Download Electrons - Mrs. B Chemistry

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Electrons and Periodicity
• Adapted from www.chemistrygeek.com
Rutherford Model of the Atom
(1911)
What is wrong with this model of the atom?
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Bohr Model of the Atom (1913)
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•Electrons are found in circular paths around the nucleus
•The paths are called Energy Levels (n)
•Each level has a specific energy
•The lowest level is called Ground State
•Electrons can jump to higher levels
by absorbing energy (Excited State)
•Electrons can return to a lower
level by releasing energy (light)
•Bohr’s model only works for
atoms of Hydrogen
Wave-Particle Duality
• Photons of light and most subatomic
particles (i.e. electrons) behave as both
PARTICLES and WAVES.
https://www.youtube.com/watch?v=MB03zWIQO1Q
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Heisenberg Uncertainty
Principle (1927)
W. Heisenberg
1901-1976
Problem of defining nature
of electrons in atoms
solved by W. Heisenberg.
Cannot simultaneously
define the position and
momentum (= m•v) of an
electron.
We can define electron
energy exactly but accept
limitation that we do not
know exact position.
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Heisenberg Uncertainty
Principle (2015)
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Electron Cloud Model (1926)
•
•
•
•
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Proposed by Erwin Schrödinger
Applied idea of electrons behaving as a wave
Electrons are no longer in defined paths/orbits
Uses Electron Clouds (Orbitals) – defined areas
where an electron is likely to be found
• Based on probability
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Electromagnetic Spectrum
In increasing energy, ROY G BIV
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Wave Behavior
wavelength
Visible light
Amplitude
wavelength
Node
Ultraviolet radiation
Visible Spectrum
Long wavelength --> small frequency
Short wavelength --> high frequency
What colors have
--long wavelengths?
--high frequencies?
--short wavelengths?
--small frequencies?
increasing
frequency
increasing
wavelength
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Wave Behavior
• Wavelength – distance between 2 crests of a wave
– Abbreviated with Greek letter lambda (
– Units are meters (m)
λ)
• Frequency – speed at which crests pass a given point
– Abbreviated with Greek letter nu (
– Units are 1/s, s-1 or Hertz
ν)
• Wavelength and Frequency are related by the equation:
• = c
– c =speed of light = 3.00 x 108 m/sec
– Inverse relationship (one goes up, the other goes down)
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Wave Behavior
• Energy and frequency are directly related
• Amount of energy released by an excited
electron corresponds to a specific frequency
of light
•E=hxν
–(E) Energy measured in Joules (J)
–(h) Planck’s constant = 6.626 x 10-34 J s
–(ν) Frequency measured in 1/s or hertz
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Wave Behavior
• What is the frequency of a wave that has a wavelength of
2.56 x 10-7 m?
c
= λ
x
ν
2.8 x 108 m/s
= (2.56 x 10-7 m)
2.8 x 108 m/s
2.56 x 10-7 m
=
ν
1.1 x 1015 1/s
=
ν
• What is the Energy of this wave?
E
=
h
x
ν
x
ν
E
= (6.626 x 10-34 Js) x (1.1 x 1015 1/s)
E
=
7.2 x 10-19 J
Learning Check
• What is the frequency of EMR having a
wavelength of 5.55 x 10-7 m?
• What is the Energy of this same wave?
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Continuous Spectrum of
White Light
•If you separate white light with a prism you will
see ALL colors of the rainbow
•Continuous because there are no gaps between
colors
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Creating an Atomic Spectrum
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• Discharge tube – tube containing atoms of a gas
connected to an electricity source
• When electricity is turned in the gas will emit
light
Atomic Spectrum of
Excited Hydrogen Gas
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•If the light from a discharge tube is separated by a prism
you will only see very specific lines of color separated by
black regions
•Called atomic emission spectra or line spectra
Atomic Emission Spectra
of Excited Atoms
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• When electricity is supplied to the gas atoms
the electrons absorb Energy and become
EXCITED
• Electrons eventually have to release this energy
in the form of light
• Excited atoms emit light of only certain
wavelengths (lines of color)
• The wavelengths of emitted light depend on the
element.
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Line Spectra of Other Elements
Arrangement of
Electrons in Atoms
Electrons in atoms are arranged as
LEVELS (n)
SUBLEVELS (l)
ORBITALS (ml)
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Quantum Numbers
•Every electron has a unique set of 4 quantum
numbers that defines it’s position in the atom
•Think of it as the electron’s address!
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Quantum Numbers
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Learning Check
• Write a set of quantum numbers for the 4f
orbital
• Which orbital is represented by the following
sets of Quantum numbers
a.
b.
c.
d.
e.
n = 3, l = 0, ml = 0
n = 2, l = 1, ml = 1
n = 4, l = 2, ml = -1
n = 3, l = 3, ml = 2
n = 3, l = 1, ml = 2
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Energy Levels
• Each energy level has a number called
the PRINCIPAL QUANTUM NUMBER
(n)
• Currently n can be 1 thru 7, because
there are 7 periods on the periodic
table
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Energy Levels
n=1
n=2
n=3
n=4
Types of Orbitals
• Orbital - most probable area to find an
electron
• Orbitals can take on different shapes
• 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
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S Orbital
(sharp)
• The s orbital is the
simplest orbital
• The shape is a
sphere
• An electron can
move anywhere
within the sphere
• L=0
• Every principal level
(n) has an s orbital
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Relative sizes of the spherical 1s,
2s, and 3s orbitals of hydrogen.
