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Unit 3: Electron Arrangement &
The Periodic Table
This chapter will introduce the
chemistry needed to understand
Section 1: History of Atomic Theory & The Bohr
Model
Section 2: Electron Configuration
Section 3: Orbital Notation
Section 4 Electron Configuration for Ions
Section 5: Light
Section 6: The Periodic Table
Section 7: Trends
REMEMBER…..
The atom is defined as the
smallest particle of an element
that retains the properties of that
element.
Section 4.0—The History of
Atomic Theory
The Beginning of the “ATOM”
In 400 B.C., Democritus, a Greek
philosopher expressed his idea that matter
was made of very small, indivisible
particles that he named “ATOMOS”
Next in line: John Dalton
 John Dalton (17661844), an English
school teacher and
chemist, studied the
results of experiments
by other scientists.
Dalton’s Atomic Theory : in 1803
1. All matter is made of tiny indivisible particles
“atoms”
2. Atoms of the same element are identical; those
of different atoms are different.
3. Atoms of different elements combine in wholenumber ratios to form compounds
4. Chemical reactions involve the rearrangement
of atoms. Atoms cannot be created, divided,
destroyed or changed into other types of atoms
Dalton’s Atomic Model: tiny solid
ball that could not be broken up
into parts.
What was wrong with Dalton’s
Theory? Theories change
 The discovery of the subatomic particles led to
the idea that the atom is not “Indivisible”
 Atoms of the same element can have different
masses
Next : J.J. Thomson & Cathode Ray
Tubes
A cathode ray is a ray of light traveling in a vacuum (no other particles inside)
The ray travels from one metal plate to another as the plates are connected to
electricity
Cathode ray
Metal plate (cathode)
releases stream
Metal plate (anode) to
which stream travels
Cathode Ray Tubes & Charge
In the late 1800’s, JJ Thomson put charged plates outside the tube
Negatively charged plate
-
+
Positively charged plate
Ray is deflected
away from negative
plate and towards
positive plate
It made no difference what type of metal he used in the tube—all material produced
this stream that curved towards the positive charge
Discovery of Subatomic Particles
1897
Thompson discovers the electron through the
use of the cathode ray tube.
 He knows it is negatively charged and has an
extremely small mass.
Cathode Ray Movie Clip
http://www.youtube.com/watch?v=O9Goys
cbazk
New Model Proposed
 Goldstein determines there are positive particles
called protons and hey have a mass 1837 times
heavier than the electron!
 Thomson’s evidence showed Dalton’s idea of
solid, uniform atoms was incorrect.
 Thomson developed the “plum pudding” model.
Since most of us aren’t familiar with plum pudding, you
can think of it as a chocolate cookie dough theory
Plum Pudding Model
A.K.A.
Cookie Dough Model
The “chips” are the negative
electrons.
The “dough” is the positive
portion
The “chips” are stationary and
don’t move within the “dough”
Remember, officially this theory
is called “plum pudding” but it’s
the same idea!
Next: Rutherford & Gold Foil
Experiment: Early 1900’s
 He led a team of
scientists
 They bombarded very
thin gold foil with
radioactive particles
(alpha particles “”)
 They expected these
relatively heavy
particles to go
through the atoms
with a small deflection
Gold Foil Experiment
Basic video
http://www.youtube.com/watch?v=5pZj0u_X
Mbc
Relates to the models
http://www.youtube.com/watch?v=XBqHkraf
8iE
What happened in the Experiment?
What did he see?
Most of the alpha particles passed straight
through with no deflection
These particles did not run into anything
Some did deflect slightly
These particles ran into something much
smaller than themselves
A few were reflected back the direction
they came from
These particles ran into something very dense
Rutherford discovers the nucleus
through the gold foil experiment &
that atoms are mostly empty space.
 Electrons (small
particles) caused the
small deflections
 A small area of the
atom with most of its
mass (the protons)
that caused the
reflections.
Rutherford’s Model
Sphere with
dense,
middle
center called
the nucleus
with
electrons
dispersed
around it
1932: Chadwick: A third particle
The protons and electrons could explain
the charges of the various parts of the
atom
They could not explain the total mass of
the atoms
Chadwick confirms the neutron which has
a mass similar to a protonand no charge.
They were located in the nucleus
Section 1: Next: Niels Bohr:1913
Bohr performed experiments with
hydrogen atoms & light
He determined that electrons are in levels
according to how much energy they have
and that only certain energy amounts were
allowed.
Think of Energy levels as rungs of a
ladder.
