Chapter 1
Introduction: Some Basic Concepts
Welcome to the World of
Chemistry
Chemistry
The study of matter – its nature, its structure
(how it is related to its atoms and
molecules), properties, transformations,
and its interactions with energy
Gold
Mercury
Matter
Anything that has mass
and occupies space
Mass vs Weight
weight = force = mg
g: gravitational acceleration
 Mass is a measurement of the
quantity of matter in a body or
sample
 Weight is the magnitude of Earth’s
attraction to such a body or sample

Physical States (Phases)
Example of bromine, Br2, a halogen
Solid
definite shape and volume
 made of particles (atoms, molecules, or
ions) held close together and rigidly in
place
 reasonably well understood.

Example:
Graphite — layer
structure of carbon
atoms reflects
physical properties.
Liquid
definite volume but
indefinite shape
 made of particles
(atoms, molecules, or
ions) held close together
but allowed to move
relative to each other
 fluid and may not fill a
container completely
 not well understood

Gas





indefinite volume and
indefinite shape
the same shape and
volume as their container
made of particles (atoms or
molecules) separated from
each other by large
distances and that move
very fast
fluid
good theoretical
understanding
Physical Property



characteristic of matter
that can be observed
without changing the
basic identity of the
matter
characteristics that are
directly observable
eg. state, size, mass,
V, color, odor, melting
point (Tm), boiling point
(Tb), density,
solubility...
Chemical Property
characteristic of matter that requires
change in identity of the matter for
observation (a chemical reaction)
 characteristic that describes the
behavior of matter


eg. flammability, corrosiveness,
bleaching power, explosiveness, ...
Scientific Method
Observation
Hypothesis
*
Law
*
*
Theory
* experiment and then modify
Scientific Method
Procedure
designed to
test an idea
Tentative explanation of a
single or small number of
observations
General explanation of
natural phenomena
Careful noting and
recording of natural
phenomena
Tro: Chemistry: A Molecular Approach,
Approach 2/e
Generally observed
occurrence in nature
12
Relationships Between Pieces of
the Scientific Method
Tro: Chemistry: A Molecular Approach,
Approach 2/e
13



Hypothesis – a tentative interpretation or
explanation for an observation
– falsifiable – confirmed or refuted by other
observations
– tested by experiments – validated or
invalidated
when similar observations are consistently
made, it can lead to a Scientific Law
– a statement of a behavior that is always
observed
– summarizes past observations and
predicts future ones
– Law of Conservation of Mass
A theory is a unifying principle that explains
a body of facts and the laws based on them.
It is capable of suggesting new hypotheses.
Classification of Matter
Mixture
A combination of pure substances in
which the components retain their
identities (no reaction)
 Can be separated into simpler
mixtures and/or pure substances by

Physical Separation Methods





mechanical: eg. sand and iron filings
filtration: eg. sand and water
extraction: eg. washing clothes, decaffeinating
coffee
distillation
chromatography
Distillation

Simple - for separation of volatile component
from non-volatile component(s)
Distillation

Fractional - for separation of multiple
volatile components from each other.
Employed in many chemistry labs,
labs, and in crude oil refining.
Chromatography

Mixture placed in mobile phase (gas or
liquid). Mobile phase flows over and
through stationary phase (solid or liquid).
Mixture components separate based on
relative affinity for mobile and stationary
phases.
Heterogeneous Mixture
inconsistent composition
 atoms or molecules mixed not uniformly
 contains regions within the sample with
different characteristics
 eg. pizza, carpet, beach sand, ...

