Chemistry 101 : Chap. 1

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Chemistry 101 Chap. 1
Matter and Measurement

• What is Chemistry and Why we study it
• (2) Classification of Matter
• (3) Properties of Matter
• (4) Units of Measurement
• (5) Uncertainty in Measurement
• (6) Dimensional Analysis

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The Study of Chemistry

• Chemistry The study of the properties of matter
  and the
• changes that matter
  undergoes.

? Matter Physical material of the universe
Anything that has mass and occupies
space ? Changes in Matter Physical or Chemical
changes

• Why Chemistry?

• ? Chemistry is the central science
• ? Chemistry is a practical science and has
  profound
• impact on our daily living

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Macroscopic vs. Microscopic

• Macroscopic World Realm of ordinary-sized
  object.
• Things we
  can see with the naked eye.

? (Sub)Microscopic World Realm of atoms and
molecules
Carbon nanotube (10-9 m)
Chemistry is the science that seeks to understand
the properties and behavior of matter
(macroscopic) by studying the properties and
behaviors of atoms and molecules (microscopic)
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Major Divisions in Chemistry
? Physical Chemistry (CHM321, CHM420) ? Organic
Chemistry (CHM211, CHM212) ? Inorganic Chemistry
(CHM 455, CHM546) ? Analytical Chemistry (CHM235,
CHM435) ? Biochemistry (CHM365, CHM568)
All divisions are interrelated and cannot
be standing alone.
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Classification of Matter pure substance vs.
mixture

• Pure Substance A sample of matter that has
  distinct
• properties and a composition that doesnt vary
  from sample
• to sample (either element or compound)

• Elements A pure substance that cannot be
• decomposed into simpler substances. The basic
  unit
• of an element is an atom.

Nitrogen atom
Nitrogen molecules
Argon gas (atoms) Nitrogen gas (molecules)
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Classification of Matter pure substance vs.
mixture

• Compound Substances that are composed of
• two or more elements. The basic unit of
  compound is
• a molecule

• Mixture Combinations of two or more substances
  
• in which each substance retains its own chemical
  identity.

nitrogen atom
Two or more elements (compound)
Two or more substances (mixture)
hydrogen atom
ammonium (molecule)
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Elements
? At the present time, there are 116 elements
Periodic Table of the Elements
H2, N2, O2, F2, Cl2, Br2, I2
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Elements
? Not all elements are equal
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Compounds
? Most elements can interact (or react) with
other elements to form compounds
Example Combine hydrogen oxygen to generate
water
Oxygen Hydrogen
water
However, elemental hydrogen and oxygen exist as
diatomic molecules (H2 and O2) in nature.

2H2O
O2 2H2
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Mixture
? Components The substances making up a mixture

• Homogeneous Mixture (solution) Uniformly
  distributed
• throughout. (air, salt solution, sugar
  solution )

• Heterogeneous Mixture Do not have the same
• composition, properties and appearance
  throughout.
• (rock, wood )

Oil on water
Air
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Classification of Matter
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Classification of Matter
? Example

• 14 K gold
• (2) Orange Juice
• (3) A cup of coffee
• (4) Mud

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Separation of Mixture
Separate a mixture into its components by taking
advantage of the difference in
their properties

• Filtration Separation is based
• on the size of particles in the
• mixture. Filtration is used with
• heterogeneous mixtures

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Separation of Mixture
? Distillation Separation is based on the
boiling points of the components in the
mixture. Distillation is typically used with
homogeneous solutions.
Water changes its states from gas to liquid
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Separation of Mixture
? Chromatography Separation is based on the
solubilities of the components in the mixture.
It is normally used with homogeneous mixture.
Paper chromatography
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Classification of Matter states of matter

• States of matter A sample of matter can have
  three
• physically
  different states

? Gas Indefinite volume and indefinite
shape (depends on the volume and shape of
its container) ? Liquid definite volume, but
indefinite shape. ? Solid definite volume and
definite shape
Pure substance can have any state depending on
the temperature and pressure
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Three States of Water
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Properties of Matter

• Physical properties They can be measured
  without
• changing the identity and composition of the
  substance
• Ex. color, order, density, boiling point

? Chemical properties They describe the way a
substance can change or react Ex.
flammability, solubility,
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Physical vs. Chemical Properties
? Example Zinc (Zn)
silver-grey metal
melting point 420oC
reacts with oxygen to form Zinc oxide (ZnO)
density (25oC) 7.13 g/cm3
generates hydrogen when dissolved in sulfuric
acid
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Properties of Matter

• Extensive properties Properties that depend on
• the quantity of a sample.

Ex. Volume ? ? ? V1
V2 V1 V2
? Intensive properties Properties that are
independent on the quantity of a sample
Ex. Temperature ? ? ?
T T T
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Extensive vs. Intensive Properties
? Example
Boiling/melting point (bp/mp) Mass Density Pres
sure
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Changes of Matter

• Physical changes
• Phase changes, but it is
• still H2O (no change in its
• composition)

• Chemical changes
• Aluminum (Al) reacts with
• Bromine (Br2). (A substance
• is transformed into a chemically
• different substance AlBr3)

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Units of Measurement SI Unit

• Système International (SI) dUnités
• International agreement on the metric units
  for the
• uses in science (1960)

