Honors Chemistry
Chapter 1
Matter and Measurement
1.1 What is chemistry?
• Chemistry is the study of the properties of
matter and their changes.
• Matter: The physical material of the
universe.
– Anything that has mass and
– Occupies space (volume)
Matter
• Matter: Anything that has mass and takes
up volume
• Mass: the amount of material (stuff) in an
object (kg, g, lb, t )
• Volume: the amount of space an object
occupies (cm3, mL, L, ci)
Elements
• Very basic elementary
substances.
• Made of unique atoms.
• Can not be broken down further
without changing the identity.
• Named for Latin or Greek
characters or places.
• Represented by 1 or 2 letter
symbols.
– First letter CAPITALIZED
– Second letter lower case
• C, O, Co
• CO vs. Co
Hydrargyrum
Greek for “liquid
silver”
Molecules
• Combinations of two
or more atoms held
together in specific
shapes.
• Properties are based
on the structure and
arrangement , and the
number and types of
atoms present.
Methane
Classifications of Matter
•
•
•
How is matter classified? (2 ways)
1. Macroscopic level (more than one
molecule)
a. Gas or vapor (g): no volume or shape, highly
compressible
•
•
•
gases expands to occupy it’s container.
b. Liquid (l):
volume independent of container, no
shape, incompressible (relatively)
c. Solid (s):
volume and shape independent of
container, rigid, incompressible.
2. Molecular level (single
molecules)
•
a. gas:
•
b. liquid:
molecules closer than gas, move rapidly,
slide over each other
•
c. solid:
molecules packed closely, definite
arrangement, don’t move.
molecules far apart, high speed, collide
often with each other, and container walls.
(kinetic molecular [KM] theory)
States of matter (4)
• 1. Solids
• Hold a particular
shape.
• Have a definite
volume.
• Particles arranged in
an orderly manner.
• High density
States of Matter
• 2. Liquid
• Does not hold its own
shape, takes the
shape of its
container.
• Has definite volume.
• Particles arranged
randomly
• High density
• Not compressible
States of Matter
• 3. Gas
• No definite shape or
volume.
• Expands to occupy the
space of its container.
• Extremely low density.
• Particles far apart.
• Highly compressible.
• Density depends on
pressure.
– High pressure forces
particles closer.
States of Matter
• 4. Plasma
• Ionized gasses.
• Ions and electrons
(charged particles)
• Found in the center of
stars and space.
• Emit light when “excited”
• Eg. Lightning, Northern
Lignts
Matter
Matter can be classified into four
groups.
Heterogeneous
Matter
Homogeneous
AKA:
Solution
Mixtures
Pure
Substance
Compound
Element
Pure
Substances
(2 types)
1. Pure substance is a
substance with
constant composition
2. Pure substances have
fixed measurable
properties
3. Pure substances
cannot be found in
nature.
C6H12O6
H 2O
Elements (pure substance #1)
1. Simplest form of matter.
2. Can not be broken down
by ordinary chemical
means.
3. 92 naturally occurring
elements.(118 total)
4. Arranged in the periodic
table.
5. Solids, Liquids, Gases
(at room temp)
Symbol
Atomic Number
Mass Number
http://www.funbrain.com/periodic/
(Atomic Weight)
Element names and symbols to
memorize. (on periodic table)
2A
Be
Mg
Ca
Sr
Ba
3A
B
Al
4A
C
Si
Sn
Pb
5A
N
P
6A
O
S
7A
F
Cl
Br
I
8A
He
Ne
Ar
Kr
Rn
K
R
b
1A
H
Li
Na
K
Rb
Transition Metals to Memorize
•
•
•
•
•
•
•
•
•
•
Cr
Mn
Fe
Co
Ni
Cu
Zn
Ag
Hg
Au
•
•
•
•
•
•
•
•
Diatomic Elements
H2
N2
O2
F2
Cl2
Br2
I2
• BrINClHOF
Compounds (pure substance #2)
Two or more elements
combined in a
chemical reaction.