Homogeneous Mixture
solution
 consistent composition throughout
 atoms or molecules mixed uniformly
 eg. air in a room, glass of tap water

Compound
can be broken down to 2 or more
elements by chemical means
 constant composition
 eg. water, H2O, by mass H:O = 1:8
hydrogen peroxide, H2O2, H:O = 1:16
 elements combined lose individual
identities
 more than 20 million compounds are
now known

Elements





basic substances of
which all matter is
composed
pure substances that
cannot be decomposed
by ordinary means to
other substances.
made up of atoms
~ 117 known at this time
given name and chemical
symbol
Aluminum
Bromine
Element Symbols
1, 2 or 3 letters:
 first letter always capitalized
 usually first letter(s) of name

H hydrogen
O oxygen
Na sodium
Al aluminum
C carbon
N nitrogen
Cl chlorine Mg magnesium
P phosphorus K potassium
Po polonium
learn Latin names where appropriate,
antimony - Sb - stibium
gold - Au - aurum
tungsten - W – wolfram
sodium – Na – natrium
potassium – K - kalium
 elements from 104 to 111 are named
after scientists; 112-118 have 3 letter
symbols based on Latin name for
number

112
113
114
115
116
Uub
Uut
Uuq
Uup
Uuh
ununbium
ununtrium
ununquadium
ununpentium
ununhexium
Homework: learn the names of
first 36 elements in the periodic table
Periodic Table
a listing of the elements arranged
according to their atomic numbers,
chemical and physical properties
 VERY useful and important

Physical Change
transformation of matter from one
state to another that does not involve
change in the identity of the matter
 examples: boiling, subliming,
melting, dissolving (forming a
solution), ...

Chemical Change


transformation of matter from one state to another
that involves changing the identity of the matter
examples: rusting (of iron), burning (combustion),
digesting, formation of a precipitate, gas forming,
acid-base neutralization, displacing reactions...
Intensive Property
independent of amount of matter
 eg. density, temperature,
concentration of a solution, specific
heat capacity...

Extensive Property
depends on amount of matter
 eg. mass, volume, pressure, internal
energy, enthalpy, ...

Density



mass (g)
mass (g)
Density =  = 
volume (cm3)
volume (mL)
density of H2O is 1.00 g/cm3 (pure water at ~ 4 °C)
1cm3 = 1mL
Mercury
Platinum
Aluminum
liquid
13.6 g/cm3
21.5 g/cm3
They sink in water
2.7 g/cm3
Know and Own and Practice
Well
Metric System
 SI Units
 Unit Conversions
 Learn a Conversion Factor Between
English and Metric for
– length, mass, volume, pressure

