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Unit 1: Lecture Notes
Unit 1: Atomic and Molecular Structure
L. G.
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Topic
Intro to matter: Definition of Matter
Comparing Elements and Compounds
Elements
Element Symbols
Elements: The first 20
States of matter
Physical and Chemical Properties
Elements: The second 20
Mixtures and Pure Substances
Separation of mixtures
Periodic Table: Introduction
Elements: The Remaining
Metals, Nonmetals, and Metalloids
Major Groups
Periodicitiy (Moseley's Periodic Law)
Subatomic Particles
Dalton's Atomic Model
Thomson's model (Discovery of electrons)
Rutherford (Discovery of nucleus)
Bohr model: Introduction
Wave model: Introduction
Standard model
Fundamental particles
Atomic Structure: Atomic number and mass
Isotopes
Light Properties & Waves
Electromagnetic Spectrum
Photoelectric effect
Emission Spectra
Orbitals (Electrons as waves)
Quantum numbers
Aufbau principle, Pauli exclusion, Hund's rule
Electron configurations
Valence Electrons
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Text
2.1
2.2
3.1
3.2
Suppl
2.3
2.4
Suppl
2.5
2.6
3.8
Suppl
11.11
Suppl
Suppl
3.6
3.3
3.5
3.5, 11.1
11.5
11.6
Suppl
Suppl
3.7
3.7
11.2-11.6
11.2
Suppl
11.3
11.7-11.8
Suppl
Suppl
11.9-11.10
Suppl
Lab
Mixtures
Chromatography
Flame test
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Unit 1: Lecture Notes
Standard 10
15
16
56
57
58
59
60
61
62
63
Nuclear Radiation
Half-Life
Nuclear Reactions
Nuclear Decay
Balancing Nuclear Reactions
Fission
Fusion
Effects of radiation
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Chp 19 Intro
19.3
19.1
19.1, 19.2 The Alchemical race
Suppl
19.6-19.7
19.9
19.1
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Unit 1: Lecture Notes
LESSON 22: INTRODUCTION TO MATT ER
2.1 The Particulate Nature of Matter
1. Matter
a. Definition
i. Matter takes up space (volume) and has mass. Matter is loosely defined in
order to provide room for explaining matter’s behavior at the “quantum”
(extremely small) level.
2. Atoms
a. All matter consists of tiny particles called atoms (p. 22).
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Unit 1: Lecture Notes
LESSON 23: COMPARING ELEMENTS A ND COMPOUNDS
2.2 Elements and Compounds
1. Elements
a.
b.
c.
d.
contain only one kind of atom
atoms with the same number of protons
do not decompose into other elements
properties do not vary
2. Compounds / Molecules
a. elements combined chemically in law of definite proportions
b. properties do not vary
A molecule is formed when two or more atoms join together chemically
Compounds always have different elements. Molecules may or may not have
different elements.
ALL COMPOUNDS ARE MOLECULES BUT NOT ALL MOLECULES ARE COMPOUNDS.
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Unit 1: Lecture Notes
LECTURE NOTES
2. Compounds / Molecules
Elements combined chemically in law of definite proportions
Water is always H2O  Two hydrogen atoms and one oxygen atom
Properties do not vary
B. Mixtures
1. Definition
A substance consisting of two or more substances placed together but not in
fixed proportions and not with chemical bonding.
2. Homogeneous
Mixtures without distinguishing parts
Sodas, milk, solder
Salt water, air, brass
3. Heterogeneous
Mixtures with distinguishing parts
Cement, iron filings in sand
Sand and water, gravel
Think of a pizza 
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Unit 1: Lecture Notes
LESSON 24: THE ELEMENTS
Text 3.1
A. History
a. Greeks proposed that all matter was composed of four fundamental substances:
fire, earth, water, and air.
b. Alchemists discovered mercury, sulfur, and antimony.
c. Robert Boyle insisted science should be grounded in experiments; defined an
element as a substance that could not be broken down.
1. 88 elements occur naturally; other elements are man-made (made in particle
accelerators)
2. Living systems
a. Elements in living systems: CHNOPS
b. Trace elements in living systems
i. Elements in small amounts
3. Multiple meanings of elements
a. Single atom
b. Large sample enough to weigh on a balance
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Unit 1: Lecture Notes
LESSON 25: SYMBOLS FOR THE ELEM ENTS
Text 3.2
A. Names
1. Some element names are derived from Latin, Greek, or German describing the
element’s property
a. Gold (Au) from aurum, Latin from “shining dawn”
b. Lead (Pb) from plumbum, meaning “heavy”
2. Some element names are derived from a person or place
B. Symbols (abbreviations of chemical elements)
1. Rules must be followed
a. First letter is capitalized
b. Second letter must be lowercase
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Unit 1: Lecture Notes
LESSON 26: ELEMENTS – THE FIRST 20
Students should memorize the first twenty element names and symbols.
