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1
Meet the Elements!
2
WHAT IS CHEMISTRY?
• Chemistry is the study of matter and the
changes that it undergoes.
3
WHAT IS MATTER?
• Matter is anything that has mass and
volume.
• Mass refers to the measure of the amount
of material in an object ( measures the
resistance to an object being moved).
• Volume refers to the amount of space an
object occupies.
4
MATTER HAS 4 STATES
• Solid
• Liquid
• Gas
• Plasma
5
3 STATES OF MATTER
6
7
SOLIDS
• State of matter in which the particles are
closely packed.
• Solids have definite shape and definite
volume
• Solids essentially cannot be compressed.
8
LIQUIDS
• Particles are arranged
so they can slide past
one another.
• Liquids take the shape
of their container and
have definite volume.
• Liquids essentially
cannot be compressed.
9
GASES
• Particles are spread out widely.
• Gases have neither definite shape nor
definite volume.
• Are very compressible.
10
PLASMA
• An electrically neutral gas of ions (charged
particles) and electrons (negatively
charged particles)…VERY HIGH
ENERGY!! Thought to be found in stars.
• Is present when nuclear fusion
occurs..requires a temp of 100,000,000
0C.
11
The smallest unit of matter is
the atom
• are the smallest unit of matter capable of
existing by themselves.
• 2 or more atoms chemically bonded
together are referred to as a molecule.
12
MODELS OF THE ATOM
orbits
13
Atoms are composed of the
following subatomic particles.
• Protons- Positively charged particles located in
the nucleus of the atom.
• Neutrons- Non-charged nuclear particles
located in the nucleus of the atom.
• Electrons- Negatively charged particles that
move around the nucleus located in the electron
cloud
14
Matter with only one type of
atom are called elements.
• Each element has a
unique number of
protons in the
nucleus of its atoms.
This is what defines
that particular
element.
15
ELEMENTS (CONT).
• Chemists use elemental symbols as a
shorthand way of representing elements.
• These symbols consist of 1, 2, or 3 letters
(only the 1st is uppercase).
• Many elemental symbols show the Latin
origin of the element’s name.
16
Periodic Table of Elements
• 114 elements
• Vertical column =
groups
• Horizontal rows =
periods
• Traits are organized
by similar
properties
17
ALL MATTER CAN BE
CLASSIFIED AS EITHER A
SUBSTANCE OR A MIXTURE
• A substance is matter that has a fixed
composition and distinct properties.
• A mixture is a combination of 2 or more
substances, each of which retains it’s
properties.
18
Substances can be either
elements or compounds.
• An element is a substance in which all atoms
have the same number of protons.
• Elements cannot be decomposed into
simpler substances.
19
Compounds
• Compounds are substances in which 2 or
more elements chemically combine to
form a new substance with new
properties.
• Compounds can be decomposed into
simpler substances.
20
COMPOUNDS CAN BE
CHEMICALLY BROKEN DOWN
• This can be accomplished through the use
of heat energy (thermal decomposition) or
electric current (electrolysis).
21
EVERY SUBSTANCE HAS A
UNIQUE SET OF PROPERTIES
• Properties are characteristics that allow us
to recognize and distinguish a substance
from other substances.
22
PHYSICAL PROPERTIES
• Physical properties are properties that can
be observed or measured without
changing the identity and composition of
the matter.
• Physical properties include color, odor,
density, melting pt., boiling pt., malleability,
ductility, hardness, etc.
23
CHEMICAL PROPERTIES
• Chemical properties describe the way a
substance may change or react to form
other substances.
• To observe this property you must carry
out a chemical change.
• An example would be burning in the
presence of oxygen.
24
MATTER UNDERGOES CHANGE
• This may be physical change or chemical
change.
25
PHYSICAL CHANGES
• In physical changes, the substance
changes its physical appearance but not
its composition.
• An example would be changing states.
– Evaporation of water into water vapor
26
CHEMICAL CHANGES
• In a chemical change, a substance is
changed into 1 or more new substances
with new properties.
• Chemical change involves rearrangement
of atoms.
– Aluminum rusting, fireworks exploding
27
PHYSICAL CHANGE
28
INTENSIVE PROPERTIES
• Are independent of sample size.
• Examples include density, temperature,
melting point, etc. (1 gal of water or 1 cup
of water: both have same boil point.)