• Elements are in fixed
proportions to each
other.
• Written using symbols
with subscripts to
denote the ratio of
elements.
• Can not be separated
by physical means.
H2O:
2 Hydrogen: 1 Oxygen
Mixtures
• A combination of two
or more pure
substances.
• Two types of
mixtures.
Heterogeneous mixture
(hetero: different)
• A MIXTURE is a combination
of
– two or more pure
substances that are
– not chemically united and
– do not exist in fixed
proportions to each other.
– Do not have uniform
composition.
• Most natural substances are
mixtures.
• A heterogeneous mixture
consists of visibly different
substances or phases.
• A mixture can be physically
separated into pure
compounds or elements.
Methods of separating mixtures
• 1. Decantation:
pouring a less dense
liquid off the top of a
more dense liquid.
• 2. Filtration:
Separating particles
based on size or
phase. Larger
particles are trapped
in the filter.
• Liquid passes through
filter, solid is trapped.
Separating
Funnel
Homogeneous Mixture
(Homo: Same) Solution
• A homogeneous mixture has
the same uniform
appearance and
composition throughout.
Many homogeneous mixtures
are commonly referred to as
solutions.
• All components are all in the
same phase.
• Particles are uniform in size
(atoms or molecules)
• Can not be separated by
physical means.
• Can be separated based on
differences in properties of
components.
Methods of separating
solutions
• 3. Crystallization:
Evaporating one
component to leave
another.
• 4. Chromatography:
Separating components
based on specific gravity.
• 5. Distillation:
Separation of liquids
based on differences in
boiling point.
Properties of compounds vs.
elements.
• A compound may have different properties
than the elements composing it.
• Hydrogen H2: Gas
• Oxygen O2:
Gas
• Water H2O: Liquid
Joseph Proust (1799)
• French Chemist
• Developed The Law of
Definite Proportions
• Aka the Law of
Constant Composition
• Compounds always
contain the same
elements in the same
proportion by mass.
Law of Definite Proportions
• H20 (by mass is
always)
• 88.9% Oxygen, 11.1%
Hydrogen
• If we had an 80g
sample of H20 how
much is O?
• .889 x 80 = 71g
• How much is H?
• .111 x 80 = 9g
1.2 Properties of Matter
• Physical Properties:
• Characteristics that
can be observed and
measured without
altering the identity
of the substance.
• Density, color, melting
point, odor, boiling
point, etc.
Ice, water and steam.
Physical Properties: 2 Types
• 1:Intensive Properties:
• Do not depend on the
amount of matter present.
• Examples:
–
–
–
–
–
–
Density
Conductivity
Melting point
Boiling point
Temperature
Pressure
Physical Properties: 2 Types
• 2: Extensive
Properties
• Depends on how
much material is on
hand.
• Changes with the
amount of material
present.
• Examples:
– Mass
– Volume
– Weight
Chemical Properties
• The way matter
behaves when
brought into contact
with other
substances, or a
source of energy.
• Describes how
substances react with
other substances.
• Example:
– Flammability
– Inert (nonreactive)
“Oh the humanity”
Changes in Matter
• Physical Changes:
• Alter the form, but do not
alter the identity of the
substance.
• Change in physical
appearance.
• Examples:
– Crushing
– Tearing
– Changes of state (phase)
Chemical Changes
(Reactions)
• Chemical Changes
(reactions)
• Changes a substance
into chemically different
substances.
• Irreversible
• Alters the identity of the
substance being
changed.
• Examples:
– Wood burning,
– Food cooking,
– Iron rusting
Homework Practice
• Exercises #s 2, 3, 5, 10, 12, 15
• On page 29 in Textbook
1.4 Units of Measurement & the Metric System
A.
Why Measure? In order to get complete observations,
quantifying is essential.
B.
Consists of:
1.
2.
C.
Quantity
a.
