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Write about what you know of the following
people:
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John Dalton
J. J. Thomson
Ernest Rutherford
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Key Learning: The number, type and arrangement of subatomic particles
differs with each element.
Concept
CHOOSE 1: HOW TO DO VOCABULARY:
 Write word and definition.
Underline and highlight
vocabulary word.
 Write word on front and
definition on back of index
cards.
 Create a poem or rap of
vocabulary words and
definitions.
 Create foldable with words
and definitions.
VOCABULARY
Subatomic particles
 Orbits
 Nucleus
 Atom***
 Electron cloud
 Ground state
 Excited state
 Absorption
 emission
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John Dalton (1766–1844)
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Dalton’s Atomic Theory
First to propose that elements were
comprised of atoms
Write 5 parts to Dalton’s Atomic Theory pg.
68
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Modern Atomic Theory
***modification of Dalton’s Atomic
Theory***
1. All Matter is composed of atoms
2. Atoms of any one element differ in
properties from atoms of another element
How is this different than Dalton’s Theory?
 there is no mention of an atom being
smallest particle
Investigated cathode rays.
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J.J. Thomson
Discovered electrons ~1897 by using cathode
ray tubes
“plum-pudding” model
Thought the atom was solid
Cathode (-) Anode (+) Electric plates (one +, one -)
partially evacuated glass tube
that produces a beam between
the cathode and anode
Discovered:
Alpha particles
and
Beta particles
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Ernest Rutherford
 Credited with the development of the nuclear model for
the atom
Used gold foil method. This disproved Thomson’s
model.
 3 postulates
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 Most of atom’s mass and all of its positive charge are in a
localized small core---nuclear theory
 Most of atom’s volume is empty space in which tiny,
negatively charged electrons are dispersed
 Since electrons are electrically neutral, there are as many
(+) particles (protons) with the nucleus as there are (-)
electrons outside the nucleus – or – protons = electrons
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10/08/2012
Write Question and Answer
What did Rutherford do that disproved
Thomson’s model?
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Developing a New Atomic Model
EM Spectrum Diagram
 Energy moves as a wave
 Includes all electromagnetic radiation
 Visible light is only a small part of the EM
spectrum (400-700 nm)
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EM radiation
1.Properties of Light
a. visible light
 Form of energy traveling in waves
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b. speed of light (c)
 Wavelength X frequency
 The shorter the wavelength, the greater the
frequency
c. Characteristics of Waves
 Wavelength
 Distance from peak to peak of a wave
 Measured in nm
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Frequency
 Number of waves that pass through a set point in
a given amount of time (waves/sec)
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Photoelectric effect
Emissions of electrons from metals that have
absorbed photons
Photon
 A particle of EM radiation that has a zero mass
and a quantum amount of energy
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Quantum
 Minimum amount of energy that can be gained or
lost by an atom
Bohr model
 Lowest energy state of an atom  ground state
 An atom having a higher potential energy 
excited state
 Energy must be absorbed to move electrons to a
higher orbit (ground state  excited state)
 Energy is emitted (given off as EM radiation)
when electrons are moved to a lower orbit
(excited state  ground state)
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When an element is heated up, it
gives off light of specific colors:
emission spectrum
When light goes through an element, it
absorbs light of specific colors:
absorption spectrum
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Bohr model continued…
Electrons circle nucleus in orbits (energy
levels)
When electrons are in orbit, its energy is fixed
Lowest energy levels are closest to the
nucleus while Highest energy levels are
farther away from nucleus
Electrons can move to higher energy levels
when hit by a photon
Structure of the atom
3 subatomic particles identify the element
1. protons
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 Ernest Rutherford (1909)
 (+) charge
 Mass 1.674 X 10 -²7 g – or – 1 amu
 Relative mass 1.007 276
 Held by nuclear forces
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2. Neutron (N)
 James Chadwick (1932)
 No charge; neutral
 Has mass of 1.675 X 10-27 g – or – 1 amu
 Relative mass = 1.008 665
 In nucleus and held by nuclear forces
 Responsible for isotopes- atoms of same element
with different #’s of neutrons
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3. Electron (e-)
J.J. Thomson (1897)
( – ) charge
Negligible mass (1/1836 amu –or- 9.11 X1031g
Relative mass 0.000 5486
In orbitals within the electron cloud
Makes volume of atom
# in last energy level determines the chemical
activity
 Things
to Note:
 Mass of atom is in the nucleus
 Atomic radius (Angstrom) is distance from
nucleus to outer edge of electron cloud
 Volume of atom is from electron cloud
Concept
 Write word and
definition. Underline and
highlight vocabulary
word.
 Write word on front and
definition on back of
index cards.
 Create a poem or rap of
vocabulary words and
definitions.
 Create foldable with
words and definitions.
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Atomic number
Average atomic mass
Isotopes
Protons
Neutrons
Electrons
Mass number
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Write question and answer.
Write the correct nuclear symbol for the
following:
An atom has 23 protons and 29 neutrons.
An atom has 42 protons and 49 neutrons.
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You can use information from periodic table
to determine the structure of an atom
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Atomic number (identifies the element)
Atomic symbol
Element name
Average atomic mass  round to nearest
whole number to get the mass # (total
number of protons and neutrons that make
up nucleus of an isotope.