SI Units


Système International d’Unités
A different base unit is used for each quantity.
Prefixes
A prefix
 in front of a unit increases or decreases the size
of that unit.
 makes units larger or smaller than the initial unit
by one or more factors of 10.
 indicates a numerical value.
prefix
1 kilometer
1 kilogram
=
=
=
value
1000 meters
1000 grams
Metric and SI Prefixes
Learning Check
Indicate the unit that matches the description.
1. A mass that is 1000 times greater than 1
gram.
1) kilogram 2) milligram 3) megagram
2. A length that is 1/100 of 1 meter.
1) decimeter 2) centimeter
3) millimeter
3. A unit of time that is 1/1000 of a second.
1) nanosecond
2) microsecond
3) millisecond
Learning Check
Select the unit you would use to measure
A. your height.
1) millimeters 2) meters
3) kilometers
B. your mass.
1) milligrams 2) grams
3) kilograms
C. the distance between two cities.
1) millimeters 2) meters
3) kilometers
D. the width of an artery.
1) millimeters 2) meters
3) kilometers
Volume
1 m = 10 dm
(1m)3 = (10 dm)3
1m3 = 1000 dm3 = 1000 L
1 dm = 10 cm
(1dm)3 = (10 cm)3
1dm3 = 1000 cm3 = 1000mL
Equalities
Equalities
• use two different units to describe the same
measured amount.
• are written for relationships between units of the
metric system, U.S. units, or between metric and
U.S. units.
For example,
1 in = 2.54 cm
1 m = 1000 mm
1 ft = 12 in
1 lb = 16 oz
1 mile = 5280 ft
2.205 lb = 1 kg
1 lb = 454 g
1L
= 1.057 qt
1 gal = 4 qt
1 hour = 60 min
1 cm3 = 1 cc = 1 mL
Conversion Factors
A conversion factor
• is a fraction obtained from an equality.
Equality:
1 in. = 2.54 cm
• is written as a ratio with a numerator and
denominator.
• can be inverted to give two conversion factors
for every equality.
1 in.
and 2.54 cm
2.54 cm
1 in.
Learning Check
Write conversion factors for each pair of units.
A. liters and mL
Equality: 1 L = 1000 mL
B. hours and minutes
Equality:
1 hr = 60 min
C. meters and kilometers
Equality:
1 km = 1000 m
D. micrograms and grams
Equality:
1 µg = 10-6 g
Conversion Factors in a Problem
A conversion factor
• may be obtained from information in a word
problem.
• is written for that problem only.
Example 1:
The price of one pound (1 lb) of red peppers is
$2.39.
1 lb red peppers
and
$2.39
$2.39
1 lb red peppers
Example 2:
The cost of one gallon (1 gal) of gas is $3.95.
1 gallon of gas
and
$3.95
$3.95
1 gallon of gas
Percent as a Conversion Factor
A percent factor
• gives the ratio of the parts to the whole.
% =
•
•
•
Parts x 100
Whole
uses the same unit to express the percent.
uses the value 100 and a unit for the whole.
can be written as two factors.
Example: A food contains 30% (by mass) fat.
30 g fat
100 g food
and
100 g food
30 g fat
Density as a conversion factor
Density of a mineral oil = 0.875 g/mL
0.875 g oil
1 mL
and
1 mL
0.875 g oil
Learning Check
Write the equality and conversion factors for
each of the following.
A. square meters and square centimeters
B. jewelry that contains 18% (by mass) gold
C. One gallon of gas is $4.00
Solving: Given and Needed Units
To solve a problem
• Identify the given unit
• Identify the needed unit.
Example:
A person has a height of 2.0 meters.
What is that height in inches?
The given unit is the initial unit of height.
given unit = meters (m)
The needed unit is the unit for the
answer.
needed unit = inches (in.)
Problem Setup: Dimensional Analysis
• Write the given and needed units.
• Write a unit plan to convert the given unit
to the needed unit.
• Write equalities and conversion factors that
connect the units.
• Use conversion factors to cancel the given
unit and provide the needed unit.
Unit 1
x
Given
unit
x
Unit 2
= Unit 2
Unit 1
Conversion = Needed
factor
unit
Setting up a Problem
How many minutes are 2.5 hours?
Given unit
=
2.5 hr
Needed unit =
min
Unit Plan =
hr → min
Setup problem to cancel hours (hr).
Given Conversion
Needed
unit
factor
unit
2.5 hr x 60 min
= 150 min (2 SF)
1 hr
Learning Check
A rattlesnake is 2.44 m long. How many
centimeters long is the snake?
1) 2440 cm
2) 0.0244 cm
3) 24.4 cm
4) 244 cm
Using Two or More Factors
• Often, two or more conversion factors are
required to obtain the unit needed for the answer.
Unit 1
→
Unit 2
→
Unit 3
• Additional conversion factors are placed in the
setup to cancel each preceding unit
Given unit x factor 1 x factor 2
Unit 1 x
Unit 2
Unit 1
x
Unit 3
Unit 2
= needed unit
= Unit 3
Example: Problem Solving
How many minutes are in 1.4 days?
Given unit: 1.4 days
Factor 1
Plan:
days
→
Factor 2
hr
→
min
Set up problem:
1.4 days x 24 hr x 60 min = 2.0 x 103 min
1day
1 hr
2 SF
Exact
Exact
= 2 SF
Learning Check
A bucket contains 4.65 L of water. How many
gallons of water is that?
Unit plan: L
→
qt
→ gallon
Equalities: 1.06 qt = 1 L
Set up Problem:
4.65 L x
3 SF
1.06 qt
1L
Exact
1 gal = 4 qt
x 1 gal = 1.23 gal
4 qt
Exact
3 SF
Learning Check
If a ski pole is 3.0 feet in length, how long is the ski
pole in mm?
Solution:
3.0 ft x
12 in x 2.54 cm x 10 mm =
1 ft
1 in.
1 cm
Calculator answer: 914.4 mm
Needed answer:
910 mm (2 SF rounded)
Check factor setup: Units cancel properly
Check needed unit: mm
Learning Check
If your pace on a treadmill is 65 meters per minute,
how many minutes will it take for you to walk a
distance of 7500 feet?
Solution:
Given: 7500 ft 65 m/min Needed: min
Plan:
ft
→ in. →
cm →
m → min
Equalities: 1 ft = 12 in. 1 in. = 2.54 cm 1 m = 100 cm
1 min = 65 m (walking pace)
Set Up Problem:
7500 ft x 12 in. x
1 ft
2.54 cm x 1m
x 1 min
1 in.
100 cm
65 m
= 35 min final answer (2 SF)
(# 11): Ethylene glycol, C2H6O2, is an ingredient of
automobile antifreeze. Its density is 1.11 g/cm3 at 20
°C. If you need exactly 500. mL of this liquid, what
mass of the compound, in grams, is required?
Needed: m(g)
Given: d(g/cm3) and V(mL)
1 cm3 = 1 mL
1.11 g
500. mL  ───── = 555 g
1 mL
3 SF
(# 13): A chemist needs 2.00 g of a liquid compound
with a density of 0.718 g/cm3. What volume of the
compound is required?
Needed: V(cm3)
Given: d(g/cm3) and m(g)
1 cm3
2.00 g  ────── = 2.78 cm3
0.718 g
3 SF
(# 15): A sample of 37.5 g of unknown metal is placed in a
graduated cylinder containing water. The levels of the
water before and after adding the sample are 7.0 and 20.5
mL respectively. Which metal in the following list is most
likely the sample?
Metal
d(g/mL)
Metal
d(g/mL)
Mg
1.74
Al
2.70
Fe
7.87
Cu
8.96
Ag
10.5
Pb
11.3
The volume of sample = volume of water
displaced in cylinder = 20.5 – 7.0 = 13.5 mL
one dec. place
After
placing the
piece of
metal
Before
(# 15):
Needed: d(g/cm3)
Given: V(mL) and m(g)
m
37.5 g
d = ── = ────── = 2.78 g/mL
V
13.5 mL
3 SF
From the list, the metal is Al.

