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Electrons in Atoms
A chapter dedicated just for you
electrons!
Chapter 5
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Light Waves
Amplitude = wave’s height
Wavelength = wave’s distance
Frequency = wave cycles (hertz)
Electromagnetic radiation
= A big word for Light
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4
Electromagnetic Radiation
• Waves have a frequency

• Use the Greek letter “nu”, , for frequency,
and units are “cycles per sec” or Hertz (Hz)
  = c
• All radiation:
•
where c = velocity of light = 2.998 x 108 m/sec
Electromagnetic Spectrum
Long wavelength --> small frequency
Short wavelength --> high frequency
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INVERSE
RELATIONSHIP
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Electromagnetic Spectrum
In increasing energy, ROY G BIV
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White light is actually all the
colors together
A prism bends
the different
colors to where
we can see them
Atomic Emission Spectra
and Niels Bohr
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Bohr’s greatest contribution
to science was in building a
simple model of the atom.
It was based on an
understanding of the
Niels Bohr
(1885-1962)
ATOMIC EMISSION
SPECTRA of excited
atoms
- Remember the Bohr-rings!
Atomic Emission Spectra
Excited Atoms
• Excited atoms emit light of only
certain wavelengths
• The wavelengths of emitted light
depend on the element
Each element has its own Atomic Emission
Spectra… like a FINGERPRINT
ELECTRICITY or HEAT EXCITED ELECTRONS!
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5.2 = More about electrons
Heisenberg Uncertainty Principle
Problem of defining nature
of electrons in atoms
solved by W. Heisenberg
Cannot simultaneously
define the position of an
electron
W. Heisenberg
1901-1976
We define e- energy exactly
but do not know exact
position
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We know area
the propeller
is, but we
cannot locate
its exact
position
Arrangement of
Electrons in Atoms
Electrons in atoms are arranged as
LEVELS
ORBITALS
SUBLEVELS
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QUANTUM NUMBERS
The shape, size, spin, and energy of each orbital is a
function of 4 quantum numbers which describe the
approximate location of an electron
n (principal) ---> energy level = FROM BOHR!
l (orbital) ---> shape of orbital
ml (magnetic) ---> designates a particular
suborbital
s (spin)
---> spin of the electron
(clockwise or counterclockwise)
Think of the 4 quantum numbers as the address of an
electron… Country > State > City > Street
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QUANTUM NUMBERS
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So… if two electrons are in the same place at
the same time, they must be repelling
(like charges repel!)
The Pauli Exclusion Principle says that an
atomic orbital may describe at most 2
electrons. If two electrons are in the same
energy level, the same sublevel, and the same
orbital, they must repel = spin (S)
Pauli = spinning electrons (repel)
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Energy Levels
• Each energy level has a number
called the PRINCIPAL
QUANTUM NUMBER,
n (energy level)
• Currently n can be 1 thru 7,
because there are 7 periods on
the periodic table
• The period #s = energy levels
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Energy Levels
n=1
n=2
n=3
n=4
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Types of Orbitals
• The most probable area to find
these electrons takes on a shape
• So far, we have 4 shapes
They are named s, p, d, and f
• No more than 2 e- assigned to an
orbital – one spins clockwise, one
spins counterclockwise
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Types of Orbitals (l)
s orbital
p orbital
d orbital
Relative sizes of the spherical 1s,
2s, and 3s orbitals of hydrogen.
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Aufbau’s Principal
• Electrons occupy the orbitals
of lowest energy first!
• Aufbau = electron order….
Least to most energy!
Orbitals and the Periodic
Table
• Orbitals grouped in s, p, d, and f orbitals
(sharp, proximal, diffuse, and fundamental)
s orbitals
f orbitals
d orbitals
p orbitals
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Electron Configurations
A list of all the electrons in an atom
• Must go in order (Aufbau principle)
• 2 electrons per orbital, maximum (Pauli Ex.)
• We need electron configurations so that we
can determine the number of electrons in the
outermost energy level
These are called valence electrons
• The number of valence electrons determines
how many and what this atom can bond to in
order to make a molecule
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14… etc.
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Electron Configurations
4
2p
Energy Level
(n)
Number of
electrons in
the sublevel
Sublevel
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6
6s2 4f14… etc.
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Diagonal Rule
Steps:
1s
2s
3s
1.