Happy Henry, the Little Beach Boy, CaN dO FiNe
Naughty Megan, the Alpine Sister, Pretends to Ski at ClArK Canyon
Element Name
Symbol
Element Name
Symbol
Hydrogen
H
Sodium
Na
Helium
He
Magnesium
Mg
Lithium
Li
Aluminum
Al
Beryllium
Be
Silicon
Si
Boron
B
Phosphorous
P
Carbon
C
Sulfur
S
Nitrogen
N
Chlorine
Cl
Oxygen
O
Argon
Ar
Fluorine
F
Potassium
K
Neon
Ne
Calcium
Ca
Practice writing the names and symbols. Make and use flash cards to help.
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Unit 1: Lecture Notes
LESSON 27: STATES OF MATTER
2.3 States of Matter
1. Described macroscopic and microscopic views
a. Macroscopic
Solid
Liquid
Gas
Shape
Fixed
Of container
No
Volume
Fixed
Fixed
No
3.
[Image from http://www.grc.nasa.gov/WWW/K-12/airplane/state.html]
Microscopic
Particles in a:
□
□
□
Gases are well separated with no regular arrangement.
Liquids are close together with no regular arrangement.
Solids are tightly packed, usually in a regular pattern.
Particles in a:
□
□
□
Gases vibrate and move freely at high speeds.
Liquids vibrate, move about, and slide past each other.
Solids vibrate (jiggle) but generally do not move from place to place.
Liquids and solids are also known as condensed phases.
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Unit 1: Lecture Notes
LESSON 28: PHYSICAL AND CHEMICAL PROPERTIES
2.4 Physical and Chemical Properties and Change
2. Properties of Matter
Properties are characteristics of matter.
A. Definitions
1. Physical Properties
Can be observed with the senses and determined without destroying the object
2. Chemical Properties
A substance reacts with something else, creating a new substance. Heat changes
usually accompany chemical changes.
B. Substances, or matter, have two types of properties (physical/chemical).
1. Consider gold (Au):
a. Shiny (lustrous)
b. Conducts heat and electricity
c. Malleable (hammer into thin sheets)
d. Ductile (pull into wire)
Your book refers to odor, color, volume, state (phase), density, etc
C. Physical properties that do not depend on the system size or amount
1. Of material are intensive properties
* Temperature, density, melting or boiling points
2. Properties that depend on the system size are extensive
* Mass, volume
Consider two beakers of water: 100g and 10g of water
The mass and volume are different
Yet, the temperature and density of water are the same.
D. Chemical properties
At this point we’ve been talking about physical properties. Chemical properties that
change the nature of the substance are chemical
* PH, reactivity with other substances
E. Physical changes
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Unit 1: Lecture Notes
The most common physical change is a phase change:
* Solid
liquid
gas.
The molecules, compounds, or elements still exist and nothing new is created.
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Unit 1: Lecture Notes
LESSON 29: ELEMENTS SECOND 20
Students should memorize the second twenty element names and symbols.
Scary Tina a Vicious Crow, was Mean and Fierce and Could Nibble Cuts from Zn.
Gabriel Gemp wAs Seeing his Brother Krimped in Rubber, he said SorrY loZer
Element Name
Symbol
Element Name
Symbol
Scandium
Sc
Gallium
Ga
Titanium
Ti
Germanium
Ge
Vanadium
V
Arsenic
As
Chromium
Cr
Selenium
Se
Manganese
Mn
Bromine
Br
Iron
Fe
Krypton
Kr
Cobalt
Co
Rubidium
Rb
Nickel
Ni
Strontium
Sr
Copper
Cu
Yttrium
Y
Zinc
Zn
Zirconium
Zr
Practice writing the names and symbols. Make and use flash cards to help.
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Unit 1: Lecture Notes
LESSON 30: MIXTURES AND PURE SUBSTANCES
2.5 Mixtures and Pure Substances
I. Classification of Matter
Two general categories of matter: pure substances and mixtures.
A. Pure substances
Definition:
Have uniform composition (cannot be separated by physical means)
Pure water is H2O; pure Gold (Au) consists of just Gold atoms
* Percentage composition are always the same from sample to sample
* Pure substances melt and boil at a characteristic temperature
Two types of pure substances: elements and compounds
1. Elements
If a pure substance cannot be decomposed into something else, then the substance
is an element.
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Unit 1: Lecture Notes
LESSON 31: SEPARATION OF MIXTURES
2.6 Separation of Mixtures
Physical means can be used to separate a mixture into its pure components
A. Distillation
Definition
Distillation is a method of separating mixtures based on differences in
their volatilities in a boiling liquid mixture.
B. Filtration
Definition
Filtration is a mechanical or physical operation which is used for the separation of
solids from fluids (liquids or gases) by interposing a medium through which only the
fluid can pass.
