Transcript Chapter 1 Basic Concepts of Matter

Chapter 1
Basic Concepts of Matter
Classifying Matter
Matter
– Anything that has mass and occupies space
Mass vs. Weight
Kinetic-Molecular Theory
All matter consists of extremely tiny
particles in constant motion
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States of Matter
• Solid
– -Closely packed together with a definite ridged shape
– -Vibrate back and forth in a confined space
– -the particles are not able to move past one another
• Liquid
– -arranged randomly with a definite volume
– -“fluid”
– -the particles are not confined in space and can move
past one another
• Gas
– -no definite shape or volume
– -“fluid”
– -the particles are far apart and move very rapidly
colliding with other particles and the container walls
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Categorizing Matter
• Elements
– -cannot be decomposed into simpler form via
chemical reactions
– -found on periodic chart
– -atoms are the smallest particle that retains
the characteristic properties of the elements
• Pure Substance
– -consists of all the same substance (pure
gold, distilled water, etc)
– -have a set of unique properties that
identifies it
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Categorizing Matter
• Chemical Compounds
– -two or more elements in a definite ratio
by mass with unique properties that
separate them from the individual elements
– -can be decomposed into the constituent
elements by chemical reactions
– -chemical compounds are held together by
a chemical bound
• Water– hydrogen and oxygen
• Carbon dioxide – carbon and oxygen
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Categorizing Matter
• Mixtures
– two or more pure substances in the same
container
homogeneous mixtures (solution)
• -uniform composition throughout
• -single phase
• -cannot be separated easily
heterogeneous mixtures
• -nonuniform composition thoughout
• -easily separated
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Physical and Chemical Changes
Physical changes
• changes in physical properties
• -melting, boiling, and cutting
Chemical changes
changing one or more substances into one or
more different substances (chemical reaction)
• 2H2 + O2 -> 2H2O
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Chemical and Physical Properties
Chemical Properties
observed during a chemical reaction (change
in chemical composition)
– -rusting, oxidation, burning…
– -chemical reactions
Physical Properties
observed without changing the substance’s
composition
– -allow for identification and classification
– -density, color, solubility, melting point…
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Classification of Physical Properties
Extensive Properties
depend on the amount of substance
present
-mass or volume
Intensive Properties
do not depend on the amount of substance
present
-melting point, boiling point, density…
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Density
• Describes how compact a substance is
• Who “discovered” density?
• Density = mass/volume or D = m/V
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Density
• Example: Calculate the density of a
substance if 742 grams of it occupies
97.3 cm3.
• 1cm3 = 1mL => 97.3cm3 = 97.3mL
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Density
• Example: You need 125 g of a corrosive
liquid for a reaction. What volume do you
need?
•
– liquid’s density = 1.32 g/mL
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Units of Measure
• Qualitative measures
– Nonnumerical experimental observations
describing the identity of a substance in a
sample
• Quantitative measures
– Numerical experimental observations describing
how much of a particular substance is in a
sample
System International d’Unites (SI)
measurement system used in the sciences
based on the metric system
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Math Review and Measurements
• We make measurements to understand our
environment:
– Human senses: sight, taste, smell, hearing…
• Our senses have limits and are biased
– Instruments: an extension of our senses meter sticks,
thermometers, balances
• These are more accurate and precise
– All measurements have units
• METRIC SYSTEM vs. British System
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SI units
Quantity
 length
 mass
 time
 current
 temperature
 amt. substance
Unit
meter
kilogram
second
ampere
Kelvin
mole
Symbol
m
kg
s
A
K
mol
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Measurements in Chemistry
Name
• mega• kilo• deka• deci• centi-
Symbol
M
k
da
d
c
Multiplier
106 (1,000,000)
103 (1,000)
10
10-1 (0.1)
10-2 (0.01)
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Measurements in Chemistry
Name
• Milli• Micro• Nano• Pico• Femto-
Symbol
m

n
p
f
Multiplier
10-3(0.001)
10-6(0.000001)
10-9
10-12
10-15
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Units of Measurement
Length
– Measure of space in any direction
– -derived unit cm
– -standard length is a meter (m)
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Units of Measurement
• Volume
– Amount of space occupied by matter
– -derived unit: mL or cm3 (cc)
– -liter (L) is the standard unit
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Units of Measurement
• Time (t)
– Interval or duration of forward events
– -standard unit is the second (s)
• Mass (m)
– measure of the quantity of matter in a body
1 kg = 1000g
1 kg = 2.2 lbs
1 g = 1000mg
• Weight (W)
– measure of the gravitational
– attraction (g) for a body (w=m x g)
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Heat and Temperature
• Heat (q) vs. Temperature (T)
• 3 common temperature scales:
• all use water as a reference
• -Fahrenheit (F)
• -Celsius (C)
• -Kelvin (K)
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Temperature Reference Points
• Fahrenheit
•
Celsius
•
Kelvin
Melting Point
of water
32 oF
0.0 oC
273 K
Boiling Point
of water
212 oF
100 cC
373 K
• Body temperature 37.0 oC or 98.6 oF
– 37.2 oC and greater—sick
– 41 oC and greater, convulsions
– <28.5 oC hypothermia
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Temperature Scales
Kelvin and Centigrade Relationsh ips
K  C  273.15
or
o
o
C  K  273.15
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Temperature Scales
Fahrenheitand
to Centigrade
Centigrade Relationsh
Relationships
ips
Fahrenheit
o
9 o
F  1.8  C  32  x C  32
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or
o
F  32 5