Write the energy levels top to bottom.
2.
Write the orbitals in s, p, d, f order. Write
the same number of orbitals as the energy
level.
3.
Draw diagonal lines from the top right to the
bottom left.
4.
To get the correct order,
2p
3p
3d
follow the arrows!
4s
4p
4d
4f
5s
5p
5d
5f
5g?
6s
6p
6d
6f
6g?
6h?
7s
7p
7d
7f
7g?
7h?
By this point, we are past
the current periodic table
so we can stop.
7i?
Let’s Try It!
Write the electron configuration for
the following elements:
H 1s1
Li 1s2 2s1
N 1s2 2s2 2p3
Ne 1s2 2s2 2p6
K 1s2 2s2 2p6 3s2 3p6 4s1
Zn 1s2 2s2 2p6 3s2 3p6 4s2 3d10
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Orbital Diagrams
• Graphical representation of an
electron configuration
• One arrow represents one
electron
• Shows spin and which orbital
within a sublevel
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Orbital Diagrams
• Hund’s Rule
– In orbitals of EQUAL
ENERGY (p, d, and f), place
one electron in each orbital
before making any pairs
– All single electrons must
spin the same way
• I nickname this rule the
“Monopoly Rule”
• In Monopoly, you have to build
houses EVENLY. You can not
put 2 houses on a property
until all the properties have at
least 1 house.
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Lithium
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Group 1A
Atomic number = 3
1s22s1 ---> 3 total electrons
Carbon
Group 4A
Atomic number = 6
1s2 2s2 2p2 --->
6 total electrons
HUND’S RULE!
Same energy level
= 1 electron until
All filled
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Draw these orbital diagrams!
• Lithium (Li)
• Beryllium (Be)
• Boron (B)
• Carbon (C)
• Nitrogen (N)
• Neon (Ne)
• Sodium (Na)
it helps to do the
e- configuration
first
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Ch. 6.1 & 6.2
Dmitri Mendeleev (1869)
Mendeleev published
classification schemes for
elements known to date.
The periodic table has
organized the similar
properties and reactivities
by certain elements
Later, Henri Moseley
established that each
elements has a unique
atomic number
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The Periodic Table
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http://www.chemsoc.org/viselements/pages/periodic_table.html
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Periodic Law
When elements are arranged in
order of increasing atomic
number, there is a periodic
repetition of their physical and
chemical properties
Periodic Table Expanded View
The Periodic Table can be arrange by
subshells
The s-block is Group IA and & IIA
The p-block is Group IIIA - VIIIA
The d-block is the transition metals
The f-block are the Lanthanides and Actinide
metals
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Periodic Table: Metallic arrangement
Layout of the Periodic Table:
Metals vs. nonmetals
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METALS
• 80 percent of all elements
• Are GOOD conductors of
heat and electric current
(Give up electrons)
• Have a high luster or sheen
• Reflects light
• All metals are solids at room temperature
(except mercury Hg)
• They are ductile (can be drawn into wires)
• They are malleable (can be hammered down)
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NONMETALS
• Upper-right corner
• Variation in physical properties:
• Accept electrons
– Most are gas at room temperature (N and O)
– Few are solids (sulfur and phosphorus)
– One is a liquid (bromine—dark red)
• Most nonmetals have properties
opposite of metals:
• Poor conductors (except carbon)
• Brittle when solid
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METALLIODS
• Hug the stair-step line
• Have properties similar to both metals
and nonmetals, depending on certain
conditions
• Give up or accept electrons
• EXAMPLE:
Pure silicon is a poor conductor of electrical
current
But… if silicon is mixed with boron, it is a great
conductor of electrical current
Used to make computer chips
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Across the Periodic Table
Periods: Are arranged horizontally across
the periodic table These elements have
the same number of valence shells.
