Chemistry
Fall 2003
Dr Supplee
Chapter 1- Definitions
• Science
– Methodical exploration of nature followed by a
logical explanation of observations
• Scientific Method
– A systematic investigation of nature and
requires proposing an explanation for the
results of an experiment in the form of a
general principle (hypothesis)
Chapter 1 - Definitions
• Hypothesis
– Initial explanation of observations
• Theory
– Sufficient evidence in support of the hypothesis
– Model that scientifically explains the behavior of
nature
• Law
– Does not explain behavior
– States a measurable relationship under different
experimental conditions
Chapter 1 – Definition Examples
• Hypothesis
– Dalton proposed that all matter was composed of
small individual particles (atoms)
• Theory
– 100 years later
Atomic Theory which
explains the composition of substances as well as the
behavior of gases
• Law
– Boyle’s Law P1V1= P2V2 at constant temperature
• If volume decreases than pressure increases at constant
temperature
Chapter 1- Definitions Summary
Scientific Theory
Natural Law
Analyze more data
Hypothesis
Analyze initial observations
Experiment
Chapter 1 – Modern Chemistry
• Organic Chemistry
– Chemistry of carbon containing compounds
• (C, H, O, and N)
• Inorganic Chemistry
– Chemistry of all other substances
• Biochemistry
– Chemistry of substances derived from plant
substances
Chapter 1 – Modern Chemistry
• All three have in common
– Analytical Chemistry
• Qualitative (what) and quantitative (how much)
analyses
– Physical Chemistry
• Theoretical and mathematical explanations of
chemical behavior
Relevance to daily life
Interesting topics
Fun experiments
CHEMISTRY
Career Opportunities
Benefits to society
Applications
Chapter 2- Scientific Measurements
• Introduction to Laboratory
– Work alone
– Handout
– Due 9/15/03
• Measurement Uncertainty
– Plus/minus factor ( error)
• Metric versus English Units
– Conversion factors
• Significant Figures
– Rounding rules
Precision versus Accuracy
Precise, not accurate
True Value
Accurate, not precise
Precision –how close two measurements
of the same quantity are to each other
Accuracy – how close an experimental
observations to the true value
Chapter 2- Scientific Measurements
• Measurement
– a number with units
• Uncertainty
– the estimated unit amount
– plus/minus associated with measurement
• Mass
– Amount of matter an object possesses
• Weight
– Force exerted by gravity on an object
Chapter 2- Scientific Measurements
• Volume
– Amount of space occupied by a solid, gas or
liquid
Significant Digits/ Figures
• Digits are significant when the do more
than hold a decimal place
– A place holder zero is NEVER significant
• determines measurement uncertainty
(error analysis)
• Does not apply for exact numbers, only
measured numbers
Significant Digits/ Figures Rule
• Rule #1
– Count the number of nonzero digits left to
right
– Do not count place holder zeros
Significant Figure Rounding Rules
• After all calculations are complete
determine significant figures and then
round
– 5 or greater round-up to the nearest whole
number
– less than 5 truncate
Scientific Notation
• Exponential numbers (power of 10)
Base
•
•
•
•
10exponent
The number 10 is raised to the nth power
Numbers greater than 1 the exponent is positive
Numbers less than 1 the exponent is negative
The decimal is placed after the first significant
digit and sets the size of the number by using a
power of 10.