Accuracy: nearness of the
measurement to accepted
value of the quantity.

Precision: reproducibility;
how well several
determinations of the same
quantity agree.
Consider a sample that was analyzed for
lead content and was known to contain 49.3 ppm
lead. Two analyses
Analysis A
Analysis B
diff. from
Trial
ppm Pb
Trial
ppm Pb average
1
38.9
1
48.9
4.6
2
23.2
2
59.8
6.3
3
55.9
3
54.5
1.0
4
80.1
4
49.0
4.5
5
46.9
5
55.3
1.8
average = 49.0 ppm Pb
53.5 ppm
average diff. 15.5 ppm
More accurate
3.6 ppm
More precise
Numbers
magnitude, value
 direction: sign (+ or −)
 type of measurement: units
 precision of original measurement:
significant figures

Measured Numbers
A measuring tool
is used to determine
a quantity such as
height or the mass
of an object.
• provides numbers
for a measurement
called measured
numbers.
~4.56 mL
Reading a Meter Stick
. l2. . . . l . . . . l3 . . . . l . . . . l4. .
cm
• The markings on the meter stick at the end
of the orange line are read as
the first digit
2
•
•
plus the second digit
2.7
The last digit is obtained by estimating.
The end of the line might be estimated
between 2.7–2.8 as half-way (0.5) or a little
more (0.6), which gives a reported length
of 2.75 cm or 2.76 cm.
Known & Estimated Digits
In the length reported as 2.76 cm,
• The digits 2 and 7 are certain
•
•
(known).
The final digit 6 was estimated
(uncertain).
All three digits (2.76) are
significant including the estimated
digit.
Significant Figures in
Measured Numbers
Significant figures
• obtained from a measurement
include all of the known digits
plus the estimated digit.
• reported in a measurement
depend on the measuring tool.
Significant Figures
Examples of Counting SF
143.22
 143.0
 300592
 0.0020930
 100.0
 100.
 100



