. 0.000888
______________________
3. 602,000,000,000,000,000,000,000
______________________
SI Units
All scientific units use Système International d’Unités (SI units) as of 1960 from which smaller and
larger units are obtained by decimal fractions or multiples of the base units.
Table 1.1 SI base units
Physical Quantity
Mass
Length
Time
Temperature
Amount of substance
Electric current
Luminous intensity
Name of Unit
Kilogram
Meter
Second
Kelvin
Mole
Ampere
Candela
Abbreviation
kg
m
s
K
mol
A
cd
SOME UNIT CONVERSIONS
Length
2.54 cm = 1 in
12 in = 1 ft
Mass
2.2 lbs = 1 kg
Temperature
K = °C + 273
°C =
Volume
1 L = 1000 cm3
5
(°F - 32)
9
1 mL = 1 dm3
9
°F = 5 °C + 32
Table 1.2 Selected prefixes used in metric system
Prefix
Abbreviation
Meaning
GigaG
109
MegaM
106
KiloK
103
Decid
10-1
Centic
10-2
Millim
10-3
Microμ
10-6
Nanon
10-9
Picop
10-12
Femtof
10-15
EXERCISE 1.2
Complete the following unit equations:
1. 1 kg = ______ g
2. 1 cm = ______ m
3.
4.
1 s = ______ ns
1 L = ______ ML
Uncertainty in Measurement
Two types of numbers:
1. Exact numbers - obtained from counting or defined as conversion factor.
CHM 013 Chapter 1-4
2.
Inexact numbers - derived from measurement. All measurements have some degree of
uncertainty or error associated with the instrument used and is thus inexact.
Precision and Accuracy
Precision - how well measured quantities agree with each other.
Accuracy - how well measured quantities agree with the “true value.”
Rounding off non-significant digits
How do we eliminate nonsignificant digits?
1. If first nonsignificant digit < 5, just drop ALL nonsignificant digits.
2. If first nonsignificant digit ≥ 5, raise the last significant digit by 1 then drop ALL nonsignificant
digits.
EXERCISE 1.3
Express each of the following with the number of SF indicated. Express measurements in scientific
notation when necessary to make it clear how many significant figures there are in the
measurement.
a. 376.276 to 3 SF
_______________________
b. 500.072 to 4 SF
_______________________
c. 0.00654321 to 3 SF
_______________________
d. 1,234,567 to 5 SF
_______________________
e. 2975 to 2 SF
_______________________
Figure 1.4 Illustration on accuracy and precision
Significant Figures in Calculations
1. Multiplication and Division - report to the least number of significant figures.
e.g. 6.221 cm x 5.2 cm = 32 cm2
2.
Significant Figures
In a measurement, it is useful to indicate the exactness of the measurement. This exactness is
reflected in the number of significant figures. The number of significant figures is the number
of digits known with certainty plus one uncertain digit. For example, 2.2405 g means we are sure
the mass is 2.240 g but we are uncertain about the nearest 0.0001 g. Final calculations are only
as significant as the least significant measurement.
Rules on determining the number of significant figures:
1. Nonzero numbers and zeros between nonzero numbers are always significant.
2. Zeros before the first nonzero digit are not significant (e.g. 0.0003 has one SF).
3. Zeros at the end of the number after a decimal point are significant (e.g. 1.00 has three SF).
4. Zeros at the end of a number before a decimal point are ambiguous. Exponential notation
eliminates this ambiguity (e.g. 10 300 - number of SF cannot be determined; suppose you
wish to express this number in three SF, you write 1.03 x 104).
Method:
1. Write the number in scientific notation.
2. The number of digits remaining is the number of significant figures.
2.50 x 102 cm has three SF
1.03 x 104 g has three SF
4
1.030 x 10 lb has four SF
1.0300 x 104 in has five SF
Addition and Subtraction - report to the least number of decimal places.
e.g. 20.4 g – 1.322 g = 19.1 g
EXERCISE 1.4
a. 106.61 + 0.25 + 0.195 = 107.055
107.055 to the correct number of SF
b. 725.50 – 103 = 622.50
622.50 to the correct number of SF
____________________
____________________
In multiple step calculations always retain an extra significant figure until the end to prevent
rounding errors.
Multiplying or dividing with exponential numbers
When multiplying or dividing measurements with exponents, use the digit term (N in N × 10n) to
determine number of significant figures.
(6.02 × 1023)(4.155 × 109) = 2.50131× 1033
How do you calculate this using your scientific calculator?
1. Enter “6.02 × 1023” by pressing 6.02 then EE or EXP (which corresponds to “× 10”) then 23.
CHM 013 Chapter 1-5
2.
3.
4.
Your calculator should read or something similar.
Multiply by pressing ×.
Enter “4.155 × 109” by pressing 4.155 then EE or EXP (which corresponds to “× 10”) then
9.
Get the answer by pressing =.
Your calculator should now read:
Or something similar indicating:
Thus, the answer with the correct SF
2.50131 33 or 2.50131 E 33
2.50131 × 10 33
____________________
Be sure you can do exponential calculations with your calculator. Many of the calculations we do
in chemistry involve very large and very small numbers with exponential terms.