• Can be used to identify a substance.
29
EXTENSIVE PROPERTIES
• Are properties that depend on sample
size.
• Examples include mass, volume, and
length
• 1 gal of water will have a greater mass
than 1 cup of water.
30
Mixtures can be either
homogeneous or
heterogeneous.
• Homogeneous mixtures are uniform
(have the same composition) throughout.
Homogeneous mixtures are also called
solutions.
Ex: salt, air, water
• Heterogeneous mixtures are nonuniform throughout.
Ex: sand, wood, rock
31
MIXTURES CAN BE
PHYSICALLY SEPARATED
• Filtration
• Distillation
• Chromatography
32
FILTRATION
• In filtration, the
mixture is separated
based on differences
in particle size.
33
DISTILLATION
• Distillation separates components based on
differences in boiling point.
35
DISTILLATION APPARATUS
36
CHROMATOGRAPHY
• Separates mixture
components based on
differences in
attraction to a
surface.
37
40
Name Those Molecules
41
CHEMICAL CHANGE
42
Measurement
1.7
43
SCIENTIFIC DATA FALLS INTO 2
CATEGORIES
• Qualitative data
• Quantitative data
44
Qualitative
• Qualitative data consists of descriptive
terms. There is no use of numbers.
• Determines the presence or absence of a
particular substance in a mix.
– Ex: water is a liquid
45
Quantitative
• Quantitative data consists of
measurements (numbers)
• Determines the amount of substance that
is in a sample.
– Water has a boiling point of 100degrees
Celsius
46
Scientific Method
• Hypothesis : possible explanation of an
observation.
• Theory: set of tested hypothesis that give
an explanation of a natural phenomena
• Law: concise statement or explanation
47
Apples, Volkswagens, Pigeons or
Metric?
UNITS ARE EVERYTHING!
Chemistry is in the details. The slightest
change or mistake can mean a world of
difference. This IS going to drive you nuts
this year, and you will hate chemistry, me,
and the universe, but we go’tta do it!
48
MEASUREMENT
• At one time measurement was inconsistent and
therefore, unreliable.
• Scientists need to be able to repeat each other’s
experiments to verify/modify scientific
knowledge.
• To assist in this process, all scientist use the
SI system of measurement. (AKA your new
best friend!)
49
Base Units
mass
length
time
electric current
temperature
light intensity
amt. Subst.
gram
meter
seconds
ampere
Kelvin
candela
mole
g
m
s
A
K
cd
mol
l
50
MEASUREMENT (CONT)
• Scientist often use prefixes in conjunction
with the basic unit of measure.
• The indicate decimal fractions or multiples
of various units.
– Ex: milli = 10-3
1 milligram = 1 mg = 10 -3 grams
51
Like what is a centimeter again?
52
MEASUREMENT (CONT)
Prefix
Giga
megakilohectodekadecicentimillimicronanopico-
Symbol
G
M
k
h
da
d
c
m
µ
n
p
Exp. Represent.
109
106
103
102
101
10-1
10-2
10-3
10-6
10-9
10-12
MEMORIZE THIS TABLE !!!! ROYO pg 17 table 1.3
53
Question
• What is the name for:
• 10-9 gram =
• 10-6 seconds =
• 10-3 meter =
54
USEFUL STUFF
• Back cover of your text “useful
conversions”
• Memorize:
454 grams =1 lb
2.54 cm = 1 in
1.6 km = 1 mi
1 m = 3.33 ft
1 hr = 3600 sec
55
TEMPERATURE
• Think of temperature as a measure of the
“hotness” of an object.
• Scientist use both the Kelvin (SI) and
Celsius scale to express temperature.
• Kelvin scale is an absolute scale.
Idea: All molecular motion would stop at
0 K.
56
TEMPERATURE (CONT)
• The Celsius scale is based on the freezing
point and boiling point of water.
• 1 Kelvin unit of measure is equal to 1 unit
of measure on the Celsius scale…the 2
scales simply have different “starting
points”
• There is no Degrees symbol in Kelvin
57
58
TEMPERATURE (CONT)
K = 0C +
273
MEMORIZE THESE CONVERSIONS!!!
59
Scientific notation
• Break VERY LARGE or very small
numbers into manageable bits
• The number 123,000,000,000 in scientific
notation is written as :
61
To write a number in scientific
notation:
Put the decimal after the first digit and drop the
zeroes.