Unit Equations, Factors and
Conversions
• Problem Solving Technique
• Equivalent relationships
• Unit equation
– A simple statement of two equivalent
quantities
• Unit Factor
– A ratio of two equivalent quantities
Unit Dimensional Analysis Problem
Solving
• Three steps
1) write down the units asked for in the
answer
2) write down the value given in the problem that is
related to the required answer
3) Apply a unit factor to convert the units in the given
value to the units in the answer
Given Value
x
Unit
Factor
= units asked for
Percent Concept
• amount of a single quantity compared to
the entire sample
• one part per 100 parts
one quantity
total sample
x 100
= %
Review
Significant Digits/ Figures
• Digits are significant when the do more
than hold a decimal place
– If the number is less than 1, a place holder
zero is NEVER significant
• determines measurement uncertainty
(error analysis)
• Does not apply for exact numbers, only
measured numbers
Exact Numbers
• Infinite significant figures
• English to English conversion factors
• Metric to metric conversion factors
Unit Equations, Factors and
Conversions
• Problem Solving Technique
• Equivalent relationships
• Unit equation
– A simple statement of two equivalent
quantities
• Unit Factor
– A ratio of two equivalent quantities
Chapter 3 – The Metric System
• Single basic unit for each quantity
measured
• Decimal system that uses a system of
prefixes to enlarge or reduce a basic unit
Metric System Definitions
• Meter equals one ten-millionth of the
distance from the North Pole to the
equator
• Kilogram equals the mass of one a cube of
water one-tenth of a meter on a side
• Liter equals the volume occupied by a
kilogram of water at 4 oC
The Metric System
Physical Quantity
Basic Unit
Symbol
Length
meter
m
Mass
gram
g
Volume
liter
L
time
second
s
Metric Prefixes
Prefix
gigamegakilodecicentimillimicronano-
Symbol
G
M
k
d
c
m
m
n
Multiple/Fraction
1 x 109
1 x 106
1 x 103
1 x 10-1
1 x 10-2
1 x 10-3
1 x 10-6
1 x 10-9
Metric Conversion Factors Practice
• 1 kg =? g
k = kilo = 1000 basic units
1kg = 1000g
• 2s =? ns
n=nano=1 1 x 10-9
2s=2 x 10-9 ns
Unit Conversion Factors
• Ratio of two equivalent quantities
• The quantity in the numerator is equal to
the quantity in the denominator
• If 100cm = 1 m, then the factor becomes
100 cm
1m
or
1m
100 cm
Metric- English Conversions
Physical
Quantity
English Unit
Metric
Equivalent
length
1 inch (in.)
1 in = 2.54 cm
mass
1 pound (lb)
1 lb = 454 g
volume
1 quart (qt)
1 qt = 946 mL
time
1 second (sec)
1 sec = 1.00 s
Unit Analysis
• Recall:
– Problem Solving Technique
Units Given
Unit
Factor
New unit
Unit
Factor
Units asked for
Practice Problems
• Work in groups of 3-4
• One student from each group puts solution
in board and explains to class
Quiz # 4
• See Chemistry Current News Slides
• Presentation to be given on Oct 6, 2003.
Density - Review
•
•
•
•
•
•
•
•
Lab Experiment 2
Physical property
Defined as mass per unit volume
Liquids and solids expressed in g/ml (g/ cm3)
Gases expresses in grams per liter
Density of water is 1.00 g/ml
Floats in water
density <1.00 g/ml
Sinks in water
density >1.00g/ml
Estimating Density
(page 59 and 60 )
Water, chloroform and ethyl ether are poured into a tall glass cylinder.
Three known solids are added. Identify the liquids.
Solid 1 = ice
Solid 2 = rubber
Solid 3 = aluminum
Liquid 1
Liquid 2 = water
Liquid 3
Temperature
Fahrenheit, Celsius and Kelvin
• Measure of the average energy of
individual particles in a system
– Warmer temps = more molecules moving thus
more energy
– Cooler temps = slow moving molecules thus
less energy
• Fahrenheit
• Celsius
• Kelvin
oF
oC
K
Temperature
• oF
– Freezing point of water 32 oF
– Boiling point of water 212 oF
• oC
– Freezing point of water 0 oC
– Boiling point of water 100 oC
• K ( SI unit)
– Absolute zero
– Equal to -273.15oC
0K
Temperature Conversions
• oF to oC
( oF - 32 oF ) x 100 oC / 180 oF = oC
• oC to oF
( oC x 180 oF / 100 oC ) +32 = oF
• Kelvin
oC
+273
Heat
•
•
•
•
Heat measures the total energy
Temperature measures the average energy
Heat energy units calories or kilocalories
A calorie (cal) is defined as the amount of heat
needed to raise 1 g of water 1 oC
• Food Calorie equals 1 kcal = 1000 cal
• SI unit = joule (J)
1 cal = 4.184 J
Specific Heat
• Amount of heat required to bring about a
given change in temperature
• Observed amount
• Unique for each substance
• Specific heat of water is high
– Change in temperature is minimal as water
gains or losses heat
– Surface of earth is covered in water so water
helps to regulate the climates
Specific Heat
• Amount of heat required to raise the
temperature of 1 g of substance 1 oC
• Units are cal/g oC
Water
1g
1.0 oC
Ice
1g
2.0 oC
Iron
1g
9.3 oC
Silver
1g
17.7 oC
Specific Heat
• gain or loss of heat divided by mass and
temperature change = specific heat
How many calories are required to raise the
temperature of a 3 inch iron nail weighing 7.05 g
form room temperature to 100 oC?