The Bohr Model: 1913
The farther away an energy level is from
the nucleus, the more energy it attains.
The circle closest to the nucleus contains
the lowest energy electrons
Energy Levels
Electrons can move from one
energy level to the next by
gaining or losing energy
(quanta).
Ground State: An electron is
as close to the nucleus as it
can get.
“excite
d state”
Excited State: An electron in a
higher energy level than it
should be.
“ground
state”
Drawing Simple Bohr Models
 When assigning electrons, a max of 2 electrons
are placed in the first shell, up to 8 in the 2nd
shell, up to 18 in the 3rd shell, etc.
 Only 8 electrons can be placed in the 3rd shell at
first, then 2 electrons will move into the 4th shell
and the remaining of the 18 will be placed back
in the 3rd shell for a total of 18. **(Just know this : it
will be explained in the next section)**
 Valence electrons are the outermost electrons
is found in the highest, outermost energy level of
the atom
Bohr Atomic Model
http://www.youtube.com/watch?v=ZkNPBy
2i0ss
http://www.youtube.com/watch?v=Cn6v5y
gyZHQ
Lewis Dot Diagrams
 A way to show the # of valence electrons in an
atom.
 The symbol represents the nucleus and inner
core electrons. The dots represent the valence
electrons.
 Only 2 dots per side. Only 4 sides. Max. of 8
valence electrons.
 Add dots one at a time on each side until each
side is full. Then add a second dot to make a
pair when needed. Exception is Helium. It can
look like this: He: or

He 
Bohr Model In and then Out!
We no longer believe electrons are in
concentric circles, but this is still a
convenient way to show energy levels on
2-dimensional paper
1920: At Last: Modern Atomic
Theory: Quantum Mechanical Model
Modern atomic theory uses calculus
equations to show how the subatomic
electrons act as both particles and waves
These equations show the most probable
location of electrons in the atom (known as
atomic orbitals)
Quantum Mechanical Model
Section 2—Electron Structure
The Electron Hotel Analogy
Shopping Center
Parking Garage
Restaurant
A man built an hotel for electrons with a restaurant next door.
But he was making so much money that he decided to add on with some more
rooms and a parking garage.
He still had high demand and decided to add on some more rooms and a
shopping center.
He used the last space he could to put some rooms above the shopping center.
How the Electron Hotel Fills
Shopping Center
Parking Garage
Restaurant
This man had some very strange ideas about how to run his hotel. He insisted
four things:
• The lowest possible must be used first (actually it was the fire inspector
that insisted on this one)
• There can only be one person in a room until all rooms at that level
have someone
• No more than 2 people to a room
• When two people are in a room, they must be of opposite sex
If 8 people come to the hotel, where would he put them?
Another Example
Shopping Center
Parking Garage
Restaurant
This man had some very strange ideas about how to run his hotel. He insisted
four things:
• The lowest possible must be used first (actually it was the fire inspector
that insisted on this one)
• There can only be one person in a room until all rooms at that level
have someone
• No more than 2 people to a room
• When two people are in a room, they must be of opposite sex
If 21 people come to the hotel, where would he put them?
You Try
Shopping Center
Parking Garage
Restaurant
This man had some very strange ideas about how to run his hotel. He insisted
four things:
• The lowest possible must be used first (actually it was the fire inspector
that insisted on this one)
• There can only be one person in a room until all rooms at that level
have someone
• No more than 2 people to a room
• When two people are in a room, they must be of opposite sex
If 42 people come to the hotel, where would he put them?
Where do electrons really live?
They don’t live in a hotel…They are in
the area outside of the nucleus, the
ELECTRON CLOUD
Electron Clouds can be broken into:
LEVEL = SHELL
Electron Hotel
Electron
cloud
Which section
of the hotel
Principal
energy levels
The electron cloud is
made of energy levels.
Which floor
Sublevels
Energy levels are
composed of sublevels
Which room
Orbitals
Sublevels have orbitals.
Energy Levels
There are 7 energy levels.