Anything that can be measured
b.
Involves a number followed by a unit
c.
Examples: Volume, Mass, Length, Time
Unit (can be English (customary) or metric)
a.
Is what unit the quantity is measured in
b.
Examples: Liter, Gallon, Gram, Foot, Meter, Second
Why Use the Metric System?
1.
It’s easier because it is based on units of 10
2.
Everyone uses it so it gives consistency
3.
There are standards that everyone can compare to
D.
International System of Units – called SI
1.
Composed of fundamental or base units.
Chart 1-1
Quantity
Standard Unit
Symbol
Length
Meter
m
Mass
Kilogram
kg
Time
Second
s
Amount/Count
Mole
mol or m (bar)
Temperature
Kelvin
K
Electric Current
Ampere
A
Luminous Intensity
Candela
cd
Prefixes used in the Metric System
Greek mu
Page 14 in your book
Length and Mass
• SI unit of length is the meter (m) (ca. 39
inches)
• Mass: The measure of the amount of
material (matter, stuff) in an object.
• SI base unit of mass is the kilogram (kg)
(ca. 2.2 lbs)
• This is odd because it uses a prefix
instead of gram alone.
Temperature
• The measure of how fast molecules in a
substance are moving.
» Or
• The measure of the average kinetic
energy of particles in a substance.
Temperature scales
• Fahrenheit
• Named after Gabriel
Fahrenheit (1686-1736)
• Thermometer maker who
devised his own
temperature scale.
• Based on the
temperature of an equal
ice-salt mixture. (0ºF)
• Average temperature of a
healthy horse. (100 º F)
• H2O freezes at: 32º F
• H20 boils at 212º F
Why me?
Say aaah.
Temperature Scales
• Celsius (Centigrade)
• Anders Celsius
• Developed a
temperature scale
more compatible with
the metric system.
• Based on H20
– H2O freezes at: 0º C
– H20 boils at 100º C
(1701- 1744)
Thanks Mr.
Celsius!!
Fahrenheit vs Celsius
• Fahrenheit
– 212º - 32º = 180º from
boiling to freezing.
• Celsius
– 100º - 0º = 100º from
boiling to freezing.
• 180/100 = 9/5
• Celsius degrees are 9/5
bigger than Fahrenheit
• The Fahrenheit scale also
begins 32º above the
Celsius scale.
Converting between Fahrenheit
and Celsius
• Taking the degree size, and starting point
in consideration, we get:
• ºF = 9/5ºC + 32 and/ or
• ºC = 5/9(ºF – 32) *watch order of ops*
Conversion Practice
• The temperature of a healthy human is 98.6ºF.
Convert this to Celsius.
• ºC = 5/9 (98.6ºF – 32)
• 37ºC
• A baby has a fever of 39.5ºC. What is her
temperature in Fahrenheit?
• ºF = 9/5 x 39.5ºC + 32
• 103.1 ºF
The Kelvin Temperature Scale
• The SI unit used to
measure temperature.
• Named for William
Thomson, Lord Kelvin
• Not Melvin Kelvin
• The º symbol is not
used for kelvin (K)
• 1ºC = 1K
1824-1907
The Kelvin Temperature Scale
(aka Absolute Scale)
• The zero point on the
Kelvin scale is absolute
zero
• 0K = -273ºC
• Absolute zero is the
point at which all
molecular motion ceases.
Lowest temperature
possible.
• Can absolute zero ever
be reached?
Converting between Celsius and
Kelvin
• ºC = K – 273
• At 50K air will freeze to a solid. Convert this
to ºC.
• ºC = 50 – 273 = -223 ºC
• K = ºC + 273
• Antifreeze (ethylene glycol) boils at 199ºC.
Convert this to K.
• K = 199 + 273 = 472 K
Comparing temperature scales
• Note that in ºC, and K,
there are 100 degrees
between the freezing and
boiling points of H20.