The specific heat of iron is 0.108 cal/g oC
Solution
• Specific Heat = Heat/ (mass x D t)
cal/g oC
= cal / g
x oC
• 0.108 cal/g oC = cal / 7.05 x (100-25oC)
• Solving for Heat ( energy required )
• Rearrange
(0.108 cal/g oC) x 7.05 g x 75 oC
= 57 cal
Chapter 4
Matter and Energy
• Matter is any substance that has mass
and occupies volume
• Physical State changes
– Melting
– Sublimation
– Condensation
– Deposition
– Freezes
– Vaporization
solid into liquid
solid into gas
gas into liquid
gas to solid
liquid to solid
liquid to gas
Increasing temperature
ice
steam
water
melting
vaporizing
freezing
condensing
Sublimation
Deposition
Chapter 4
Matter and Energy
Property
Solid
Liquid
Gas
Shape
Definite
Indefinite
Indefinite
Volume
Fixed
Fixed
Variable
Compressibility
negligible
negligible
significant
Elements, Compounds and Mixtures
• Properties of matter may be consistent
throughout or they may vary
• Melting point
– Gold (Au)
– Quartz
1064 oC
1000 – 1600 oC
• Gold is homogenous – properties consistent
• Quartz is heterogeneous – properties vary
Mixtures
• Heterogeneous
– Usually Solids
– Separated into pure substances by physical methods
which take advantage of different physical properties
– Properties are not the same throughout the sample
• Homogeneous
– Gases or liquids
– Separated into pure substances by either chemical or
physical methods which take advantage of different
physical properties
– Properties a the same for any given sample, but can
vary sample to sample
Mixtures
• Alloy
– Homogeneous mixture of two or more metals
– Gold ( Au)
10 K
42 %
14 K
18 K
75%
• Substance
– Matter with definite composition and constant properties
– Compound or an element
• Compound
– Broken down into elements by chemical reactions
• Element
– Cannot be broken down further by chemical reactions
Matter
Mixtures
Heterogeneous
Physical Separate
Homogeneous
Substances
Compounds
Elements
Names and Symbols of the Elements
• 81 stable elements that
occur in nature
• Only 10 account for 95%
of the mass of the earths
crust, water and
atmosphere
Element
Mass
Mass
Element
Percent
Percent
O
49.5
Na
2.6
Si
25.7
K
2.4
Al
7.5
Mg
1.9
Fe
4.7
H
0.9
Ca
3.4
Ti
0.6
All other elements 0.5 %
Names and Symbols
• Names are from various sources
– Hydrogen
– Carbon
– Calcium
(hydro, Gr. = water former)
(carbo, Lt. =coal)
(calcis, Lt. = lime)
• Chemical Symbols
– Dalton in 1803 proposed that elements are composed
of indivisible spherical particles or atoms (atomos, Gr.
= indivisible)
– Suggested the use of circles with markings for
symbols ( pg 83)
– Berzelius in 1813 proposed our current system of
symbols = using the first letter of the name and if the
first letter is already in use two letters
Metals, Nonmetals and Semimetals
• Predict by position in Periodic Table
• Metals
– solid element
– Bright metallic luster
– Good conductor of heat and electricity
– High density
– High melting point
– Malleable ( thin sheets)
– Ductile ( fine wire)
Metals, Nonmetals and Semimetals
• Nonmetals
–
–
–
–
–
–
Solid or gas element
Dull appearance
Low density
Low melting point
Poor conductor of heat and electricity
Crush to a powder, if solid
• Semimetal
– metalloids
– midway between metal and nonmetal
– Semiconductor
Periodic Table of the Elements
(page 86)
• Atomic number
– Number of protons
•
•
•
•
•
Metals are placed on the left
Nonmetals on the right
Separated by semimetals starting at B
Solids are to the left (most all elements)
Gases to the right
Compounds & Chemical Formulas
• 1799 Proust
• Law of Definite Composition
• Law of Constant Proportion
“Compounds always contain the same
elements in a constant proportion by
mass.”