C
9
1.8
o
o
 F - 32
o
•Example: Convert 211 oF to degrees Celsius.
Example: Express 548 K in Celsius degrees.
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Precision and Accuracy
Accuracy
how closely measured values agree with the correct value
Precision
how closely individual measurements agree with each other
Precise
Neither
Accurate
Both
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Mathematics in Chemistry
• Exact numbers (counted numbers)
– 1 dozen = 12 things
• Measured Numbers
– Use rules for significant figures
– Use scientific notation when possible
• Significant figures
– digits in a measured quantity that reflect the
accuracy of the measurement
– -in other words, digits believed to be correct by
the person making the measurement
– Exact numbers have an infinite number of
significant figures
12.000000000000000 = 1 dozen
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Significant figures (numbers/digits)
Why use significant numbers?
• -Calculators give 8+ numbers
• -People estimate numbers differently
• -Dictated by the precision (graduation) on
your measuring device
• -In the lab, the last significant digit is the
digit you (the scientist) estimate
Scientists have develop rules to help
determine which digits are “significant”
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Rules for Significant Figures
1. All Nonzero numbers are significant!!!
2. Leading zeroes are never significant
– 0.000357
3. Imbedded zeroes are always significant
– 3.0604
4. Trailing zeroes may be significant
- You must specify significance by how the
number is determined or even written
– 1300 nails - counted or weighed?
– 1.30000 –How many significant figures?
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Significant Figures
• Multiplication & Division rule:
• The product retains the number of significant figures
that corresponds to the multiplier with the smallest
number of significant figure (sig. fig.)
4.242
x 1.23
2.7832
x 1.4
5.21766
round off to 5.22
3.89648
round off to 3.9
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Significant Figures
• Addition & Subtraction rule:
– Answer retains the smallest decimal place
value of the addends.
3.6923
 2.02
18.7937
 2.123
6.9463
round off to 6.95
16.6707
round off to 16.671
 1.234
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Scientific Notation
Express answers as powers of 10 by moving
the decimal place right (-) or left (+)
• Use of scientific notation is to remove doubt
in the Significant Figures:
2000  2 x 103
15000  1.5 x 10?
0.004  4 x 10-3
0.000053  __.__ x 10?
Key to Sig. Figs…
Locating the decimal and deciding when to
count Inthe
zeros!!!
scientific
notation, zeros are given if they are
significant!!!
1.000 x 103 has 4 significant figures
2.40 x 103 has ? significant figures
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Review #2
• Units of Measure
–
–
–
–
–
-length
-volume
-time
-mass
-weight
• Heat vs. Temperature
– -three temperature scales
– -temperature conversions
• Precision vs. Accuracy
• Significant Figures
• Scientific Notation
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Conversion Factors
• Length
– 1 m = 39.37 inches
– 2.54 cm = 1 inch
• Volume
– 1 liter = 1.06 qt
– 1 qt = 0.946 liter
• See Text for more conversion factors
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Conversion Factors
Why do conversions?
• -Scientists often must convert between units
Conversion factors can be made for any
relationship of units
-Use known equivalence to make a fraction
that can be used to “convert” from one unit
to the other
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Dimensional Analysis
•
1 inch = 2.54 cm
1in
2.54 cm
2.54 cm
or
1in
– Use the ratio to perform a calculation so the
units will “divide out”
Example: Convert 60 inches to centimeters
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Dimensional Analysis
• Example: Express 9.32 yards in
millimeters.
3 ft = 1 yard
1 ft = 12 in
1 in = 2.54 cm
100 cm = 1 m
1000 mm= 1 m
12 in
1 ft
or
1 ft
12 in
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Dimensional Analysis
• Example: Express 627 milliliters in
gallons.
1 liter = 1.06 qt
1 qt = 0.946 liter
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Practice on your Own
1kg = 2.20 lbs
– Convert 25 g to lbs
– Convert 1 mL to Liters
– Convert 20 meters to cm
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Dimensional Analysis
Area = length x width
Area is two dimensional thus units must be
in squared terms:
• Express: 2.61 x 104 cm2 in ft2
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Dimensional Analysis
Volume =length x width x height
• Volume is three dimensional thus units
must be in cubic terms
Express: 2.61 ft3 in cm3
– this volume is used in medical
measurements--cc
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Percentage
• Percentage is parts per hundred of a
sample
g of substance
• %=
x100
total g of sample
• Example: A 335 g sample of ore yields 29.5 g
of iron. What is the percent of iron in the ore?
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