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Where does average atomic mass come
from?
Average of atomic masses of naturally
occurring isotopes of that element
Example: We have a box of 2 types of marbles
(100 total). If 25% of these marbles have a
mass of 2 g and 75% have a mass of 3 g, then
what is the average mass of the marbles?
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APPLY to elements!!!
We have 69.15% of Cu-63, with an atomic
mass of 62.929601 amu and 30.85% of Cu-65
with an atomic mass of 64.927794. What is
the average atomic mass?
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You APPLY to elements!!!
There is 99.757 % of O-16 with an atomic
mass of 15.994915 amu and .038% of O-17
with an atomic mass of 16.999132 amu and
0.205% of O-18 with an atomic mass of
17.999160 amu. Calculate the average atomic
mass.
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Isotopes
They have the same number of protons and
electrons as a neutral element; but, different
number of neutrons (therefore different
masses).
Even though this affects average atomic
mass –or- mass number, is does not
significantly alter their chemical behavior
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Nuclear Symbol
 Superscript indicates the mass number (p+n)
 Subscript indicates the atomic number (p)
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Hyphen Notation
 Element symbol, hyphen, mass number
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Mathematical Equations
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Atomic number = # of p (also tells you # of e)
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Mass number = # of p + # of N
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Number of Neutrons = mass # - atomic #
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Complete worksheet 2 and 3
Concept
CHOOSE 1: HOW TO DO VOCABULARY:
 Write word and definition.
Underline and highlight
vocabulary word.
 Write word on front and
definition on back of index
cards.
 Create a poem or rap of
vocabulary words and
definitions.
 Create foldable with words
and definitions.
VOCABULARY
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Quantum numbers
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With the Bohr model, we can show what
energy level the electrons are in. However,
the Bohr model does not show where the
electron is exactly located or what orbital
they are in.
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Where are electrons located in an atom?
Electron cloud
Electrons are located in specific locations
within energy levels. These levels are known
as orbital “home”
Atomic orbitals vary in amount of energy,
shape and location outside of nucleus
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Quantum numbers are used to describe the
atomic orbitals and the electrons which
occupy them
“address” of an electron
Quantum Theory-describe proportion of
electrons mathematically
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4 Quantum Numbers
Principle Quantum Number (n)
 Energy level of orbital
 7 energy levels (7 periods of table)
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Angular Momentum Quantum Number (l)
 Indicates the shape of the orbital
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Magnetic Quantum Number (m)
 Indicates the orientation of the orbital
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Spin Quantum Number (+ ½ or – ½ )
 Electrons are constantly spinning
(clockwise/counter-clockwise)
Concept
CHOOSE 1: HOW TO DO VOCABULARY:
 Write word and definition.
Underline and highlight
vocabulary word.
 Write word on front and
definition on back of index
cards.
 Create a poem or rap of
vocabulary words and
definitions.
 Create foldable with words
and definitions.
VOCABULARY
Orbital notation
 Electron configuration
notation
 Noble gas notation
 Hund’s Rule
 Aufbau’s Principle
 Pauli Exclusion Principle
 Heisenberg Uncertainty
Principle
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Pauli Exclusion Principle
No two electrons can have the same 4
quantum numbers
Atomic orbitals
 Electrons are not only located in energy levels;
but, they are more specifically located in orbitals
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4 types of Atomic Orbitals
s, p, d, and f
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S-orbital
Shaped like a sphere
1 per energy level
Can hold 2 electrons
S-block is groups 1-2
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P orbital
Shaped like a dumbbell
3 per energy level
One on x-axis, y-axis, and z-axis
Each one holds 2 electrons for a total of 6
electrons
Groups 13-18
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D-orbital
Begins at energy level 3
5 different orbitals
Each hold 2 electrons for a total of 10
electrons
Groups 3-12
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F-orbital
Begins at 4th energy level
7 different orbitals
Each has 2 electrons for a total of 14 electrons
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Be looking over notes from board and
yesterday’s powerpoint
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Aufbau Principle
Electrons fill the atomic orbitals in a very
specific way; based on energy
An electron will occupy the lowest energy
orbitals first
Follow the arrows to fill atomic orbital
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Hund’s Rule
Orbitals of equal energy are filled by one
electron before a second electron is added
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Using orbital notation (combining Aufbau
and Hund’s)
Use dashes to represent orbitals and arrows
to represent electrons
1. Use atomic number to know how many
electrons are present
2. Fill in orbitals based on Aufbau’s chart and
Hund’s Rule
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10/18/2012
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Write the orbital configuration for:
 Sodium
 Magnesium
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Electron Configuration Notation
Instead of using dashes and arrows, use
coefficients and superscripts
1. use atomic number to know how many
electrons are present
2. exponents show number of electrons in
each orbital
3. coefficients represent energy level
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Noble Gas Notation
Use noble gases to make notation shorter
Noble gas  has a full s and p orbital
 octet
To read: Add the electrons to the atomic
number of the noble gas
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Valence Electrons
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Electrons located in the last energy level
Determines the chemical properties of that
element
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