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p Orbitals
(proximal)
• A p orbital has a dumb-bell
shape
• There are three different
directions for p-orbitals
• These are called x, y, and z
• The electron can move
anywhere in the nodes
• L=1
• Every level above n = 1 has a
p orbital
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There is a PLANAR
NODE thru the
nucleus, which is
an area of zero
probability of
finding an electron
p Orbitals
• The three p orbitals lie 90o apart in space
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d Orbitals
(diffuse)
• A d orbital has a
clover shape
• There are 5 different
directions for a d
orbital
• L= 2
• Every level above
n=2 has a d orbital
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The shapes and labels of the
five 3d orbitals.
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f Orbitals
(fundamental)
• f orbitals have a flower shape
• There are 7 directions for d
orbitals
• L= 3
• Every level above n = 3 has an
f orbital
Electron Configurations
• The arrangement of electrons in an atom into
atomic orbital
• Electron configurations are filled according to
three rules (Aufbau, Pauli, Hund)
• 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
the number and types of bonds and atom can
make
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14… etc.
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Electron Configurations
4
2p
Energy Level
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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Electron Configuration Rules
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• Aufbau Principle – electrons are filled starting from the
lowest energy levels first
• Pauli Exclusion Principle – no more than 2 electrons
can fill into orbitals at one time and they must have
opposite spin (clock-wise/counter-clockwise)
• Hund Rule – In p, d and f sublevels, electrons will have
to fill one clockwise electron into all available orbitals
before filling in counter-clockwise electrons
Energy Level (n)

Sublevel (l) 
Orbital (ml)
Aufbau Principle
• Aufbau Principle states that electrons
fill from the lowest possible energy to
the highest energy
• The Aufbau Diagram is a memory
device that helps you remember the
order of the filling of the orbitals from
lowest energy to highest energy
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Aufbau Diagram
Why are d and f orbitals always
in lower energy levels?
•d and f orbitals require LARGE
amounts of energy to fill
•It’s easier to fill an s orbital
before a d or f
•This is the reason for the
diagram
•BE SURE TO FOLLOW THE
ARROWS IN ORDER!
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Pauli Exclusion Principle
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How many total electrons can be in a sublevel?
Remember: A maximum of two electrons can
be placed in an orbital.
s orbitals p orbitals d orbitals f orbitals
Number of
orbitals
Number of
electrons
Lithium
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Group 1A
Atomic number = 3
1s22s1 ---> 3 total electrons
3p
3s
2p
2s
1s
Carbon
3p
3s
2p
2s
1s
Group 4A
Atomic number = 6
1s2 2s2 2p2 --->
6 total electrons
Here we see HUND’S
RULE.
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Let’s Try It!
• Write the electron configuration for
the following elements:
H
He
B
N
F
Ne
Na
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Orbitals and the Periodic Table
• Orbitals grouped in s, p, d, and f orbitals
(sharp, proximal, diffuse, and fundamental)
• The d orbitals are n-1 for Rows 4-7
• The f orbitals are n-2 for Rows 6-7
s orbitals
f orbitals
d orbitals
p orbitals
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Shorthand Notation
• A way of abbreviating long electron
configurations
• Use Noble gases as placeholders
for completely filled levels
• Only write out the configuration for
the Valence (outermost) Level
He Ne Ar Kr Xe Rn
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Shorthand Notation
• Step 1: Find the closest noble gas
to the atom (or ion) that has a
lower atomic number.
• Step 2: Write the symbol for the
Noble Gas in brackets [ ]
• Step 3: Find the energy level
below your Noble Gas
• Step 4: Write the configuration for
that energy level and stop when
you reach your target element
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Shorthand Notation
• Chlorine
– Longhand is 1s2 2s2 2p6 3s2 3p5
You can abbreviate the first 10
electrons with a noble gas,
Neon. [Ne] replaces 1s2 2s2 2p6
The next energy level after Neon
is 3
Write out the configuration for
level 3 until you reach Cl
[Ne] 3s2 3p5
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Practice Shorthand Notation
• Write the shorthand notation for
each of the following atoms:
Br
K
Sb
Zn
Sn
Np
Valence Electrons
Electrons are divided between core and
valence electrons
B 1s2 2s2 2p1
Core = [He] , valence = 2s2 2p1
Br [Ar] 3d10 4s2 4p5
Core = [Ar] 3d10 , valence = 4s2 4p5
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Electron Configurations for Ions
• Remember…
– Negative ions have gained electrons
– Positive ions have lost electrons
• Electrons that are lost or gained should
be added/removed from the highest
energy level (not the highest orbital in
energy!)