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Units of Measurement Prefixes
? Prefixes They are used to indicate decimal
fractions or multiples of various units.
A Megabyte of memory 106 bytes of
memory Femtochemistry chemistry that occurs on
the time scale of 10-15 second check out
http//www.lms.caltech.edu (prof. Zewails
homepage)
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Length and Mass
Length 1 meter (m) 100 cm Mass 1
kilogram (kg) 1000 g
Metric to English conversion 1 m
1.093613 yard 1 cm 0.393701 inch 1 kg
2.204623 lb
Check out http//www.digitaldutch.com/unitconverte
r/
NOTE Mass and weight are not the same thing.
Mass is an intrinsic property of
matter, but weight depends on the gravity.
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Temperature
Water
freezing Water boiling Celsius scale (oC)
0 100
Fahrenheit scale (oF) 32
212
oC 5/9 (oF ? 32) oF 9/5(oC) 32
Kelvin K oC 273.15 (exact)
Absolute zero temperature 0 K ? 273.15
oC The lowest attainable temperature in our
universe
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Temperature
William Thomson Kelvin (1824-1907)
On an Absolute Thermometric Scale
Philosophical Magazine, vol. 1 pp. 100-106 (1848)
(98.6 oF ? 32)?5/9 37 oC 37 oC 273.15
310.15 K
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Derived Units
Use the defining equation for the quantity of
interest and substitute the appropriate SI units
? Volume a?b?c (length)3 m3
a
b
c
In chemistry, we normally use smaller units.

• Liter (10 cm)3 1 L 1 dm3 10-3 m3
• 1 gal 3.8 L

(2) Milliliter 1 mL 10-3 L 1 cm3 1 cc
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Derived Units

• Density The amount of mass in a unit volume of
  
• substance

SI unit of density
? In chemistry, we typically use g/mL g/cm3
g/cc
? Density depends on temperature
? Dont be confused about density and weight
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Density, Volume and Mass
(1) 1.00 ? 102 g of mercury occupies a volume of
7.36 cm3. What is the density of mercury?
(2) The density of liquid methanol is 0.791 g/mL.
What is the volume of 65.0 g of liquid
methanol?
(3) The density of gold is 19.32 g/cm3. What is
the mass in gram of a cube of gold if the
length of the cube is 2.00 cm?
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Uncertainty in Measurement
We need to distinguish two different types
of number in science
? Exact Number Defined number
1 dozen 12, 1 m 100 cm
Counted number
There are 120 students in the
class.
? Inexact Number Numbers from measurement
(human errors, machine
errors..)
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Precision and Accuracy
? Precision How closely individual measurements
agree with one another.
? Accuracy How closely individual measurements
agree with the correct or
true value.
good precision poor accuracy
good precision good accuracy
poor precision poor accuracy
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Significant Figures

• Measured quantities are generally reported in
  such a
• way that only the last digit is uncertain.
• mass of a dime 2.2405 g

Uncertain. Could be 6 or 4
(2) Sometimes, ? sign is used to specify the
uncertainty. mass of a dime
2.2305 ? 0.0002 g
Significant Figures All digits of a measured
quantity, including the uncertain one.
2.2405 g ? 5 significant figures
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Rules for Significant Figures

• All non-zero digits are significant

(2) Zeros at the beginning of a number are never
significant ? count the digits starting with
the first non-zero digit 0.0026 has
TWO significant figures
(3) Zeros between non-zero digits are
significant ? 0.00206 has THREE significant
figures
(4) Zeros at the end of a number are
significant. ? 0.002060 has FOUR significant
figures 2060 has FOUR significant
figures
2.06 x 103 has THREE significant figures
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Significant Figures in Calculation
The number with the fewest number of significant
figures limits the certainty of the calculated
quantity.

• Multiplication Division The final answer can
  have
• no more significant figures than the fewest
  number of
• significant figures in any number in the
  problem.

? Addition Subtraction The final answer can
have no more decimal places than the fewest
number of decimal places in any number in the
problem
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Significant Figures in Calculation
Example 1 Area of a rectangle whose measured
edge lengths are 6.221 cm and
5.2 cm
Area (6.221 cm) x (5.2 cm) 32.3492 cm2
Include only 2 significant figures
Only 2 significant figures
Example 2 Addition of three measured numbers
20.42 1.322 83.1
104.842 ?
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Significant Figures in Calculation

• When calculation involves multiple steps
• Retain at least one more extra digit (past the
  number
• of significant figure) in each step

? When you use a calculator Enter the numbers
one after another (without worrying about
significant figures) and rounding only the
final answer
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Significant Figures in Calculation
Example 3 863 ?1255 ? (3.45 ? 108)
863 ? 1255 ? 372.6 863 ? 882.4 761511.2
Example 4 (0.0045 ? 20000.0) (2813 ? 12)
90.0 33800 33890
From calculator 33846
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Dimensional Analysis
We carry units through all calculations. Units
behave like numbers they are multiplied
together, divided into each other, or canceled.
Example How many inches are in 10 cm?
Correct
Wrong

• Advantages of dimensional analysis
• (1) It ensures that your answer has the correct
  unit
• (2) It makes it easier to find out possible
  errors

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Unit Conversion
Conversion factor
Example The speed of N2 in air at 25 oC is 515
m/s. Convert the speed into
mile/hour
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Unit Conversion
Example The density of water is 1.00 g/mL.
What is the mass 1.00 gal of water in
grams?
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An example
The density of gold is 19.32 g/cm3. If 2.00 g
of gold wire has 0.12 mm radius, how long
the wire is?