Energy levels are also called shells
The period number on the periodic table
corresponds to the energy level
Example: The energy level of the furthest
electrons is:
Ca is in energy level 4
Cl is in energy level 3
Sublevels (subshells): a set of
orbitals with equal energy
4 sublevels exist
“S” subshell
Spherical shaped
Only 1 orientation (position)= 1 orbital
Maximum of 2 electrons
Represented on the periodic table as
groups 1A and 2A + helium
First seen in the 1st energy level
“P” Subshell
Dumbbell shaped
There are 3 orientations (positions)= 3 orbitals
Maximum of 6 electrons
Represented on the periodic table as groups
3A -8A
First seen in the 2nd energy level
“D” Subshell
Four lobed shaped
There are 5 orientations (positions)= 5 orbitals
Maximum of 10 electrons
Represented on the periodic table as the
transition metals (group B)
First seen in the 3rd energy level
“F” Subshell
Too complex to name shaped
There are 7 orientations (positions)= 7 orbitals
Maximum of 14 electrons
Represented on the periodic table as the
inner transition metals (lower block)
First seen in the 4th energy level
Within a Sublevel are Orbitals
Orbital– Area of high probability of the
electron being located.
Each orbital can hold 2 electrons
 To calculate the total number of orbitals,
use 2(n)2
Summary
Electron Configuration
Is an address of an electron
Electrons must be placed in
the lowest possible energy levels first (ground state)
4p
1
Energy Level
# of electrons
subshell
3 rules that govern electron
configurations
Aufbau Principle: Electrons must fill the lowest
1
available subshells and orbitals before moving on
to the next higher energy subshell/orbital.
Filling order is:
1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f
5d 6p
Can use the periodic table as a guide to fill from lowest to
highest sublevels or memorize filling order
An orbital’c cloud space can overlap within different energy
levels
The 4s is lower in energy than the 3d
Energy and Subshells-Filling Order
6p
6s
5p
5d
4f
4d
5s
4p
3d
4s
3p
3s
2p
Energy
2s
Subshells are filled from
the lowest energy level to
increasing energy levels.
Does this look familiar? Electron Hotel!
1s
How Electron Configurations Relate to the
Organization of the Periodic Table
s
p
d
f
Figure 11.31: Orbitals being filled for elements in various parts of the periodic
table.
Hund’s Rule
2
Hund’s Rule: Place electrons in unoccupied
orbitals of the same energy level before doubling
up.
How does this work?
If you need to add 3 electrons to a p subshell, add 1 to each before
beginning to double up.
Pauli Exclusion Principle
3
Pauli Exclusion Principle: Two electrons that
occupy the same orbital must have opposite spins.
“Spin” is designated with an up
or down arrow.
How does this work?
If you need to add 4 electrons to a p subshell, you’ll need to double
up. When you double up, make them opposite spins.
Electron Configurations
1
Determine the number of electrons to place
2
Follow Aufbau Principle for filling order
3
Fill in subshells until they reach their max (s = 2, p = 6, d = 10,
f = 14) or use periodic table as a guide.
4
The total of all the superscripts is equal to the number of
electrons.
Example:
Write
electron
configuration
for S
No charge written  Charge is 0
Atomic number for S = 16 = # of protons
0 = 16 - electrons
1s 2 2s 2 2p 6 3s 2 3p 4
Electrons = 16
2 + 2 + 6 + 2 + 4 = 16
Electron Configurations
Example:
Write
electron
configuration
for K
No charge written  Charge is 0
Atomic number for K = 19 = # of protons
0 = 19 - electrons
Electrons = 19
1s 2 2s 2 2p 6 3s 2 3p6 4s1
2 + 2 + 6 + 2 + 6 + 1 = 19
Electron Configurations
Example:
Write
electron
configuration
for Ti
No charge written  Charge is 0
Atomic number for Ti = 22 = # of protons
0 = 22 - electrons
Electrons = 22
1s 2 2s 2 2p 6 3s 2 3p6 4s2 3d2
2 + 2 + 6 + 2 + 6 + 2 + 2 = 22
Unit 3: Other Notations
What is Orbital Notation?
It shows the grouping and position of
electrons in an atom.
Orbital Notation use boxes or lines for
orbitals and arrows for electrons.
Drawing Orbital Notation (boxes &
arrows)
1
Aufbau Principle: Electrons fill subshells (and orbitals) so that
the total energy of atom is the minimum
2
Hund’s Rule: Place electrons in unoccupied orbitals of the
same energy level before doubling up.
3
Pauli Exclusion Principle: Two electrons that occupy the same
orbital must have different spins.
Example:
Write the
orbital
notation for
Cl
Drawing Orbital Notations
1
Aufbau Principle: Electrons fill subshells (and orbitals) so that
the total energy of atom is the minimum
2
Hund’s Rule: Place electrons in unoccupied orbitals of the
same energy level before doubling up.
3
Pauli Exclusion Principle: Two electrons that occupy the same
orbital must have different
3p spins.