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Electron Configurations for Ions
Tin (Sn)
Atom: [Kr] 5s2 4d10 5p2
Sn+4 ion: [Kr] 4d10
Sn+2 ion: [Kr] 5s2 4d10
Note that the electrons came out of
the highest energy level, not the
highest energy orbital!
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Electron Configurations for Ions
Bromine
Atom: [Ar] 4s2 3d10 4p5
Br- ion: [Ar] 4s2 3d10 4p6
Note that the electrons went into
the highest energy level, not the
highest energy orbital!
Ion Configurations
To form anions from elements, add 1 or more
e- from the highest sublevel.
P [Ne] 3s2 3p3 + 3e- ---> P3- [Ne] 3s2 3p6 or [Ar]
3p
3p
3s
3s
2p
2p
2s
2s
1s
1s
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Try Some Ions!
• Write the longhand notation for these:
FLi+
Mg+2
• Write the shorthand notation for these:
O-2
Ba+2
Al+3
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(HONORS only)
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Exceptions to the Aufbau Principle
• Remember d and f orbitals require LARGE
amounts of energy
• If we can’t fill these sublevels, then the next
best thing is to be HALF full (one electron in
each orbital in the sublevel)
• There are many exceptions, but the most
common ones are the elements whose
configurations end in
d4 – Cr, Mo, W
d9 – Cu, Ag, Au
(HONORS only)
Exceptions to the Aufbau Principle
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d4 is one electron short of being HALF full
In order to become more stable (require
less energy), one of the closest s
electrons will actually go into the d,
making it d5 instead of d4.
Example:
Cr (theoretical) [Ar] 4s2 3d4
Cr (actual)
[Ar] 4s1 3d5.
Procedure: Find the closest s orbital. Steal
one electron from it, and add it to the d.
(HONORS only)
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Exceptions to the Aufbau Principle
OK, so this helps the d, but what about the
poor s orbital that loses an electron?
Remember, half full is good… and when an
s loses 1, it too becomes half full!
So… having the s half full and the d half full
is usually lower in energy than having the
s full and the d to have one empty orbital.
(HONORS only)
Exceptions to the Aufbau Principle
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d9 is one electron short of being full
Just like d4, one of the closest s electrons
will go into the d, this time making it d10
instead of d9.
Example: Au (theoretical) [Xe] 6s2 4f14 5d9
Au (actual)
[Xe] 6s1 4f14 5d10.
Procedure: Same as before! Find the
closest s orbital. Steal one electron from
it, and add it to the d.
(HONORS only)
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Try These!
• Write the shorthand
notation for:
Cu
W
Au
General Periodic Trends
• Atomic and ionic size
• Ionization energy
• Electronegativity
Higher effective nuclear charge
Electrons held more tightly
Larger orbitals.
Electrons held less
tightly.
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Atomic Size
• Size goes UP on going down a group.
• Because electrons are added further
from the nucleus, there is less
attraction. This is due to additional
energy levels and the shielding effect.
Each additional energy level “shields”
the electrons from being pulled in
toward the nucleus.
• Size goes DOWN on going across a
period.
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Atomic Size
Size decreases across a period owing
to increase in the positive charge from
the protons. Each added electron feels
a greater and greater + charge because
the protons are pulling in the same
direction, where the electrons are
scattered.
Large
Small
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Which is Bigger?
• Na or K ?
• Na or Mg ?
• Al or I ?
Ion Sizes
Li,152 pm
3e and 3p
Does+ the size go
up+ or down
Li , 60 pm
when
losing
an
2e and 3
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.
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.
• Trends in ion sizes are the same as
atom sizes.
Trends in Ion Sizes
Figure 8.13
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Which is Bigger?
• Cl or Cl- ?
• K+ or K ?
• Ca or Ca+2 ?
• I- or Br- ?
Ionization Energy
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IE = energy required to remove an electron
from an atom (in the gas phase).
Mg (g) + 738 kJ ---> Mg+ (g) + e-
This is called the FIRST
ionization energy because
we removed only the
OUTERMOST electron
Mg+ (g) + 1451 kJ ---> Mg2+ (g) + eThis is the SECOND IE.
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
(Shielding Effect)
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Which has a higher 1st
ionization energy?
• Mg or Ca ?
• Al or S ?
• Cs or Ba ?
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Electronegativity, 
 is a measure of the ability of an atom
in a molecule to attract electrons to
itself.
Concept proposed by
Linus Pauling
1901-1994
Periodic Trends:
Electronegativity
• In a group: Atoms with fewer
energy levels can attract electrons
better (less shielding). So,
electronegativity increases UP a
group of elements.
• In a period: More protons, while
the energy levels are the same,
means atoms can better attract
electrons. So, electronegativity
increases RIGHT in a period of
elements.
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Electronegativity
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Which is more electronegative?
• F or Cl ?
• Na or K ?
• Sn or I ?