Download Chapter 4 Atomic Structure

9/18/2012
Defining the Atom
Atomic Structure
– smallest particle of an
___________ that retains the
chemical __________ of that
element
 Atom
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Defining the Atom
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Dalton’s Atomic Theory (experiment based!)‫‏‬
The Greek philosopher Democritus (460
B.C. – 370 B.C.) was among the ______
to suggest the existence of ________
(from the Greek word “atomos”)‫‏‬
 He
believed that atoms were indivisible and
indestructible
 His ideas did agree with later scientific
theory, but did not explain chemical
behavior, and was not based on the
scientific method – but just _____________
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John Dalton
(1766 – 1844)
1) All elements are composed of
tiny indivisible particles called
__________
2) Atoms of the same element are
__________. Atoms of any one
element are ___________ from
those of any other element.
1) Atoms of different elements combine in
simple whole-number ratios to form
chemical _____________
2) In chemical reactions, atoms are combined,
separated, or rearranged – but __________
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changed into atoms of another element.
Sizing up the Atom
 Elements are able to be subdivided into
smaller and smaller particles – these are
the ___________, and they still have
___________ of that element
If you could line up 100,000,000
copper atoms in a single file, they
would be approximately 1 cm long
Despite their small size, individual
atoms are observable with instruments
such as scanning tunneling (electron)
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microscopes
Structure of the Atom
 One
change to Dalton’s atomic
theory is that atoms are
___________ into subatomic
particles:
 ____________,
_________, and
____________ are examples of these
fundamental particles
 There are many other types of
particles, but we will study these
three
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Thomson’s Atomic Model
 Although
Thompson had the
positive and negative particles in
the wrong place, he is credited
with discovering ___________.
J. J. Thomson
Thomson believed that the electrons
were like plums embedded in a
_________ charged “pudding,” thus it
was called the “____ pudding”
model.
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Earnest Rutherford’s
Experiments
Conclusions:
a) The nucleus is _______
b) The nucleus is _______
c) The nucleus is __________
charged
This “fixed” Thompson’s plum
pudding model. Rutherford is
credited with discovering the
______________. 9
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The Rutherford Atomic Model

Based on his experimental evidence:
 The atom is mostly _________ space
 All the positive charge, and almost all the
_______ is concentrated in a small area in
the ________. He called this a “nucleus”
 The nucleus is composed of _______ and
___________ (they make the nucleus!)‫‏‬
 The
____________ distributed around the
nucleus, and occupy most of the _________
 His model was called a “nuclear model”
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Bohr Model

Modern (Wave) Model of the Atom
Neils Bohr discovered _______ _______
of electrons. He said that electrons
orbited the nucleus like planets orbit the
___…he was wrong about the orbiting but
right that there are particular energy levels
that electrons reside in.
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The modern model of the atom consists of
__________ that electrons move around
in…we will discuss later.
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About Atoms
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There are presently ___ elements, thus
118 _________ kinds of ________.
These atoms _________ in many different
combinations and proportions to form the
tremendous number of ___________
found.
Experiments determined that atoms
contain three ______________ particles:
protons, neutrons, and electrons.
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Atoms have a positively charged dense
central core called a nucleus. The nucleus
contains protons and neutrons.
Protons – __________ charge
Neutrons – _________ or zero charge
Protons and neutrons have about the
same mass.
Electrons move in the space around the
nucleus called the _______________.
Electrons – _________ charge
Electron’s mass is _____ _____ ____
than that of neutrons and protons.
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The nucleus is much more __________
but much ___________ than the electron
cloud. If you made a model of an atom to
scale – use __________ as nucleus, the
end of the atom would be 2 ___________
fields away from the baseball.
abbreviations commonly used: p+, no, e-
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Subatomic Particles
Particle
Electron
(e-)
Proton
(p+)
Charge
Mass (g) and
Mass (amu)
Location
9.11 x 10-28 g
Approx. _ amu
1.673 x 10-24 g
__ amu
Neutron
(no)
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Atomic Number – tells the number of
_______ in an atom. It is found on the
periodic table.
Ex: there are __ protons in an atom of
Nitrogen, there are ___ protons in an atom of
Uranium
The number of ________ in an atom makes
the atom what it is! Ex. Potassium has 19
protons it can never have any more or any
less and still be potassium.
Individual atoms are electrically _________,
which means they have the _____ number of
________ as __________. Ex. An atom of
copper has ___ protons and ___ electrons.
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1.675 x 10-24 g
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__ amu
Mass Number – ____ of _______ and
___________ of a particular atom.
Ex. “Chlorine-37”
Ex. “Aluminum-27”
Atomic number:
Atomic number:
Mass number:
Mass number:
p+:
p+:
e-:
e-:
no:
no:
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Mass Number
Naming Isotopes
 The
number written after an
element name is always the
______ number for a particular
isotope of that atom.
Mass number is the number of
protons and neutrons in the nucleus
of an isotope: Mass # = p+ + n0
p+
Nuclide
n0
Oxygen - 18
 carbon-12
Arsenic - 75
 carbon-14
Phosphorus - 31
 uranium-235
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Question: Why do you think the “mass”
number only includes protons and
neutrons?
Answer: because protons and neutrons
both have a mass of __ amu, but electrons
have a mass of __ amu
Question: Why do you think the mass
number is always a whole number?
Answer: because there are always
________ numbers of protons and
neutrons in an atom
Complete Symbols
Contain the symbol of the element,
the mass number and the atomic
number.
Mass
Superscript →
number