No charge written  Charge is 0
Example:
Write the
orbital
notation for
Cl
1s
Atomic number for Cl = 17 = # of protons
0 = 17 - electrons
Electrons = 17
2s
2p
4
9
70
6
5
3
2
1
8
17
16
15
14
13
12
11
3s
3p
Drawing Orbital Notations
No charge written  Charge is 0
Example:
Write the
orbital
notation for
Fe
Atomic number for Fe = 26 = # of protons
0 = 26 - electrons
Electrons = 26
1s
2s
2p
3s
3p
4s
3d
Shorthand Notation
(aka: Noble Gas Notation)
Noble Gas– Group 8 of the Periodic
Table. They contain full valence shells.
The Noble gas is used to represent the core
(inner) electrons and only the valence shell is
shown.
Br
Complete electron configuration
1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 5
Noble gas
[Ar] 4s 2 3d 10 4p 5
The “[Ar]” represents the core electrons and only the valence electrons are shown
Which Noble Gas Do You Choose?
How do you know which noble gas to use to symbolize
the core electrons?
Think: Price is Right.
How do you win on the Price is Right?
By getting as close as possible without going over.
Choose the noble gas that’s closest without going over!
Noble Gas
# of electrons
He
2
Ne
10
Ar
18
Kr
36
Xe
54
Noble Gas Notation Example
1
Determine the number of electrons to place
2
Determine which noble gas to use
3
Start where the noble gas left off and write electron
configuration for the valence electrons
Example:
Write noble
gas notation
for Fe
Noble Gas Notation Example
1
Determine the number of electrons to place
2
Determine which noble gas to use
3
Start where the noble gas left off and write spectroscopic
notation for the valence electrons
No charge written  Charge is 0
Atomic number for Fe= 26= # of protons
Example:
Write noble
gas notation
for Fe
0 = 26- electrons
[Ar] 4s 2 3d 6
Electrons = 26
Closest noble gas: Ar (18)
Ar is full up through 3p
18 + 2 + 6 = 26
Noble Gas Notation Example
Example:
Write noble
gas notation
for Ba
No charge written  Charge is 0
Atomic number for Ba= 56 = # of protons
0 = 56 - electrons
Electrons = 56
Closest noble gas: Xe (54)
Xe is full up through 4p
[Xe] 6s 2
54 + 2 = 56
Section 4: Electron Configuration for
Ions
1
Determine the number of electrons to place. Positive ions lose
electrons; negative ions gain electrons.
2
Follow Aufbau Principle for filling order
3
Fill in subshells until they reach their max (s = 2, p = 6, d = 10,
f = 14) or use periodic table as a guide.
4
The total of all the superscripts is equal to the number of
electrons.
Example:
Write
electron
configuration
for S-2
Atomic number for S = 16 = # of protons
-2 = 16 - electrons
1s 2 2s 2 2p 6 3s 2 3p 6
Electrons = 18
2 + 2 + 6 + 2 + 6 = 18
Example
1
Determine the number of electrons to place. Positive ions lose
electrons; negative ions gain electrons.
2
Follow Aufbau Principle for filling order
3
Fill in subshells until they reach their max (s = 2, p = 6, d = 10,
f = 14) or use periodic table as a guide.
4
The total of all the superscripts is equal to the number of
electrons.
Example:
Write
electron
configuration
for K+1
Atomic number for K = 19 = # of protons
+1 = 19 - electrons
1s 2 2s 2 2p 6 3s 2 3p 6
Electrons = 18
2 + 2 + 6 + 2 + 6 = 18
Another Example
1
Determine the number of electrons to place. Positive ions lose
electrons; negative ions gain electrons.
2
Follow Aufbau Principle for filling order
3
Fill in subshells until they reach their max (s = 2, p = 6, d = 10,
f = 14) or use periodic table as a guide.
4
The total of all the superscripts is equal to the number of
electrons.
Example:
Write
electron
configuration
for Fe+2
Atomic number for S = 26 = # of protons
+2 = 26 - electrons
6
1s 2 2s 2 2p 6 3s 2 3p 6 3d
Electrons = 24
2 + 2 + 6 + 2 + 6 + 6 = 18
Exceptional Configuration
Half-filled or completely filled d & f sublevels have
LOWER energies and are more stable than partially filled
d’s and f’s.
This means that an atom can “borrow” one of its “s”
electrons from the previous orbital to become more stable.
___
↑↓
__
↑↓
__ __
__
__↑
↑↓
↑↓
↑↓
5s
4d
becomes
↑
___
5s
↑↓
__
↑↓
__
↑↓
__
↓↑
__
↑↓
__
4d
Because the 4d sublevel is now full, the atom is
at a lower energy state and therefore more
stable.