Summary
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Unit 1: Lecture Notes
LESSON 32: INTRODUCTION TO THE PERIODIC TABLE
Text 3.8
Johann Dobereiner (1820s)
John Newlands (1864)
Dimitri Mendeleev (1869)
Henry Mosley (1914)
The Periodic Law
A.
B.
The physical and chemical properties of the elements are periodic functions of
their atomic numbers
Elements on the table are arranged in order of increasing atomic number
(number of protons)
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Unit 1: Lecture Notes
LESSON 33: REMAINING ELEMENTS
Students should memorize the remaining element symbols and names in table 3.3 (p. 51).
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Unit 1: Lecture Notes
LESSON 34: METALS, N ONMETALS, AND METALL OIDS
A.
Metals
1.
2.
3.
4.
5.
Good conductors of heat and electricity
Lustrous (shiny)
Solids (except mercury)
Ductile (can be drawn into wire)
Malleable (can be hammered into thin sheets)
B.
Nonmetals
1. Poor conductors of heat and electricity
2. Most are gaseous
3. Solids tend to be brittle
C.
Metalloids
1. Some properties of metals, some of nonmetals
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Unit 1: Lecture Notes
LESSON 35: MAJOR GRO UPS
Periods and the Blocks of the Periodic Table
A.
Periods
1. Horizontal rows on the periodic table
2. Period number corresponds to the highest principal quantum number of
the elements in the period
B.
Sublevel Blocks
1. Periodic table can be broken into blocks corresponding to s, p, d, f
sublevels
Blocks and Groups
A. s-Block, Groups 1 and 2
1. Group 1 - The alkali metals
a. One s electron in outer shell
b. Soft, silvery metals of low density and low melting points
c. Highly reactive, never found pure in nature
2. Group 2 - The alkaline earth metals
a. Two s electrons in outer shell
b. Denser, harder, stronger, less reactive than Group 1
c. Too reactive to be found pure in nature
B.
d-Block, Groups 3 - 12
1. Metals with typical metallic properties
2. Referred to as "transition" metals
3. Group number = sum of outermost s and d electrons
C.
p-Block elements, Groups 13 - 18
1.
Properties vary greatly
a. Metals
(1) Softer and less dense than d-block metals
(2) Harder and more dense than s-block metals
b.
Metalloids
(1) Brittle solids with some metallic and some nonmetallic
properties
(2) Semiconductors
c.
D.
Nonmetals
(1) Halogens (Group 17) are most reactive of the nonmetals
f-Block, Lanthanides and Actinides
1.
2.
Lanthanides are shiny metals similar in reactivity to the Group 2 metals
Actinides
a. All are radioactive
b. Plutonium (94) through Lawrencium (103) are man-made.
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Unit 1: Lecture Notes
STAR RELEASED CHEMISTRY Q UESTIONS
Iodine would have chemical properties most
like
A manganese (Mn).
B tellurium (Te).
C chlorine (Cl).
D xenon (Xe).
Which of the following ordered pairs of
elements shows an increase in atomic number
but a decrease in average atomic mass?
A Ag to Pd
B Co to Ni
C Ge to Sn
D Cr to Mo
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Unit 1: Lecture Notes
LESSON 36: ATOMIC PROPERTIES AND THE PER IODIC TABLE
Text 11.11
Periodic Trends (Periodicity)
By organizing the elements by atomic number, patterns of physical and chemical
properties are seen.
Atomic Radii (size)
A.
Atomic Radius
1.
together
B.
One half the distance between nuclei of identical atoms that are bonded
Trends
1. Atomic radius tends to decrease across a period due to increasing positive
nuclear charge
2. Atomic radii tend to increase down a group due to increasing number
energy levels (outer electrons are farther from the nucleus)
Trends in Ionization Energy
A.
Ion
1.
An atom or a group of atoms that has a positive or negative charge
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Unit 1: Lecture Notes
B.
Ionization
1.
C.
Any process resulting in the formation of an ion
Ionization Energy
1. The energy required to remove one electron from a neutral atom of an
element, measured in kilojoules/mole (kJ/mol)
A
D.
+
energy 
A
+
e-
Trends
1.
Ionization energy of main-group elements tends to increase across each
period
a. Atoms are getting smaller; electrons are closer to the nucleus
2. Ionization energy of main-group elements tends to decrease as atomic
number increases in a group
a.
Atoms are getting larger; electrons are farther from the nucleus
b. Outer electrons become increasingly more shielded from the
nucleus by inner electrons
E.
3.
Metals have characteristic low ionization energy
4.
Nonmetals have high ionization energy
5.
Noble gases have very high ionization energy
Removing Additional Electrons
Na
+
496 kJ/mol
Na
+
4562 kJ/mol ◊
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◊
Na
Na++
+
e+
e-
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Unit 1: Lecture Notes
Na++
+
6912 kJ/mol ◊
Na+++
+
e-
1.
Ionization energy increases for each successive electron
2.