Subscript →
Atomic
number
X
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Symbols

e- Mass #
Find each of these:
– number of protons
– number of
neutrons
– number of
electrons
– Atomic number
– Mass Number
Symbols
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Br
If an element has an atomic
number of 34 and a mass
number of 78, what is the:
– number of protons
– number of neutrons
– number of electrons
– complete symbol
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Symbols
 If an element has 91
protons and 140 neutrons
what is the
– Atomic number
– Mass number
– number of electrons
– complete symbol 25
Symbols
 If an element has 78
electrons and 117 neutrons
what is the
– Atomic number
– Mass number
– number of protons
– complete symbol 26
Isotopes
Isotopes
Dalton was wrong about all elements of
the same type being _____________
 Atoms of the same element can have
___________ numbers of __________.
 Thus, different _________ numbers.
 These are called isotopes – atoms
with the same number of protons but
different number of neutrons...thus
they have different masses.
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Atomic Mass
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How heavy is an atom of oxygen?
 It depends, because there are different
________ of oxygen atoms.
We are more concerned with the average
atomic mass. (this number is found on the
_________ ______)‫‏‬
This is based on the _____________
(percentage) of each variety of that
element in _________.

We don’t use ________ for this mass
because the numbers would be too small.
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Elements occur in _________ as
____________ of isotopes.
Ex. Chlorine exists as chlorine-__ and
chlorine-__. Both have __ protons
but one has __ neutrons and one has
__ neutrons.
•There is no way of predicting which
isotopes exist for each element; these
have been _______________ determined.
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Measuring Atomic Mass
 Instead
of grams, the unit we use
is the Atomic Mass Unit (_____)‫‏‬
 Each
isotope has its own atomic
mass, thus we determine the
average from percent abundance,
which is why it is a ___________
and not a _________ number.
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Atomic Masses
Atomic mass is the average of all the
naturally occurring isotopes of that element.
Isotope
Symbol
Carbon-12
12C
Carbon-13
13C
Carbon-14
14C
Composition of
the nucleus
6 protons
6 neutrons
6 protons
7 neutrons
6 protons
8 neutrons
Carbon = 12.011
% in nature
98.89%
1.11%
 To
get the “most __________
mass number”, __________ the
atomic mass off to the nearest
whole number.
 Example: What is the most
common mass number of the
following? Ag, Cu, C, Cl
<0.01%
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Draw models of lithium-5, lithium-6, and
lithium-7 atoms including the correct
numbers of protons, neutrons, and
electrons in each.
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IONS
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Ion – atom that has _______ or ____ electrons
Two types:
_________ – a positive ion, one that has
_______ electrons
_________ – a negative ion, one that has
_________ electrons
Atoms NEVER gain or lose __________ or
____________ (except in nuclear reactions)‫‏‬
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Examples
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Examples:
Mg is a magnesium atom and has ___ protons
and ___ electrons so the overall charge is ____.
Mg+2 is a magnesium ion and has ___ protons
and ___ electrons so the overall charge is ____.
Cl is a chlorine atom and has ___ protons and
___ electrons so the overall charge is ____.
Cl- is a chlorine ion has ____ protons and
___ electrons so the overall charge is ____.
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Energy Levels
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Maximum number of electrons
Who was the first to notice that the electron cloud
has different regions where electrons reside?
There are various ______________ in an electron
cloud. Within each energy level there are
_____________ and within each sublevel there are
__________. Each orbital can hold up to two
electrons.
Electrons move very ________ within their own
______________.
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The total number of __________ levels an atom
has corresponds to the _________ number of
the atom.
Periods are the ______________ rows on the
PT
Ex. An atom of bromine has ___ energy levels
because it is in period ___
The ______ energy level, called the
____________ shell, can hold ONLY up to
______ electrons.
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Number of Valence e-
Electron dot diagrams
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3-12 and the actinides and
lanthanides (B’s)
13 (3A)
14 (4A)
_____________ electrons are the electrons in
the __________ energy level.
The atoms are arranged on the periodic table so
that the ones with ___________ properties all
line up in a ____________ or __________.
The ____________ of valence electrons plays a
big role in how the atom _____________.
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1 (1A)
2 (2A)
Max num
of e-
Valence Electrons
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Group Number
Energy
level
1
2
3
4
5
6
7
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More about energy levels…