Chemical Formulas
• Most elements occur in nature as collection of individual
atoms
• Diatomic molecules
– Oxygen (O2), Hydrogen (H2), Nitrogen (N2)
– Halogens
Chlorine (Cl2),Bromine Br2, Iodine (I2)
• A chemical formula expresses the number and type of
each atom in a compound
• The number of the each atom is indicated with a
subscript. The number 1 in the subscript is implied and
therefore is omitted.
• Parentheses are used to help clarify the structure of the
compound and
Examples
• Water
H2O
– 2 hydrogen atoms and 1 oxygen atom
• Calcium Chloride
CaCl2
– 1 calcium atom and 2 chlorine atoms
• Propylene Glycol
C3H8O2
– 3 carbon atoms, 8 hydrogen atoms, 2 oxygen atoms
• Lead acetate
PbC2H3O2
– 1 lead atom, 2 carbon atoms, 2 oxygen atoms, 3
hydrogen atoms
• 4-amino-2-hydroxytoluene
C7H9NO
– 7 carbon atoms, 9 hydrogen, 1 nitrogen, 1 oxygen
Physical and Chemical Properties
• Substances ( Compounds or Elements)
– Physical and chemical properties are the
consistent throughout
– No two substances have all the same physical
and chemical properties
2 Na +
Cl2
2 NaCl
metal yellow gas
white solid
Physical and Chemical Properties
• Physical properties are measured and observed
without changing the chemical composition of
the substance
• Examples
– Appearance, color, melting point, density, solubility,
boiling point, freezing point
• Chemical properties describes how a
substances reacts with other substances
(reaction chemistry) and always involve a
chemical change
Periodic Table and Reaction
Chemistry
• Elements with similar reaction chemistry
are “Grouped” into families
– Columns in the Periodic table
•
•
•
•
•
•
•
Group IA
Group IIA
Group VIIA
Group VIIIA
Group IIlB to IIB
Lanthanides
Actinides
Alkali Metals
Alkaline Earth Metals
Halogens
Noble Gases
Transition Metals
Physical and Chemical Changes
• Physical Change – chemical composition
does not change, but the physical state
does
• Example ?
• Chemical Change – chemical composition
changes and the physical state may or
may not change ( formation of a new
substance)
• Example ?
Physical and Chemical Changes
• Chemical changes are often detected by
– Gas formation
– Color change
– Release of energy
– Precipitate formation
bubbles or odor
permanent
light or heat
solids
Conservation of Mass and Energy
• Matter is nether created or destroyed
during a chemical reaction
• Energy can not be created or destroyed. It
can however, be converted from one form
to another
• Total mass and energy of the universe is
constant
Potential vs. Kinetic Energy
• Key concepts for understanding chemical
reactions
• Potential energy is stored energy of matter
at rest
• Kinetic energy is energy that is a result of
motion
• Potential energy can be (and is) converted
into kinetic energy
Kinetic Energy as a function of Physical State
Particles
Solid
Liquid
Gas
Kinetic
Energy
Very low
High
Very high
Movement
None
Restricted unrestricted
Chemical Reactions
• Potential Chemical Energy
• Kinetic Heat Energy
• Exothermic reactions
– Reactions that give off or produce heat
– Reactants have more potential energy than products
• Endothermic reactions
– Reactions that take in or absorb heat
– Reactants have less potential energy than products
Example
Exothermic
Reactants
( high P.E.)
Endothermic
Reactants + heat
( low P.E.) + (K.E.)
Products + heat
(low P.E. ) ( K.E.)