Each electron removed experiences a stronger effective nuclear charge
3. The greatest increase in ionization energy comes when trying to remove
an electron from a stable, noble gas configuration
Trends in Ionic Size
A.
Cations
1.
Positive ions
2.
Smaller than the corresponding atom
a.
b.
B.
Less shielding of electrons
Anions
1.
Negative ions
2.
Larger than the corresponding atoms
a.
b.
C.
Protons outnumber electrons
Electrons outnumber protons
Greater electron-electron repulsion
Trends
1.
Ion size tends to increase downward within a group
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Unit 1: Lecture Notes
Trends in Electronegativity
A.
Electronegativity
1.
electrons
A measure of the ability of an atom in a chemical compound to attract
2. Elements that do not form compounds are not assigned
electronegativities
B.
Trends
1.
Nonmetals have characteristically high electronegativity
a.
2.
Highest in the upper right corner
Metals have characteristically low electronegativity
a.
Lowest in the lower left corner of the table
3.
Electronegativity tends to increase across a period
4.
Electronegativity tends to decrease down a group of main-group element
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Unit 1: Lecture Notes
LESSON 37: SUBATOMIC PARTICLES
Text 3.6
We’ll work backwards for a moment, accept there is an atom, and remind you of the three
major subatomic particles located in an atom that you learned about in 8th grade.
Symbol
Location
Charge
Relative mass
Actual mass
Electrons
e-
Outside nucleus
Negative
1/1840
9.11 x 10-28
Protons
p+
In nucleus
Positive
1
1.673 x 10-24
Neutrons
n0
In nucleus
Neutral
1
1.675 x 10-24
A. Distinguishing Among Atoms
1. Atomic Number, Mass Number, and Electrons
a. Atomic Number (Z)
i.
The number of protons in the nucleus of each atom of that element
ii.
Atoms are identified by their atomic number
iii.
Because atoms are neutral,
# protons = # electrons
iv.
Periodic Table is in order of increasing atomic number
b. Mass Number (A)
i.
The total number of protons and neutrons in the nucleus of an isotope
c. Electrons
i.
The volume of an atom is from the area in which the electrons move
ii.
The chemical properties of an atom arise from the electrons.
B. Calculating the number of electrons, protons, and neutrons (Introduction)
A = protons + neutrons
Z = protons
Therefore, to get the number of neutrons, subtract A - Z
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Unit 1: Lecture Notes
LESSON 38: DALTON’S ATOMIC THEORY
Text: 3.3
John Dalton (1766 – 1844) explained observations such as the law of constant composition
(a compound always has the same composition) using his atomic theory. The predictive
value of the theory led to its eventual acceptance.
A.
Defining the Atom
1. Atomic Theory
a. All matter is made up of very tiny particles called atoms
b. Atoms of the same element are chemically alike
c. Individual atoms of an element may not all have the same mass. However, the
atoms of an element have a definite average mass that is characteristic of the
element
d. Atoms of different elements have different average masses
e. Atoms are not subdivided, created, or destroyed in chemical reactions
1e. Students know the nucleus of the atom is much smaller than the atom yet contains most of its mass.
The volume of the hydrogen nucleus is about one trillion times less than the volume of the hydrogen atom, yet the nucleus contains almost all the mass in the form of one
proton. The diameter of an atom of any one of the elements is about 10,000 to 100,000 times greater than the diameter of the nucleus. The mass of the atom is densely
packed in the nucleus.
The electrons occupy a large region of space centered around a tiny nucleus, and so it is this region that defines the volume of the atom. If the nucleus (proton) of a hydrogen
atom were as large as the width of a human thumb, the electron would be on the average about one kilometer away in a great expanse of empty space. The electron is
almost 2,000 times lighter than the proton; therefore, the large region of space occupied by the electron contains less than 0.1 percent of the mass of the atom.
2. Sizes of Atoms
a. Atomic radius
i. 40 to 270 picometers (pm)
1. 1 pm = 10-12m
ii. Most of the atomic radius is due to the electron cloud
b. Nuclear radius
i. 0.001 pm
ii. density is 2x108 metric tons/cm3
1. 1 metric ton = 1000kg
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Unit 1: Lecture Notes
B.
Models of the Atom
A model is a representation of nature, an attempt to communicate an explanation.
Scientist
Year
Model
Experiment
Focus
Democritus
~ 400
B.C.E.
None
Suggested Atom
Dalton
1808
Solid sphere, tiny, indivisible,
indestructible particles
Weather data
Thomson
1897
Plum pudding
Cathode Ray Tube;
also invented mass
spectrometer
Electrons
Plank
1900
Energy emitted in discrete
quantities
Radiation from solids
Quanta
Rutherford
1911
Nuclear Atom; also called the
planetary model
Gold foil
Nucleus
Bohr
1913
Bohr Model, electrons travel in
discrete orbits
Spectrum of Hydrogen
Excited and Ground state
Einstein
1905
Wave mechanical model
Photoelectric Effect
Photons
Schrödinger
1926
Wave mechanical model
Schrödinger cat;
thought experiment
Schrödinger equation
Heisenberg
1929
Wave mechanical model
Heisenberg uncertainty
principle
Murray GellMann; George
Zweig
1970s
Standard Model
Quarks and leptons
(matter)
* There were many other models developed during this time period but we’ll only focus on these particular ones.