62
To find the exponent
• count the number of
places from the
decimal to the end of
the number.
In 123,000,000,000
there are 11 places.
63
UNCERTAINTY IN
MEASUREMENT
• In scientific work, numbers can be exact
or inexact.
• This will be a huge part of your research
paper and how you collect your data!!!!
64
UNCERTAINTY IN
MEASUREMENT (CONT)
• Inexact numbers – numbers obtained by
measurement. They can be inexact due to
human error and or the limitations of the
equipment that you use.
Exact numbers – known values or quantities
Ex: 1 dozen eggs = 12 eggs
All conversion factors are exact numbers!
65
UNCERTAINTY IN
MEASUREMENT (CONT)
• THERE ARE ALWAYS LIMITATIONS IN
THE EQUIPMENT USED TO MEASURE
QUANTITIES (EQUIPMENT ERRORS)
AND THERE ARE DIFFERENCES IN
HOW DIFFERENT PEOPLE MAKE THE
SAME MEASUREMENT (HUMAN
ERROR).
66
UNCERTAINTY IN
MEASUREMENT (CONT)
• UNCERTAINTIES ALWAYS EXIST IN
MEASURED QUANTITIES!!!
67
To Describe Our Measurements ,
We Use Precision and Accuracy
• Precision refers to how close a series of
measurements are to one another.
• Accuracy refers to how closely individual
measurements are to a “true” or accepted
value.
68
Precision, Accuracy, Both ?
69
In general, the more precise a
measurement, the more accurate it
will be
• This is assuming that all measuring
devices are calibrated and the person
doing the measuring is well trained.
• BECAUSE OF THIS, SCIENTISTS
PERFORM MULTIPLE TRIALS OF THE
SAME EXPERIMENT.
70
Which would you use?
71
IDEA: EVERY MEASURING
DEVICE HAS ITS LIMITATIONS
• Because of this, when recording data,
scientists always include an “uncertain”
number (the rightmost digit in the number).
• Tied in with this is the concept of
significant figures.
72
SIGNIFICANT FIGURES (CONT)
• Significant figures are all the certain digits
of a measured quantity plus one uncertain
digit (the last).
• “sig figs” indicate how precise a number is
73
SIGNIFICANT FIGURE RULES
1. All numbers are significant except zeros
at beginning of the number.
9.12
0.00912 0.0912 0.912
all have three SF
Leading Zeros
ARE NOT
significant
74
Rule 2
Zeros between non- zeros are
significant
1005 = 4 SF
1.03 = 3 SF
Captive
Zeros ARE
significant
75
Rule 3
• Zeros that are at the end of a
number and after a decimal point
are significant.
9.00
9.10
90.0
all have three SF
Trailing
Zeros ARE
significant
76
SIG FIGS RULES (CONT)
4. When a number ends in zeros but has
no decimal point, the zeros may or
may not be significant.
We use scientific notation to clarify.
e.g 100 has 1 known sig fig, but
1.00 x 102 has 3 sig figs.
NOTE!!!!!!!
77
Question
• How many Sig figs?
4.003
6.023 x 1023
5000
0.00134
Write 5000 with three sig figs.
78
SIG. FIG CALCULATION RULES
• When multiplying or dividing, the answer has
the same total number of sig figs as the
measurement with the least number of sig
figs.
Example:
1.06 m x 13.400 m = 14.204 m2
14.2 m2
3SF
5SF
3 SF
Round when the answer contains too many digits.
You must have correct SF and UNITS for credit!
79
More Rules!
• In addition or subtraction, the answer has as
many digits after the decimal point as the
measurement with the fewest number of
digits after the decimal point measurement).
– Example:
• 38.5200 cm - 11.4 cm = 27.12 cm
6 SF
3 SF
4 digits
1 digit
27.1 cm
3 SF
1 digit
• When doing multi-step calculations, round at
the end.