Energy levels (n) can
hold a maximum number
of 2n2 electrons.
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Use the number of valence electrons as _______
and space them around the ________ sides of
the element’s symbol then ________ up as
needed
Examples:
15 (5A)
16 (6A)
17 (7A)
18 (8A)
Except He (2)
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Group/Family Names to Label
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Alkali Metals
Alkaline Earth Metals
Transition Metals
Lanthanides
Actinides
Boron Family
Carbon Family
Nitrogen Family
Oxygen Family
Halogens
Noble Gases
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Periods
Ion Charge Prediction
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Octet rule – atoms will ________, _________ or
share electrons so that they have a
______________ ________________ shell
That's ____ electrons for everything except 1st
shell (H and He) which only holds __ electrons
Ex. Nitrogen has ___ valence electrons so it will
gain ___ to make ___. N-3
Ex. Sodium has___ valence electron so it will lose
that one and the _________ shell becomes the
valence shell with ___ electrons. Na48+
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Group 1 (1A) makes ____
Group 2 (2A) makes ____
Group 3-12 (B's) are __________________
Group 13 (3A) makes ____
Group 14 (4A) makes ____
Group 15 (5A) makes ____
Group 16 (6A) makes ____
Group 17 (7A) makes ____
Group 18 (8A) makes ____ (doesn't form ions!)
why?
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Electron Configurations
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Examples:
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Sublevels: s, p, d and f
AUFBAU principle – electrons fill orbitals
starting with the _____________ energy orbital
available before filling higher energy orbitals.
PAULI EXCLUSION principle – each orbital
can hold at most _ electrons and they must have
_______ spins (clockwise and couterclockwise)‫‏‬
HUND'S rule – electrons occupy equal-energy
orbitals so that a maximum number of _______
electrons results ex. __ __ __ not __ __ __
**spread out before pairing up**
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Orbitals
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Writing electron configurations

Use the periodic table as a guide.
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Abbreviated Electron
Configurations
More Electron Configurations

Use the previous _______________ to shorten
the electron configuration.
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Orbital Notation

Exceptions to Orbital Fill Order
Use ________ to represent electrons.
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Any electron configuration that ends in ____ is
too ___________ and actually takes an electron
from the previous ___ sublevel and becomes
____.
Example:
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Another Exception
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Examples:
Any electron configuration that ends in ____ is
too ___________ and actually takes an electron
from the previous __ sublevel and becomes
_____.
Example:
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Absorption and Emission
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Normally, electrons in an atom are in the
____________ state, which means they are in the
_______________ possible energy levels.
However, these electrons can be ___________ to
higher energy levels if energy is added...called
_________________.
If energy is absorbed and the electrons jump to
higher energy levels they are now in the
________________ state.
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Once in the excited state the atoms are
______________ and thus _______ energy when
they fall back down to their original energy levels.
This process is called _______________.
This __________ of absorption and emission
happens very fast over and over again.
Atoms can be excited using _________, light, or
electricity.
Emission is usually in the form of __________.
Different energies of light have different
___________. The light spectrum:
(lower energy)ROYGBIV (higher energy)‫‏‬
This is actually how we see colors! 68
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