C.
Contributed to the Models of the Atom
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Unit 1: Lecture Notes
Maxwell
1873
Visible light consists of
electromagnetic waves
Planck
1900
Energy emitted in discrete
quantities
Chadwick
1932
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Provides description of
light
Radiation from solids
Quanta; Plank’s constant
Identified subatomic
particle
Neutron
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Unit 1: Lecture Notes
LESSON 39: THE STRUCTURE OF THE ATOM / THOMSON’S MODEL
Text: 3.5
A. The Electron
a. Discovery
i. Joseph John Thomson (1897)
1. Cathode ray tube produces a ray with a constant charge to mass
ratio
2. All cathode rays are composed of identical negatively charged
particles (electrons)
B. Plum-pudding model
C. Inferences from the properties of electrons
i. Atoms are neutral, so there must be positive charges to balance the
negatives
ii. Electrons have little mass, so atoms must contain other particles that
account for most of the mass
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Unit 1: Lecture Notes
LESSON 40: RUTHERFOR D’S MODEL
Text 11.1
1a. Students know how to relate the position of an element in the periodic table to its atomic number and atomic mass.
An atom consists of a nucleus made of protons and neutrons that is orbited by electrons. The number of protons, not electrons or neutrons, determines the unique properties of
an element. This number of protons is called the atomic number. Elements are arranged on the periodic table in order of increasing atomic number. Historically, elements were
ordered by atomic mass, but now scientists know that this order would lead to misplaced elements (e.g., tellurium and iodine) because differences in the number of neutrons for
isotopes of the same element affect the atomic mass but do not change the identity of the element.
D. Structure of the Nucleus
1. The Nucleus
a. The Rutherford Experiment (1911)
b. Alpha particles (helium nuclei) fired at a thin sheet of gold
i. Assumed that the positively charged particles were bounced back if they
approached a positively charged atomic nucleus head-on (Like charges
repel one another)
Results from gold foil experiment
1. Very few particles were greatly deflected back from the gold sheet
a. nucleus is very small, dense and positively charged
b. most of the atom is empty space
2. Structure of the Nucleus
a. Protons
i.
Positive charge, mass of 1.673x10-27kg
ii.
The number of protons in the nucleus determines the atom's identity
and is called the atomic number (Z)
b. Neutrons
i.
James Chadwick (1932)
ii.
No charge, mass of 1.675x10-27kg
c. Nuclear Forces
i.
Short range attractive forces:
a. neutron-to-neutron, proton-to-proton,
proton-to-neutron
Unanswered Questions
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Unit 1: Lecture Notes
What are the electrons doing? How are the electrons arranged? How do electrons
move? Why aren’t electrons (negatively charged) attracted to the positive nucleus?
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Unit 1: Lecture Notes
LESSON 41: BOHR’S MODEL
11.5 The Bohr Model of the Atom
The energy of the electrons is restricted to certain discrete
values; that is, the energy is quantized.
Consider the rungs of a ladder. There is no “in
between” on a ladder. Your foot is either on a
rung or it is not.
The electrons move from each orbital. Photons (packets
of light) are either absorbed or released. If it is at the lowest it’s called
ground state and the highest is called excited state.
Image from http://imagine.gsfc.nasa.gov/docs/teachers/lessons/xray_spectra/background-atoms.html
The flame test shows the spectra changes based on
the elements.
A.
Electron Orbits or Energy Levels
1.
Electrons can circle the nucleus only in allowed
paths or orbits
2.
The energy of the electron is greater when it is
in orbits farther from the nucleus
3.
The atom achieves the ground state when atoms
occupy the closest possible positions around the nucleus
4.
Electromagnetic radiation is emitted when
electrons move closer to the nucleus (excited to ground)
B.
C.
Energy transitions
1.
Energies of atoms are fixed and definite quantities
2.
Energy transitions occur in jumps of discrete amounts of energy
3.
Electrons only lose energy when they move to a lower energy state
Shortcomings of the Bohr Model
1.
Doesn't work for atoms larger than hydrogen (more than one electron)
2.
Doesn't explain chemical behavior
3.
Electrons do not move in a circular motion.
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Unit 1: Lecture Notes
LESSON 42: WAVE MODEL INTRODUCTION
Text 11.6
The Bohr model explained the hydrogen spectrum very well but it failed to explain the
spectra of all other atoms. Additional spectra analysis of elements supported a new model of
the atom, called the wave model. In this model, orbits do not exist. Instead, orbitals that
match spectra are discussed. These orbitals are s (sharp), p (principal), d (diffuse), and f
(fundamental). Orbitals do not describe the path or motion of the electron. Instead, they
describe the probability of finding an electron at a particular time.