1
IA
1
18
VIIIA
2nd Period
2
IIA
13
IIIA
2
3
3
IIIB
4
IVB
5
VB
6
VIB
4
5
6
7
6th Period
7
VIIB
8
9
VIIIB
10
11
IB
12
IIB
14
IVA
15
VA
16
VIA
17
VIIA
Down the Periodic Table
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Family: Are arranged vertically down the periodic
table = These elements have the same number
electrons in the the valence shell
1
IA
1
18
VIIIA
2
IIA
13
IIIA
2
3
4
5
6
7
3
IIIB
4
IVB
5
VB
6
VIB
7
VIIB
8
9
VIIIB
10
11
IB
12
IIB
14
IVA
15
VA
16
VIA
17
VIIA
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Groups
The specific colors of the squares
distinguish the groups of elements
Group 1A = Alkali Metals
Group 2A = Alkaline Earth Metals
Alkali is the arabic word al aqali which means
wood ashes
Wood ashes are rich in Na and K
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Groups
The specific colors of the squares
distinguish the groups of elements
Group 7A = Halogens
Halogen comes from the Greek word hals
(salt) and the Latin word genesis (to be
born)
Chlorine, bromine, and iodine can be
prepared from their salts
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Electron Configurations
Elements can be sorted into noble gases,
representative elements, transition metals, or
inner transition metals based on their
electron configurations
The noble gases are group 8A… they rarely take
place in reactions
IF YOU LOOK AT THEIR CONFIGURATIONS…
THERE VALENCE SHELL IS FULL!!!
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Infamous Families of the Periodic Table
Notable families of the Periodic Table
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Representative Elements
Groups 1A- 7A
They display a wide range of physical and chemical
properties
For representative elements…
The group number is equal to the number of electrons
in the highest occupied energy level
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Transition Elements
Transition metals are characterized by electrons in the d
sublevel
Inner transition metals are characterized by electrons in
the f sublevel
The periodic table is
organized into ‘blocks’
Based on sublevels
(except for helium)
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General Periodic Trends
6.3
• Atomic and ionic size
• Ionization energy
• (electronegativity– learn in future chapter)
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Atomic Radius
Is one-half the distance between the
nuclei of two atoms of the same
element when atoms are joined
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Atomic Size
• Size increases- going down a group
• Each element has 1 more proton and 1
more electron than the preceding one
• The increasing nuclear charge pulls
the electrons in the highest occupied
energy level closer to the nucleus and
the atomic size decreases
• Size decreases- going across a period
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Which is Bigger?
• Na or K ?
K
• Na or Mg ? Na
• Al or I ?
Al
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What are IONS??
• An atom or group of atoms that has a positive
or negative charge
• An atom is electrically neutral because it has
equal numbers of protons and electrons…
– EX: Na has 11 protons and 11 electrons = 0 charge
• Positive (cation) and negative (anion) ions
form when electrons are transferred between
atoms
• ELECTRONS determine a charge!
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IONS
• Cations = positive charge
• Lose electrons
• Usually in metallic elements
• Anions = negative charge
• Gain electrons
• Usually nonmetallic elements
How do you think the trend is going go?
Ion Sizes
Li,152 pm
3e and 3p
Does+ the size go
up+ or down
Li , 60 pm
when
an
2e and 3losing
p
electron to form
a cation?
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Ion Sizes
+
Li,152 pm
3e and 3p
Li + , 78 pm
2e and 3 p
Forming
a cation
• CATIONS are SMALLER than the
atoms from which they come
• The electron/proton attraction
has gone UP and so size
DECREASES
– (forced closer to nucleus)
Ion Sizes
Does the size go up or
down when gaining an
electron to form an
anion?
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Ion Sizes
F, 71 pm
9e and 9p
F- , 133 pm
10 e and 9 p
Forming
an anion
• ANIONS are LARGER than the atoms
from which they come
• The electron/proton attraction has
gone DOWN and so size INCREASES
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Which is Bigger?
• Cl or Cl2+ ?
• K+ or K ?
• F or F-?
• I- or Br- ?
Cl
K
FI- (b/c size)
Ionization Energy
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IE = energy required to remove an electron
from an atom (in the gas phase)
This is called the FIRST
ionization energy because
we removed only the
OUTERMOST electron
Trends in Ionization Energy
• IE increases across a
period because the
positive charge increases.
• Metals lose electrons
more easily than
nonmetals
• Nonmetals lose electrons
with difficulty (they like to
GAIN electrons)
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Trends in Ionization Energy
• IE increases UP a
group
–Because size
increases
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Which has a higher 1st
ionization energy?
• Mg or Ca ?
• Al or S ?
• Cs or Ba ?
Mg
S
Ba
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Electronegativity
Is a measure of the ability of
an atom in a molecule to
attract electrons
Will talk about more later…