Products
(high P.E. )
Forms of Energy
• Six Forms
– Light
– Heat
– Chemical
– Electrical
– Mechanical
– Nuclear
Energy Examples
• Radioactive Ur vaporizing water
nuclear
heat
• Steam driving a turbine
heat
mechanical
• Lead acid battery
chemical
electrical
Chapter 5 – Models of the Atom
• Atom – indivisible
• Dalton
– Proposed that all matter was composed as
tiny particles
– Based on the Laws of Conservation of Mass
and Definite Proportion
– Compounds are simply the combination of
two or more atoms of different elements
Dalton’s Atomic Theory
1. An element is composed of tiny, indivisible
particles called atoms
2. All atoms of an element are identical and have
the same properties
3. Atoms of different elements combine to form
compounds
4. Compounds contain atoms in small whole
number ratios
5. Atoms can combine in more than one ratio to
form different compounds
Thomson Model
(Plum Pudding Model)
• Cathode Ray experiment
– Glass tubes containing a low pressure amount of gas emitted
light when electricity was applied to one end of the tube
(Fluorescence – light energy)
– Ray emanates form the negative cathode in the tube, the
radiation is referred to as a cathode ray
– Placed tube in magnetic field, light or ray curved towards the
positive
– Concluded that cathode rays were composed of tiny negatively
charged particles ( electrons, e- )
– Further experiments showed that certain rays contained small
particles that had an equal but opposite in sign charge to
electrons
protons (p+)
Relative Charges and Masses
Subatomic
Particle
Symbol
Relative
Charge
Relative
Mass
Electron
e-
-1
1/1836
Proton
p+
+1
1
Thomson Model of the Atom
Homogeneous sphere
“plum pudding”
+
+
-
-
-
+
-
+
-
+
-
-
Rutherford Model of the Atom
• Alpha ray –
– particles identical to He w/o electrons ( He +2 )
– most passed through thin Au foil
– suggested that most of the atoms were empty space,
with electrons moving about a center ( nucleus)
– Nucleus contains protons and is tiny and dense
• Beta ray • Gamma ray – not affected by magnetic fields
Rutherford Model of the Atom
Neutrons
“heaviness of the atom”
--
1 X 10
-8
++ 0
cm +++ n
1 x 10 -13 cm
----
Subatomic Particles
Relative
Location
Charge
Relative
Mass
e-
Outside
nucleus
-1
1/1836
Proton
p+
Inside
nucleus
+1
1
Neutron
n0
Inside
nucleus
0
1
Particle
Electron
Symbol
Atomic Notation
Mass Number
(protons and neutrons)
Atomic Number
( protons)
A
Z
Sy
symbol
Isotopes
• Same element with different amount of
neutrons
– Number of Protons and Electrons are the
same
• Mass of the element is different, so the
mass number is different
• For example, Hydrogen and deuterium
1
1
H
2
1
H
Isotopes
• Naming
– Some are common (hydrogen, deuterium, tritium)
– Name of the element followed by the mass number
– For example,
carbon -14
oxygen -18
cobalt-60
Hint: should be able to determine number of protons
electrons, neutrons
atomic notation, isotope
names
Atomic Mass
•
•
•
•
•
•
Atoms are to small to weigh on balance
Masses are relative to each other
Specifically, mass is relative to carbon-12
Carbon-12 has 12 amu (atomic mass unit)
So, an amu = 1/12 the mass of carbon
Weighted average
Atomic Mass
• Weighted average of all the naturally
occurring isotopes
• Given the natural abundance of the atom,
the amu can be calculated for a given
atom
• Since it’s a calculated value no atom will
actually weigh this number
Periodic Table
• Does not tell about the number of naturally
occurring isotopes
• Atomic Number = Protons
• Atomic Mass = weighted average = protons and
neutrons (whole numbers)
• Mass number is in parentheses, then the
element is unstable. Mass is given for the best
known isotope.
Hint: Should be able to tell form periodic table which
elements are stable and which are not.
Wave Nature of Light
• Wavelength – the distance the light wave travels
to complete one cycle
• Frequency - the number of wave cycles
completed in 1 s
• Velocity of light is constant = 3.00 x 103 m/s
• If wavelength decreases, frequency increases
• Low frequency – low light energy –long
wavelengths
• High frequency-high energy light –short
wavelengths ( page 124)
Light- A Continuous Spectrum
• White light passes through a prism it
separates into all the colors
– ROY G BIV
• Light – radiant energy that is visible
• Visible spectrum = 400-700 nm
• Radiant energy spectrum – a continuous
spectrum of visible and invisible light that
ranges form short to long wave lengths
Radiant Energy Spectrum
Wave Length Increases
Cosmic Rays
Gamma Rays
UV / VIS
X Rays
400 nm
Infrared
Microwaves
500 nm
Visible Spectrum
TV
600 nm
Radio
700 nm