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Unit 1: Lecture Notes
LESSON 43: STANDARD MODEL
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Unit 1: Lecture Notes
LESSON 44: FUNDAMENT AL PARTICLES
Text 3.6 and Supplement
Simplest view of the atom is that there is a nucleus (10-13 cm) and electrons that move
outside the nucleus about 10-8 cm from it.
Particle
Symbols
Relative charge
Mass number
Electron
e-
-1
0
Proton
P+
+1
1
Neutron
N0
0
1
As tools improved, chemists and physicists discovered more and more subatomic particles.
By the end of the 1950s there were hundreds of known subatomic particles. This led to a
classification of elementary or fundamental particles; that is, particles that make up the
subatomic particles. See standard model (lesson 43).
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Unit 1: Lecture Notes
LESSON 45: ATOMIC MASS AND NUMBER
Text 3.6
2.
Atomic Number, Mass Number, and Isotopes
d. Atomic Number (Z)
v.
The number of protons in the nucleus of each atom of that element
vi.
Atoms are identified by their atomic number
vii.
Because atoms are neutral,
# protons = # electrons
viii.
Periodic Table is in order of increasing atomic number
e. Mass Number
ii.
The total number of protons and neutrons in the nucleus of an isotope
f. Electrons
iii.
The volume of an atom is from the area in which the electrons move
iv.
The chemical properties of an atom arise from the electrons.
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LESSON 46: ISOTOPES
3.7 Isotopes
1.
Defining Isotopes
i.
Atoms of the same element that have
different masses
ii.
All elements of the same element have
the same # of protons, but may vary in
the number of neutrons
iii.
Although isotopes have different masses,
they do not differ significantly in their
chemical behavior
iv.
Hydrogen as an example:
Image from http://encarta.msn.com/media_461531710/hydrogen_isotopes.html
2.
Designating Isotopes
a. Hyphen notation
i. Mass number is written after the name of the element
1. Hydrogen-2
2. Helium-4
b. Nuclear Symbol
i. Composition of the nucleus using the element’s symbol
ii. For example,
2
1
H
4
2
Name
Mass number = 2 Atomic number = 1
He Mass number = 4 Atomic number = 2
Symbol
sodium
nitrogen
15
7
Atomic number
Mass number
11
23
5
11
Neutrons
N
136
56
Ba
lithium
boron
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Unit 1: Lecture Notes
LECTURE NOTES
11.6 The Wave mechanical Model of the Atom
The Bohr model is considered “old quantum mechanics” because Louis de Broglie and Erwin
Schrödinger used mathematics to demonstrate the behavior of the electrons. Both
suggested if light can behave as both a particle and a wave then maybe electrons do too.
A.
Probability and the Electron
1.
The position and direction of motion of the electron cannot be simultaneously
determined
Translated: “The more certain I am about where it is, the less certain I can be about
where it is going. The more certain I am about where it is going, the less certain I
can be about where it is.”
B.
Regions of probability in which electrons may be found in an atom are determined by
mathematical equations (probability maps). These regions are called orbitals. Orbitals
suggest nothing about the motion of electrons.
Atomic Orbitals
A.
Atomic orbital
1. A region in space where there is a high probability of finding an electron
B.
Energy Levels of electrons (n)
1.
Indicates the distance of the energy level from the nucleus
2.
Values of n are positive integers
a.
3.
C.
n=1 is closest to the nucleus, and lowest in energy
The number of orbitals possible per energy level (or "shell") is equal to n 2
Energy Sublevels
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1.
Indicates the shape of the orbital
2.
Number of orbital shapes allowed in an energy level = n
a. Shapes in the first four shells are designated s, p, d, f
D.
Electron Spin
1.
A single orbital can contain only two electrons, which must have opposite spins
2.
Two possible values for spin, +1/2, -1/2
Lecture 14 will continue the discussion about the wave-mechanical model.
LECTURE NOTES
Wave Mechanical Model suggests electrons do not move in circular orbits but move in unpredictable orbits.
Thus this model shows the general location where atoms have a higher probability of showing up.
A probability map (an orbital) for electrons with higher probability in the middle (closer to the nucleus)
11.7 The Hydrogen Orbitals
Orbital: the probability map for predicting electron whereabouts
Orbitals boundaries are approximate and electrons can be found outside boundaries rarely
Different energy levels represent different geometric orientations for orbitals.
Principal Energy Levels: discrete energy levels
n=1, n=2, n=3, n=4
S-sublevel:
Spherical orbitals get bigger as energy increases
P-sublevel:
Characterized by two "lobes" on either x axis, y axis, and z axis
Example: 4pz, 4py, 4px
"Potential Space": orbitals have potential space for electrons
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Unit 1: Lecture Notes
11.8 The Wave Mechanical Model: Further Development
Pauli Exclusion Principal: An atomic orbital can hold a maximum of two electrons, and those two electrons
must have opposite spins.