• Note that in ºF there are
180º between the
freezing and boiling
points of H2O.
• The Celsius scale begins
273.15º higher than
Kelvin.
• The Fahrenheit scale
begins 32º higher than
Celsius.
Page 15
Derived units/ Ratios
• Derived units are obtained by performing
mathematical calculations on base
units.
• Example:
Speed = distance
time
• What is the speed of a car (in mph) which
travels 53miles in 48 minutes?
– 53miles 60min = 66 miles/hour
– 48min
1 hour
Volume
• The most commonly
used metric unit used
to express volume is
the Liter (L).
• The liter is not an SI
unit.
• The SI unit of
volume is the cubic
meter(m3). Too big to
be convenient.
Other units of volume.
•
•
•
•
Cubic centimeter (cc or cm3)
1 cm3 = 1 mL
Cubic decimeter (dm3)
1 dm3 = 1 L
Density
• Ratio of mass to volume
• Calculated by dividing the mass of an
object by its volume.
• Density = mass
volume
· Expressed in g/cm3
· Based on water.
3
The
density
of
pure
H
O
is
1.00
g/cm
·
2
Density
• Example: If a metal block has a mass of
75g and a volume of 22 cubic
centimeters, what is its density?
75g
22cm3
= 3.4 g/cm3
D. Density
• An object has a volume of 825 cm3 and a
density of 13.6 g/cm3. Find its mass.
GIVEN:
WORK:
V = 825 cm3
D = 13.6 g/cm3
M=?
M = DV
M
D
V
M = (13.6 g/cm3)(825cm3)
M = 11,200 g
D. Density
• A liquid has a density of 0.87 g/mL. What
volume is occupied by 25 g of the liquid?
GIVEN:
WORK:
D = 0.87 g/mL
V=?
M = 25 g
V=M
D
M
D
V
V=
25 g
0.87 g/mL
V = 29 mL
1.5 Uncertainty in Measurement
• Exact numbers: known by counting or definition.
Whole numbers and integers.
• Inexact numbers: Derived from measurement.
There is always estimation involved. The last
digit is the uncertain digit.
Scientific Method and Measurement
Lesson 2: Scientific Measurement
I.
1.5 Uncertainty in Measurement
A.
There is ambiguity in every measurement. Why?
1.
Instruments are never completely free of flaws.
(More expensive instruments are generally more accurate)
2.
perfect
Measuring involves estimation – because humans are
reading, there may be human error!! Practice makes
a.
Digital Display – electronic balance
b.
Scale Display – grad. cylinder, thermometer, ruler
* The last number in any measurement is estimated and therefore
uncertain
Accuracy is assessed by determining percent error
% Error = [(Exp. Value – Accepted Value) / Accepted Value] x 100
* Note – the experimental value will be the average if more than one
trial is performed for the experiment.
Accuracy: how close a measurement is to the true or accepted
value
Precision: how reproducible the measurements are (how closely
related the measurements are to each other)
Accurate and
Precise
Precise, not
accurate
Accurate,
not precise
Neither
accurate nor
precise
A. Accuracy vs. Precision
• Accuracy - how close a measurement is
to the accepted value
• Precision - how close a series of
measurements are to each other
ACCURATE = CORRECT
PRECISE = CONSISTENT
Significant Figures
• Used to indicate the exactness of a
measurement.
– 1cm vs. 1.0cm
– 2.5s vs 2.52s
• The final digit represents the estimated digit.
• If using for a derived unit, the measure is only
as accurate as the least accurate
measurement.
1.
Rules for Individual measurements
a.
All non-zero digits are significant. 1234 (4 siggy figs)
b.
All zeros between non-zero digits are significant. 101(3)
c.
Zeros to the right of a non-zero digit, but left of the
decimal point are not significant unless specifically
designated with a bar above it. 1100 vs 1100
d.
Zeros to the right of the decimal point but to the left of
a non-zero digit are not significant. .00038
e.