YOU TRY: Write the exceptional configuration for
copper, Cu
This means that an atom can “borrow” one of its “s”
electrons from the previous orbital to become more stable.
↑↓
___
↑__
↓
4s
↑ ↓ __
↑ ↓ __
↑
↑ ↓ __
__
3d
becomes
↑
___
4s
↑↓
__
↑ ↓ __
↑↓
__
↓↑
__
↑↓
__
3d
Because the 4d sublevel is now full, the atom is
at a lower energy state and therefore more
stable.
Section 5—Light
Light is Electromagnetic Radiation
Electromagnetic radiation is all energy that
travels through space at the speed of light
which equals 3.0 x 108 m/s
Examples of Electromagnetic
Radiation…visible, microwaves, infrared,
ultraviolet, radio waves, gamma, x-rays
ElectroMagnetic Radiation Spectrum
Wave Properties—Wavelength
Wavelength () is the distance from trough
to trough of a wave (measured in meters
“m”)
1nm = 1 x 10-9 m
wavelength
Wave Properties—Frequency
Frequency () is the number of times a
wave completes a cycle in one second
(cycles per second is “Hertz” or “Hz”)
Lower frequency
Higher frequency
Reference Sheet
Lower
frequency
Lower Energy
Higher frequency
Higher Energy
Relationship between wave properties
Notice the shorter the wavelength, the
higher the frequency. This is an INVERSE
relationship.
c = 
Notice as frequency increases, the amount
of energy increases. This is a DIRECT
relationship.
Visible Light
The visible light region is a small region
within the spectrum that has
wavelengths/frequencies that are eyes can
detect.
ROY G BIV is a mnemonic to help you
remember the colors of the spectrum.
Red is near infrared and violet is near ultra
violet.
Reference Sheet
Examples Using Reference Sheet
1.
Which of the following forms of electromagnetic radiation has the
shortest wavelength?
a) gamma
b) visible
c) infrared
d) radio
As the frequency of electromagnetic radiation increases, its
wavelength ______________.
a) increases b) decreases c) remains constant d) is impossible to
determine.
3. Which of the following forms of radiation has photons with greatest
amount of energy?
a) red light
b) yellow light
c) green light
d) violet light
A
2.
B
D
Interesting superhero facts:
• Superman has x-ray vision.
• The Incredible Hulk was “created” by an
accidental overdose of gamma radiation.
• The Fantastic Four were “created”
•by cosmic rays.
Bohr Model
4. What type of electromagnetic
radiation is represented by a
wavelength of 1870 nm?
a)Infrared
light
c) Ultraviolet
A
b) visible
d) x-ray
Examples Using Reference Sheet
C
5. What type of electromagnetic radiation is represented by a
wavelength of 4.7x10-1m?
a) gamma rays
b)infrared
c) microwaves d) visible light
LIGHT & MATTER:
Visible Range
Wavelength increases
Frequency decreases
Energy decreases
400 nm
700 nm
Visible light
White light is made of all the colors…a prism can separate white light into
a rainbow!
Continuous Spectrum: Sun light (or white light) will
produce a range of color because there are no
specific wavelengths
Line Spectrum
Is when individual atoms emit light of only
certain wavelengths. Each element has its
own line spectrum, or fingerprint.
How can a line spectrum be explained?
Electrons Absorbing Energy
Energy packets called photons or quanta come
into contact with an atom & collide with an
electron.
+
The electron is “excited” to a higher energy level with is
newly increased energy from absorbing the photon.
Electrons Absorbing Energy
Photon coming into atom collides with
electron. Photons are energy.
+
Excitation
The process of an electron absorbing a
photon of light (energy) and being
promoted to a higher energy level from its
“ground state”
And later…
The electron cannot remain in that excited state indefinitely
+
And later…
The electron cannot remain in that excited state indefinitely
+
Energy is released during relaxation
Relaxation
The process of an electron releasing a
photon of light (energy) and falling back
down to a lower energy level.
Energy of photon and levels jumped
The higher the energy of the photon, the
greater the electron jump!
A photon of UV light has more energy than
a photon of Infrared light
The UV photon would cause a higher energy
jump (jump up more levels) than the IR photon.
Total energy in = Total energy out
However much energy was absorbed must
be released again, but it can be released
in smaller packets
A high energy photon might be absorbed,
but two lower energy photons might be
released as the electron falls in a “stepwise” manner.
Photons must match energy changes
The energy of the photon must exactly
match the energy change of the electron.