Each of the two electrons in an orbital must
have opposite charges and there can only be
two in one orbital
11.9 Electron Arrangements in the First
Eighteen atoms on the Periodic Table
The Aufbau Principle – Electrons enter the
orbitals of lowest energy first. (See diagram
in book on page 345 for order of filling.)
The Pauli Exclusion Principle – An atomic orbital may contain at most two electrons. To occupy the same
orbital, two electrons must have opposite spins (clockwise or counterclockwise).
Hund’s Rule – When electrons occupy orbitals of equal energy, one electron enters each orbital until all the
orbitals contain one electron with spins parallel.
Orbital Diagram/Box diagram:
Example: Hydrogen: configuration: 1s^1: or orbital diagram:
Valence electrons: the electrons in the outermost (highest) principal energy level of an atom.
Core electrons: inner electrons, which are not involved in bonding atoms to each other.
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1d. Students know how to use the periodic table to determine the number of electrons available for
bonding.
Only electrons in the outermost energy levels of the atom are available for bonding; this outermost bundle of energy levels is often
referred to as the valence shell or valence shell of orbitals. All the elements in a group have
the same number of electrons in their outermost energy level. Therefore, alkali metals (Group
1) have one electron available for bonding, alkaline earth metals (Group 2) have two, and
elements in Group 13 (once called Group III) have three. Unfilled energy levels are also
available for bonding. For example, Group 16, the chalcogens, has room for two more
electrons; and Group 17, the halogens, has room for one more electron to fill its outermost
energy level.
To find the number of electrons available for bonding or the number of unfilled electron
positions for a given element, students can examine the combining ratios of the elements
compounds. For instance, one atom of an element from Group 2 will most often combine with
two atoms of an element from Group 17 (e.g., MgCl2) because Group 2 elements have two
electrons available for bonding, and Group 17 elements have only one electron position open
in the outermost energy level. (Note that some periodic tables indicate an elements electron
configuration or preferred oxidation states. This information is useful in determining how
many electrons are involved in bonding.)
11.10 Electron Configurations and the Periodic Table
Writing Electrons Configurations
A. Rules
1. Aufbau Principle
a. An electron occupies the lowest-energy orbital that can receive it
2. Pauli Exclusion Principle
a. No two electrons in the same atom can have the same set of four quantum numbers
3. Hund's Rule
a. Orbitals of equal energy are each occupied by one electron before any orbital is occupied
by a second electron, and all electrons in singly occupied orbitals must have the same spin
B. Orbital Notation
1. Unoccupied orbitals are represented by a line, _____
a. Lines are labeled with the principal quantum number and the sublevel letter
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Unit 1: Lecture Notes
2. Arrows are used to represent electrons
a. Arrows pointing up and down indicate opposite spins
C. Configuration Notation
1. The number of electrons in a sublevel is indicated by adding a superscript to the sublevel
designation
Hydrogen = 1s1
Helium = 1s2
Lithium = 1s22s1
D. Exceptional Electron Configurations
1. Irregularity of Chromium
a. Expected: 1s22s22p63s23p64s23d4
b. Actual: 1s22s22p63s23p64s13d5
2. Irregularity of Copper
a. Expected: 1s22s22p63s23p64s23d9
b. Actual: 1s22s22p63s23p64s13d10
3. Numerous transition and rare-earth elements transfer electrons from smaller sublevels in order
to half-fill, or fill, larger sublevels.
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LESSON 47: LIGHT PRO PERTIES AND WAVES
11.2 Energy and Light
Properties of Light
Light, which exists in tiny "packets" called photons, exhibits properties of
both waves and particles. This property is referred to as the wave–particle duality.
A.
Electromagnetic Radiation
1. Many types of EM waves
a. visible light
b. x-rays
c. ultraviolet light
d. infrared light
e. radio waves
2. EM radiation are forms of energy which move through space as waves
a. Move at speed of light
(1). 3.00 x 108 m/s
b. Speed is equal to the frequency times the wavelength
c = νλ
(1). Frequency is the number of waves passing a given point in one
second; designated by the Greek letter nu (ν).
(2). Wavelength (lambda, λ) is the distance between peaks of adjacent
waves
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IMPORTANT: Different wavelengths of EM carry different amounts of energy.
B. Light and Energy
1.
Radiant energy is transferred in units (or quanta) of energy called photons
a. A photon is a particle of energy having a rest mass of zero and carrying a
quantum of energy.
b. A quantum is the minimum amount of energy that can be lost or gained
by an atom.
2.