80
Question
• Find the volume of a box with the following
measurements:
L = 27.3 cm
W = 15.5 cm H = 5.4 cm
81
Answer
L = 27.3 cm
3 SF
W = 15.5 cm H = 5.4 cm
3 SF
2 SF
2285.01 cm3
->
2.3 x 103 cm3
82
PEMDAS
Please Excuse My Dear Aunt Sally
( ) Zy X ÷ + -
83
Multiple function equations
8.925 - 8.904 x 100 = 0.23529
8.925
84
(9.04 – 8.23 + 21.954 + 81.0) /3.1416
85
9.2 x 100.65
8.321 + 4.026
75
86
• 0.1654 + 2.07 – 2.114 = 0.12
87
• Derived units , uncertainty in
measurement, dimensional analysis
88
RECALL
• SI units tell us our base units
•
•
•
•
Mass = grams g
Time = seconds s
Temperature = Kelvin K
Length = meter m
89
DERIVED SI UNITS (CONT)
• In the lab, scientists use graduated
cylinders, burets, pipets, and syringes to
measure volume.
90
91
PROBLEM
• Sometimes scientist have to convert from
one unit of measure to another (e.g. it
would be inappropriate to describe the
distance from the Earth to the moon in
µm).
• To do so, scientist use a tool called
dimensional analysis or factor-label
method.
92
DIMENSIONAL ANALYSIS
(CONT)
• Dimensional analysis is a problem-solving
technique where a quantity is expressed
equivalently using a different type of unit
through the use of 1 or more conversion
factors.
• This insures that we will calculate the
covert numerical value as well as the
correct units.
93
DIMENSIONAL ANALYSIS
(CONT)
• A conversion factor is a fraction whose
numerator and denominator are the same
quantity expressed in different units.
• E.g. 1 ft/12 inch or 365 days/year
94
DIMENSIONAL ANALYSIS
(CONT)
• What volume, in L, does a 37.3 mL sample
occupy?
• 37.3 mL x
__10-3_L____ = 0.0373 L
1
mL
95
DIMENSIONAL ANALYSIS
(CONT)
• How many cm exist in a 8.50 in sample?
• 8.50 in x 2.54 cm
1 in
= 21.6 cm
96
General Rule of Thumb
1. Begin the conversion by looking at the
units given and desired.
2. Identify one/ many conversion factors that
will take you from the given to the desired
units
Given unit x Desired Unit = Desired Unit
Given Unit
97
Convert 8.00 m to in
8.00 m x 1 cm x
10-2 m
1 in
= 315 in
2.54 cm
OR
8.00 m x 102 cm x 1 in
1m
2.54 cm
= 315 in
98
Question
1.) 3.85m to mm
2.) How many cm are in 6.5 miles (hint start
with what you know)
99
Answer
3.85 m x 1mm = 3850 m
10-3m
6.5mi x 1600m x 102cm = 1.0 x 106 cm
1 mi
1m
6.5mi x 5280 ft x 12in x 2.54 cm = 1.0 x 106cm
1 mi
1 ft
1 in
6.5 mi x 1 km x 105cm = 1.0 x 106cm
0.62137 mi 1 km
100
DERIVED SI UNITS
• Derived SI units are used to measure
quantities expressed as combinations of
SI units.
• Area =
length in m x width in m = m2
101
COMMON DERIVED UNITS
Want to find
Formula
Derived Unit
Area
length x width
m2
Volume
length x width x
height
m3
Density
mass/volume
g/L
Molar mass
mass/amt of
substance
g/mol
Concentration
amt. of
substance/volume
Mol/L
Energy
force x length
N-m
102
DERIVED SI UNITS (CONT)
• Volume = l x w x h = m x m x m = m3
The volume of liquids is often measured in
mL.
1 cm3 = 1 mL
1 dm3 = 1 L
103
Cubing conversation factors
• 13 cm3 to dm3
104
105
Density
• Mass in a unit volume of a substance
Density = Mass / Volume
Units = g/cm3 or g/ml
MEMORIZE!!!!
Density of Water = 1.00 g/ml
106
Density Cont.
107
Question
1. Find the density of 1.00 x 102 g of
mercury that occupies a volume of 7.36
cm3 .
2. Find the volume of 65.0 grams of
methanol with a density of 0.791 g/ml
108
Challenge question
• What is the mass in grams of a cube of
gold (d = 19.32 g/cm3). The length of the
cube of gold is 2.00 cm
109
Answer
• Find the mass from the volume of the cube
and its density. Volume of a cube is given
by its length.
Mass = Volume x Density
8.00 cm3 (19.32 g) = 154 g
1 cm3
110
• Flash Example
111