Zeros to the right of the decimal point and following a
non-zero digit are significant. 3.0027
* There will never be a bar on the right side of the decimal!!
2.
Eg: How many significant figures for the measurements?
a.
173.2 cm
4
_____
e.
.025 L
2
_____
b.
205 kg
3
_____
f.
700 mL
1
_____
c.
4000 mm
1
_____
g.
.090500 Mg
5
_____
d.
6.7 x 103 km
2
_____
3.
4.
Rules for Operations – Addition & Subtraction
a.
Perform operation, then round answer to the least
accurate measurement (place value) for the
measurements you are adding or subtracting.
b.
Eg: Perform the operation & then round the answer to
the correct # of significant figures.
1.)
6.345 + 5.78 – 3.27 + 4.927 =
13.78
________
2.)
.045 +.22 -.20 =
________
.07
Rules for Operations – Multiplication & Division
a.
b.
Perform operation, then round answer to the least
number of significant digits for the measurements
given.
Eg: Perform the operation and round answer to the
correct # of significant figures.
VII.
1.)
4.50 x .0062
.028
________
2.)
(5.622)(72.3) / (320)
1.3
________
3.)
(6.5 x 10-2) (3.27 x 10-3)
2.1
x 10-4
________
4.)
(2.35)(.45723)(546)
(379)(4.3 x 104)
3.6
x 10-5
________
1.6 Problem Solving
A.
Dimensional Analysis: technique for converting between units
1.
Within the Metric System – move the decimal for how
many jumps the multiple makes on the prefix chart 1-3
a.
Ex: Convert 43 cm to . . .
1.)
m
.43
________
2.)
km
________
3.)
um
430,000
________
.00043
b.
2.
Convert 276 g to . . .
1.)
mg
276,000
________
2.)
Mg
.000276
________
From metric to English or English to metric
a.
Write down what you are given
b.
Write down what you are working to
c.
Determine the conversion factor (unit equality)
d.
Cancel the units
e.
Do the arithmetic
f.
Convert the following;
1.)
47 miles to km
76
________
2.)
105 yards to m
96.0
________
3.)
75.2 kg to pounds
166
________
4.)
15 gallons to L
57
________
Significant Digits
B. Dimensional Analysis
• The “Factor-Label” Method
– Units, or “labels” are canceled, or “factored”
out
g
cm 

g
3
cm
3
B. Dimensional Analysis
• Steps:
1. Identify starting & ending units.
2. Line up conversion factors so units cancel.
3. Multiply all top numbers & divide by each
bottom number.
4. Check units & answer.
B. Dimensional Analysis
• How many milliliters are in 1.00 quart of
milk?
qt
mL
1.00 qt

1L
1000 mL
1.057 qt
1L
= 946 mL
B. Dimensional Analysis
• You have 1.5 pounds of gold. Find its
volume in cm3 if the density of gold is
19.3 g/cm3.
cm3
lb
1.5 lb 1 kg
2.2 lb
1000 g
1 cm3
1 kg
19.3 g
= 35 cm3
B. Dimensional Analysis
• How many liters of water would fill a
container that measures 75.0 in3?
in3
75.0
L
in3
(2.54 cm)3
1L
(1 in)3
1000 cm3
= 1.23 L
B. Dimensional Analysis
5) Your European hairdresser wants to
cut your hair 8.0 cm shorter. How many
inches will he be cutting off?
cm
in
8.0 cm
1 in
2.54 cm
= 3.2 in
B. Dimensional Analysis
6) Taft football needs 550 cm for a 1st
down. How many yards is this?
cm
550 cm
yd
1 in
1 ft
2.54 cm 12 in
1 yd
3 ft
= 6.0 yd
B. Dimensional Analysis
7) A piece of wire is 1.3 m long. How
many 1.5-cm pieces can be cut from
this wire?
cm
1.3 m
pieces
100 cm 1 piece
1m
1.5 cm
= 86 pieces