EXERCISE 1.5
1. (3.75 × 1015)(8.6 × 104) = 3.225 × 1020
with the correct SF
2.
Equivalents are exact if we can count the number of units equal to another or the units are in the
same system (metric or English). Exact equivalents have an infinite number of SF - never limit the
number of significant figures. When the relationship between two units or items is exact, the “≡”
(meaning “equals exactly”) is used instead of the basic “=” sign.
Approximate equivalents do limit the significant figures for the final answer.
Using Two or More Conversion Factors
We often need to use more than one conversion factor in order to complete a problem. When
identical units are found in the numerator and denominator of a conversion, they will cancel. The
final answer MUST have the correct units.
____________________
Suppose that we want to convert length in meters to length in inches. We could do this
conversion with the following conversion factors 1 m = 100 cm and 1 in = 2.54 cm.
(3.75 × 1015) ÷ (8.605 × 104) = 4.357931435 × 1010
with the correct SF
___________________
The calculation would involve both conversion factors and the units of the final answer will be
inches:
Dimensional Analysis/Factor Label Method
Dimensional analysis - a method of calculation utilizing the knowledge of units. Given units can
be multiplied and divided to give the desired units and conversion factors are used to manipulate
units:
Desired unit = Given unit x Conversion factor
The conversion factors are simple ratios
Conversion factor 
Desired unit
Given unit
Using dimensional analysis in problem solving:
1. Write the units for the answer.
2. Determine what information to start with.
3. Arrange all other unit factors, showing them as fractions, with correct units in the numerator
and denominator so all units cancel except for the units needed for the final answer.
4. Check for correct units and number of significant figures in the final answer.
EXERCISE 1.6
1. If a marathon is 26.2 miles, how many inches are in the marathon? (1 mi ≡ 5280 ft)
2. The speed of light is about 3.00 x 108 m/s. Express this speed in mi/hr. (1.609 km = 1 mi,
1000 m ≡ 1km)
Unit equations and unit factors
Unit equation - a simple statement of two equivalent values.
Unit factor - a ratio of two equivalent quantities.
Unit equation
 100cm  1in 
( givenmeters)

  ____ in
 1m  2.54cm 
Unit factor
2.54 cm = 1 in
2.54cm
1in
or
1in
2.54cm
1 dozen = 12 eggs
1doz
12eggs
or
12eggs
1doz
Conversions Involving Volume
We often will encounter conversions from one measure to a different measure.
For example, suppose that we wish to know the mass in grams of 2.00 in 3 of gold given that the
density of the gold is 19.3 g/cm3.
We could do this conversion with the following conversion factors: 2.54 cm = 1 in and 1 cm3 =
CHM 013 Chapter 1-6
19.3 g gold. The calculation would involve both of these factors:
b.
Gasoline
d.
Carbon dioxide
3
 2.54cm   19.3g 
(2.00in 3 )
 ______ g
 
3 
 1in   1cm 
5.
Which quantity of each pair is larger? Show your solution.
a. 1 nm or 1 m
c.
1 pg or 1 cg
b. 1 Mg or 1 mg
6.
How many significant figures does each number possess? Write the numbers using scientific
notation.
a. 225
d.
1.0025
b. 14,100.0
e.
10,004
c. 0.002500
f.
0.0025
Round the answers in the following problems to the correct number of significant figures.
a. 12.25 + 1.32 + 1.2
d.
(1.256)(2.42)
b. 13.7325 – 14.21
e.
(16.231)/2.20750
c. 12300 + 2.11
f.
[(1.1)(2.62)(13.5278)]/2.650
Note that the calculation will NOT be correct unless the centimeter to inch conversion factor is
cubed! Both the units AND the number must be cubed.
ADDITIONAL EXERCISES
1. Match the following characteristics to one or more of the three states of matter:
a. Has no shape of its own
b. Definite shape
c. Occupies the total volume of a container
d. Partially takes on the shape of a container
e. Does not take on the shape of a container
f. Readily compressible
g. Slightly compressible
h. Essentially incompressible
2.
Match the term with the best identifying phrase.
Terms:
Homogeneous mixture
Heterogeneous mixture
Mixture
Substance
Element
Compound
Phrases:
a. Any kind of matter that is pure and has a fixed composition
b. Cannot be decomposed into simpler substances by chemical changes
c. A solution of uniform composition
d. Can be decomposed into simpler substances by chemical changes
e. Any kind of matter that can be separated into simpler substances by physical means
f. Non-uniform composition
3.
Are
a.
b.
c.
d.
the following changes physical or chemical?
The vaporization of solid carbon dioxide
Explosion of solid TNT
The aging of an egg with a resultant unpleasant smell
The formation of a solid when honey is cooled
4.
Classify each of the following as an element, compound, or mixture.
a. A 100% silver bar
c.
Wine
7.
8.
The temperature on a spring day is around 22 °C. What is this temperature in degrees
Fahrenheit and Kelvin?
9.
Which has the greater mass, 2.0 cm3 of iron (density = 7.9 g/cm3) or 1.0 cm3 of gold (density
= 19.32 g/cm3)?
10. Make the following conversions:
a. 1 ML to cm3
b. 10.5 L to mL
c.
d.
0.0001 g to kg
50 kg to lbs