If the photon is not an exact match, the
photon will pass through unabsorbed.
+
Hydrogen Line Spectrum
The colored lines are the wavelengths of
light that are emitted when an electron
moves from a higher E level to a lower E
level,
This was proof that atoms had fixed
energy levels!
Emission Spectrum
Hydrogen
Spectrum
Neon
Spectrum
How hydrogen
produces the
four visible
photons
Reference Sheet
Examples
 1. On the energy level diagram below, draw an arrow
representing the electron in hydrogen’s ground state
being excited to the fourth energy level.
Examples
2. An
electron in the hydrogen atom makes the transition
n = 5  n = 3.
a. Determine the wavelength of light associated with this transition.
Include units.
a) 434 nm b) 434 m
A
c) 1282 nm
b. Classify the type of electromagnetic
radiation this wavelength represents:
a) infrared
b) visible light
D
c) ultraviolet
d) x-ray
c. Is this energy emitted by the atom or
absorbed by the atom ?
emitted
d) 1282 m
Flame Tests
Metals can be identified by the wavelength
of light they emit. When metals absorb
energy from a flame, the electrons absorb
energy and are raised to higher energy
level.
When they return to their ground state,
they release the energy they absorbed in
the form of radiation. The wavelength of
light for some metals fall in the visible light
portion of the spectrum. This allows us to
see their color.
Ways of producing light
Fluorescence: visible light is absorbed and
visible light is emitted at the same time—
the relaxation happens very quickly after
excitation
Phosphorescence: Visible light is
absorbed and then a while later is
emitted—relaxation occurs after a period
of time
Section 6—The Periodic Table
History of the Periodic Table
Different scientists organized the
elements differently—this lead to
confusion
In 1869, Dimitri Mendeleev designed a
periodic table based on atomic mass.
This way showed patterns in properties
that repeated across rows and similarities
down columns
He couldn’t find elements to fit all the
property trends, so he left holes
Mendeleev’s Periodic Table
History of the Periodic Table
 The holes Mendeleev left were later
filled in as more elements were
discovered
 The modern periodic table is arranged
by atomic number rather than atomic
mass.
This caused a few “switches” in placement, but
overall is very similar to Mendeleev’s
 Henry Mosley is given credit for the modern
periodic table.
Modern Periodic Law: Elements in columns
have similar properties!
Organization: Groups and Periods
Periods
Rows are called
“periods”
Groups
Columns are called
“groups” or “families”
Information for Each Element
Most periodic tables give the following information,
but it can be in a different location
Atomic Number
Element Symbol
If there’s a second
letter, it’s lower-case
6
C
Carbon
12.01
Average Atomic Mass
Number with decimals
Gives the mass for 1 mole of atoms, in
grams
THIS IS NOT THE MASS NUMBER!
Whole number—
elements are ordered
by this on the periodic
table.
Element Name
Modern Periodic Table Divisions
The rows at the bottom
Most periodic tables are written with 2 rows at the bottom.
This is done to allow the font to be bigger on a piece of paper.
The rows at the bottom
Most periodic tables are written with 2 rows at the bottom.
This is done to allow the font to be bigger on a piece of paper.
But they really belong here!
Follow the atomic numbers on your periodic table to see it!
Properties of Metals
 All solid except for mercury
 Formability
Malleable: can be flattened
into thin sheets
Ductile: can be drawn into fine wire
 Great conductors of heat & electricity
 High luster
 High strength
Properties of Nonmetals
Mostly gases, few solids and 1 liquid, Br
Poor formability
Poor conductors of heat & electricity
dull
brittle
Sulfur
Properties of Metalloids(semi-metals)
All solids
Share properties of both metals &
nonmetals
Semiconductors
Electron Configuration Patterns
Look at the electron configurations for the Halogens
F
1s2 2s22p5
Cl
1s2 2s2 2p6 3s23p5
Br
1s2 2s2 2p6 3s2 3p6 4s2 3d104p5
I
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d105p5
All of the elements in Group 7 end with 5 electrons in a p subshell.
Every Group ends with the same number of electrons in the
highest energy subshell. (Similar Valence electrons of a group
explains why elements in the same column have similar
properties.)