Exact numbers have an
unlimited number of significant
figures
A number whose value is
known with complete certainty
is exact
– from counting individual
objects
– from definitions
1 cm is exactly equal to 0.01 m
– from integer values in
equations
in the equation for the radius
of a circle, the 2 is exact
SF in Calculations
Addition and/or Subtraction

perform operation

round answer to same number of
digits after decimal as number in
calculation with the fewest
Example
132.09 + 35.94376 – 0.0173 =
132.09 + 35.94376 – 0.0173 =
168.01646
must have 2 digits after decimal
168.02
Multiplication and/or
Division

perform operation(s)

round answer to same number of
significant figures as number in the
calculation with the fewest
Example
(26.894)(0.0837)/13 =
(26.894)(0.0837) = 0.1731560
13
must have 2 SF
0.17
Log and/or Antilog
number of digits in mantissa of log =
number of significant figures in antilog
Example
log (14.8003) =
log (14.8003) = 1.17027051857
antilog
1.170271
mantissa
Rounding
first digit to be eliminated is ≥ 5,
round preceding digit up one
 if
 if
first digit to be eliminated is <5,
truncate
Examples
Round each of the following to 4 SF
 10.02700
10.03
 10.02495
10.02
 10.02502
10.03
10.02500
10.03
 10.01500
10.02
 10.02500000000000000000000001
10.03

Exponentials
Scientific notation: very large or very small
numbers are expressed in the following
general form:
exponent term, n = ± integer
N x 10n
digit term, between ± 1 and 9.9999…
(coefficient)
eg
−12,760,000 = −1.276 x 107
0.000012760 = 1.2760 x 10-5
Write the following in scientific notation:
22,400 = 2.24 x 104
22,400. = 2.2400 x 104
892 x 105 = 8.92 x 107
-0.00198 x 10-10 = -1.98 x 10-13
127.60 x 10-5 = 1.2760 x 10-3
Write in fixed notation:
5.720 x 10-2 = 0.05720
-1.982 x 104 = -19,820
Exponentials in calculations
5.750 x 103 + 7.25 x 102
=
1.75 x 10-3 x 6.45 x 102
57.50 x 102 + 7.25 x 102
=
1.75 x 10-3 x 6.45 x 102
64.75 x 102
=
1.75 x 10-3 x 6.45 x 102
64.75
1.75 x 6.45
x
102
10-3 x 102
=
5.74 x 103
3 SF
(0.000345 – 0.0001273) x 6.730x103
 =
154.00
6 dec places (we keep the 77, though)
0.0002177 x 6.730x103
 =
154.00
2.177 has 3 SF only
2.177x104 x 6.730x103
 = 0.009513 = 9.51x103
1.5400 x 102
3 SF
Temperature
Temperature
• is a measure of how hot or
cold an object is compared
to another object.
• indicates that heat flows
from the object with a
higher temperature to the
object with a lower
temperature.
• is measured using a
thermometer.
Temperature Scales
Temperature
Scales
 are Fahrenheit,
Celsius, and
Kelvin.
 have reference
points for the
boiling and
freezing points
of water.
Temperature Scales

Fahrenheit (°F )

Celcius or Centigrade (°C )

9
°F = ── °C +32 = 1.8 °C + 32
5

5
(°F - 32)
°C =  (°F -32) = 
9
1.8

K = °C + 273
Kelvin (K)
273.15 (exact)
ΔT(K) = ΔT(°C) variation of temperature
Learning check
a. The normal temperature of a chickadee is
105.8°F. What is that temperature on the
Celsius scale?
1) 73.8°C 2) 58.8°C
3) 41.0°C
b. A pepperoni pizza is baked at 235°C. What
temperature is needed on the Fahrenheit
scale?
1) 267°F
2) 508°F
3) 455°F
c. Convert 204.3 K into °C.
Other elements to
remember
Sc, Ti, V, Cr, Mn, Fe, Co, Ni, Cu, Zn
 Ru, Rh, Pd, Ag, Cd
 Pt, Au, Hg