Energy of a photon is directly proportional to the frequency of radiation
a. E = h ν
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(h is Planck’s constant, 6.62554 x 10
-34
J sec)
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LESSON 48: ELECTROMA GNETIC SPECTRUM
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LESSON 49: PHOTOELECTRIC EFFECT
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LESSON 50: EMISSION SPECTRA
11.3 Emission of Energy by Atoms (Atomic Spectra)
See Demonstration notes of metals in methyl alcohol and Lab 13 (the flame test).
Atoms receive energy and become excited. These atoms release energy by photons, in
which the energy of the photons = energy change.
High energy photon = short wavelength
Low energy photon = long wavelength
For example, red light less energy than blue light.
A.
Ground State
1.
B.
Excited State
1.
C.
The lowest energy state of an atom
A state in which an atom has a higher potential energy than in its ground state
Bright line spectrum
1.
Light is given off by excited atoms as they return to lower energy states
2.
Light is given off in very definite wavelengths
3.
A spectroscope reveals lines of particular colors
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LESSON 51: ORBITALS (ELECTRONS AS WAVES)
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LESSON 52: QUANTUM N UMBERS
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LESSON 53: AUFBAU PRINCIPLE, PAULI EXCLU SION, AND HUND’S RULE
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LESSON 54: ELECTRON CONFIGURATION
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LESSON 55: VALENCE ELECTRONS
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LESSON 56: NUCLEAR RADIATION
Nuclear Composition
The radius of a nucleus is 10-14m, ten thousand times smaller than the radius of an atom. Nuclei are composed of
two protons and neutrons, called nucleons.
19.6 Nuclear Energy
The protons and neutrons are bound with forces much greater than chemical bond forces, reaching energies
greater than 1 million times. There are two types of nuclear processes which produce energy: fission and fusion.
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LESSON 57: HALF-LIFE
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LESSON 58: NUCLEAR REACTIONS
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LESSON 59: NUCLEAR D ECAY
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LESSON 60: BALANCING NUCLEAR REACTIONS
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LESSON 61: FISSION
19.7 Nuclear Fission
Heavy nuclides fragment in a process called fission. The process of fission was discovered in 1938 by Lise Meitner
and Otto Hahn in Germany. The first atomic bomb was detonated in the New Mexico desert at 5:30 AM on July 16,
1945.
A. Nuclear Fission
1. A very heavy nucleus splits into more stable nuclei of intermediate mass
2. The mass of the products is less than the mass of the reactants. Missing mass is converted to energy
a. Small amounts of missing mass are converted to HUGE amounts of energy (E = mc2)
B. Nuclear Chain Reaction
1. A reaction in which the material that starts the reaction is also one of the products and can start
another reaction
C. Critical Mass
1. The minimum amount of nuclide that provides the number of neutrons needed to sustain a chain reaction
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LESSON 62: FUSION
19.9 Nuclear Fusion
A. Nuclear Fusion
1. Light-mass nuclei combine to form a heavier, more stable nucleus
B. Fusion Reactions
1. More energetic than fission reactions
2. Source of energy of the hydrogen bomb
3. Could produce energy for human use if a way can be found to contain a fusion reaction (magnetic
field?)
Nuclear reactions
Ordinary chemical reactions involve changes in the outer electronic structures of atoms or molecules. In contrast,
nuclear reactions result from changes taking place within atomic nuclei. The nuclear reactions are represented by
nuclear equations:
14
7N
1
+ 0n 
14
6C
1
+ 1H
In which, the atomic numbers add to 7 on both sides and the mass numbers add to 15 on both sides.
147
61 Pm
64
29 Cu


0
1 e
150
61 Pm

27
13 Al
4
2
+
0
1 e
4
2
147
62 Sm
+
+
64
30
He +
He 
(beta particle emission)
Zn
146
59 Pr
30
15
(alpha particle emission)
1
P+ 0n
Application: Neutron bombardment was used in the preparation of transuranium elements, elements
above atomic numbers of 92.
Example: Smoke detectors contain a small amount of americium-241. Its decay product is neptunium-237.
Identify the emission from americium-241.
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LESSON 63: EFFECTS OF RADIATION
19.10 Effects of Radiation
The extent of damage is based on the amount of radiation absorbed and the type of radiation. The biological effect
of radiation is expressed in rems (radiation equivalent for man).
A. Penetrating Ability
1. Alpha Particles
a. Least penetrating ability due to large mass and charge
b. Travel only a few centimeters through air
c. Cannot penetrate skin
d. Can cause harm through ingestion or inhalation
2. Beta Particles
a. Travel at speeds close to the speed of light
b. Penetrating ability about 100 times greater than that of alpha particles.
c. They have a range of a few meters in air.
3. Gamma rays
a. Greatest penetrating ability
b. Protection requires shielding with thick layers of lead, cement, or both
C. Penetrating ability of radiation
Alpha
Least harmful
Beta
Gamma
Most harmful
C. Radioactive Elements
1. All isotopes of all man-made elements are radioactive
2. Some naturally isotopes are radioactive
a. All isotopes of all elements beyond bismuth (atomic #83) are radioactive
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