Configurations and the Periodic Table
s-block
p-block
d-block
s1 s2
p1 p2 p3 p4 p5 p6
d1 d2 d3 d4 d5 d6 d7 d8 d9 d10
f-block
f1
f2
f3
f4
f5
f6
f7
f8
f9 f10 f11 f12 f13 f14
2 Types of Elements
 Representative elements(main group): Group A
 See special families info. In the next set of slides
 Transition Elements: Group B
Properties
 Have more than 1
oxidation number or ion charge
 Typically the least reactive of the metals
 Hardest, densest, and highest melting points
 Greatest conductors of heat and electricity
Parts of the Periodic Table
Parts of the Periodic Table
Family: Alkali Metals
 Group 1A
 Most reactive metals; always
found in compound form
 Reacts violently with water to
produce hydrogen gas and a
base
 Soft and silver in color
 Low density & melting points
 Has 1 valence electron and
forms +1 ions
Cool Clip Showing Reactivity of Alkali
Metals
http://www.youtube.com/watch?v=m55kgy
ApYrY
Family: Alkaline Earth
Metals
Group 2A
 less reactive than 1 A but
still extremely reactive; always found in
compound form
Reacts with water to produce hydrogen gas
and a base – Not Explosive
Extracted from ores-mineral rocks
Harder,higher density & melting points than
1A
Has 2 valence electrons and forms +2 ions
Family: Halogens
Group 7A
 all are non metals
most reactive nonmetals;
always found in compound
form
Most form the diatomic
molecules
Has 7 valence electrons and
forms -1 ions
Family: Noble Gases
Group 8A
All nonmetals
 exist as gases
nonreactive elements- most stable
Called inert gases
All have 8 valence electrons except
Helium which only has 2 and does not
form ions
Other Important Facts
Hydrogen is a diatomic non metal & in a
group by itself; can have a +1 or -1 charge
58-71 Lanthanides: 1st row of lower block;
called Rare-Earth metals
89-103 Actinides: 2nd row of lower block;
all are radioactive and most are manmade
Section 7—Trends &Periodicity
What is periodicity on the periodic table?
The predictable pattern by which
properties of elements change across
or down the periodic table.
There are always exceptions to these periodicity trends…each of
the trends is a “general” trend as you move across a period or
down a group.
Three Important Terms to help explain
TRENDS
FOR A PERIOD TREND ONLY: “NUCLEAR
CHARGE”: the number of protons in the nucleus;
determines the strength of attraction to the electrons
FOR A GROUP TREND ONLY: “ENERGY LEVEL”:
location of the electrons; the higher the energy level,
the farther distance the electrons are from the nucleus
FOR A GROUP TREND ONLY:”SHIELDING” : energy
levels that are between the nucleus and valence
electrons cause this
Trend 1 :Atomic Radius
Defined as: Half of the distance between the
nuclei of two bonded atoms.
H
H
Distance between nuclei
Atomic radius of hydrogen atom
Atomic Radii Trends
Period trend
:Decreases
Group
trend:
Increase
s
Atomic Radii Period Trend: DECRESES
WHY?
As the number of protons increases, the nuclear
charge increases.
e
e
e
n
p
n p
n p
e
Lithium atom
Move across the
periodic table
Radius decreases
e
e
pn
p p
p
n nn
e
Beryllium atom
As the nuclear charge increases, the attraction
between the positive nucleus and negative electron
cloud increases.
This attraction “pulls” in on the electrons.
Atomic Radii: Group Trend: INCREASES
WHY?
Nuclear Charge also increases as you go down a group.
HOWEVER: The electrons are added into new energy levels.
e
e
e
+
Move down the
periodic table
e
e
Radius increases
e
e
e
e
e
e
e
Lithium atom
e
+
e
Sodium atom
The inner electrons “shield” the new outer electrons from the
pull of the nucleus, therefore it doesn’t pull in like the last slide.
Examples
 List the following in
INCREASING order of atomic
radius .
Li, Cs, K
 List the following in
DECREASING order of atomic
radius .
Ca, Be, Ba, Sr
 Who has the largest atomic
radius?
Mg, Cl, Na, P
Li < K < Cs
Ba > Sr > Ca > Be
Na
Trend 2: Ionization Energy
Defined as: The energy needed to remove the
outermost electron.
Ionization Energy Trends
Period Trend: Increases
Group
Trend:
Decreases
Ionization Energy Period Trend: INCREASES
WHY?
Moving left to right, the size of the atom decreases as more
protons pull harder on more electrons.
Move across the
periodic table
e
e
n
p
n p
n p
e
e
Radius decreases
IE increases
Lithium atom
e
e
pn
p p
p
n nn
e
Beryllium atom
When an atom is smaller, the electrons are closer to the nucleus,
and therefore feel the pull more strongly.
Thus, it requires more energy to remove electrons away from these
smaller atoms.
Ionization Energy Group Trend: DECREASES
WHY?
As you move down a group, the radius increases as more
electrons shells are added.
e
e
e
+
Move down the
periodic table
Radius increases
IE decreases
e
e
e
e
e
e
e
e
e
Lithium atom
e
+
e
Sodium atom
As the outer electrons (those involved in bonding) are farther from
the nucleus, they will feel the “pull” of the nucleus less.
It is easier to remove an electron from a larger atom thus requires
less energy.
Examples
 List the following in
INCREASING order of
ionization energy .
Li, Cs, K
 List the following in
DECREASING order of
ionization energy .
Ca, Ba, Be, Sr
 Who has the highest
ionization energy?
Cl, Na, I, In
Cs < K < Li
Be > Ca >Sr > Ba
Cl
Trends in the Periodic Table
“Successive Ionization Energies”
“Successive Ionization Energies” means the energy required
to remove a 2ndor a 3rdelectron from an atom.
Removing more and more e-’s requires more& more
energy.
Why?
nucleus.
The remaining e-’s are more tightly bound tothe
Trend 3: Electronegativity
Defined as: Tendency of an atom to steal an
electron when combining with another element
F is the highest
Fr is the lowest
Electronegativity Trends
Period trend:
Increases
Group
Trend:
Decreases
Electronegativity Period Trend: INCREASES
WHY?
Moving left to right, the radius of the atom decreases as more
protons pull on more electrons.
Move across the
periodic table
e
e
n
p
n p
n p
e
e
Radius decreases
EN increases
Lithium atom
e
e
pn
p p
p
n nn
e
Beryllium atom
When an atom is smaller, the nucleus pulls more strongly.
This can attract & draw an electron away from a different
atom.
Electronegativity Group Trend: DECREASES
WHY?
As you move down a group, the radius increases as more
electrons shells are added.
e
e
e
+
Move down the
periodic table
Radius increases
e
e
e
e
e
e
EN decreases
e
+
e
e
e
Lithium atom
Sodium atom
The larger atom has weaker attractions so it cannot draw
electrons away very easily from different atoms.
e
Examples
Who has the highest
electronegativity?
Ba, Br, Ca
Br
List the following in
DECREASING order of
electronegativity.
I, Cl, Br
Cl > Br > I
Trend 4: Ionic Radius
Review Some Definitions
An Ion is defined as– atom with a charge.
Cation– are positively charged ions. Results
from a metal losingelectrons.
Anion– are negatively charged ions. Results
from a nonmetals gaining electrons.
Cations Are Smaller than a Neutral Atom
WHY?
Atoms lose electrons to create positive ions
e
e
e
+
Lithium atom
Creating a cation,
losing electrons
Radius decreases
e
e
+
Li+ ion
When electrons are lost, there are now more protons than
electrons
Therefore, the protons have a greater “pull” on each of the
electrons – making a smaller ion.
Trends in the Periodic Table
Atomic Size vs. Ion Size
Anions are larger than a neutral atom
WHY?
Atoms gain electrons to create negative ions
e
e
e
e
Creating an anion,
gaining electrons
e
+
e
e
e
e
e
e
e
e
+
Radius increases
e
e
e
Oxygen atom
e
e
O2- ion
When electrons are gained, there are now more electrons
than protons
Therefore, the protons have a weaker “pull” on each of the
electrons. They loosen up & spread out making a larger ion
Ionic Radii
Period Trend for METALS & NONMETALS: decreases
Group
Trend:
Increases
Examples
Arrange in order of
decreasing ionic
radius?
P -3 Mg +2 Cl-1
Arrange in order of
INCREASING ionic
size
K+1 Cs+1 Li +1
P -3 > Cl-1 > Mg+2
Li +1 < K+1 < Cs+1
Trend 5: Reactivity
Defined as: How chemically active an atom
is to another
Metal Reactivity Trends
Period Trend Decreases
Group
Trend
Increases
METAL Reactivity Trend: DECREASES
WHY?
Move across the
periodic table
e
e
n
p
n p
n p
e
e
Radius decreases
IE increases
Lithium atom
e
e
pn
p p
p
n nn
e
Beryllium atom
The most reactive metals have a low ionization
energy.
Nonmetal Reactivity Trends
Period Trend Increases
Group
Trend
decreases
NON Metal Reactivity Trend: INCREASES
WHY?
Move across the
periodic table
e
e
n
p
n p
n p
e
e
EN increases
IE increases
Lithium atom
e
e
pn
p p
p
n nn
e
Beryllium atom
The most reactive nonmetals have a high Ionization energy &
electronegativity!