Download Section 4.4*The Periodic Table

Section 4.4—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 he 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 of the Periodic
Table
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
Whole number—
elements are ordered by
this on the periodic table.
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!
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
Sulfur
brittle
Electron Configurations and the
Periodic Table.
Configurations Within a Group
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
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=m55kgyA
pYrY
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
• Lanthanides: 1st row of lower block; called
Rare-Earth metals
• Actinides: 2nd row of lower block; all are
radioactive and most are man-made
Section 4.5—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.
Trend 1: Atomic Radii
What is atomic radius?
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:
Increases
Atomic Radii Period Trend: WHY?
Why do atomic radii decrease across a period?
As the number of protonsincreases, 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: WHY?
Why do atomic radii increase down a group?
Nuclear Charge also increases as you go down a group.
e
e
e
+
Move down the
periodic table
Radius increases
e
e
e
e
e
e
e
e
e
Lithium atom
e
+
e
Sodium atom
HOWEVER: The electrons are added into new energy levels.
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
Li < K < Cs
• List the following in DECREASING
order of atomic radius .
Ba>Sr> Ca > Be
Ca Be BaSr
• Who has the largest atomic
radius?
Mg Cl, Na P
Na
Atomic Radii Trends
Decreases
Increases
Trend 2: Ionization Energy
What is Ionization Energy?
The energy needed to remove the
outermost electron.
Ionization Energy Trends
Period Trend: Increases
Group
Trend:
Decreases
Ionization Energy Period Trend: WHY?
Why does Ionization Energy increase across a period?
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
Lithium atom
e
e
Radius decreases
IE increases
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.
It is harder to pull electrons away from these smaller atoms.
Ionization Energy Group Trend: WHY?
Why does ionization energy decrease down a group?
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
e
e
Lithium atom
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.
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
Ionization Energy Trends
Increases
Decreases
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-’srequiresmore& more energy.
– Why?
The remaining e-’s are more tightly bound tothe
nucleus.
Trend #3: Electronegativity
What is Electronegativity?
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: WHY?
Why does electronegativity increase across a period?
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
Lithium atom
e
e
Radius decreases
EN increases
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 another atom.
Electronegativity Group Trend: WHY?
Why does electronegativity decrease down a group?
As you move down a group, the radius increases as more electrons
shells are added.
e
e
e
+
Move down the
periodic table
e
e
Radius increases
EN decreases
e
e
e
e
e
e
e
Lithium atom
e
+
e
Sodium atom
The larger atom is less able the nucleus is to attract electrons away
from another atom.
Examples
• Who has the highest
electronegativity?
Ba, Br, Ca
Br
• List the following in
DECREASING order of
electronegativity.
I Cl Br
Cl
Electronegativity Trends
Increases
Decreases
Ionic Charge & Radii
Review Some Definitions
Ion– atom with a charge.
Cation– positively charged ion. Results
from a metal losingelectrons.
Anion– negatively charged ion. Results
from a nonmetals gainingelectrons.
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
e
e
+
Radius decreases
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.
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
e
Oxygen atom
O2- ion
e
When electrons are gained, there are now more electrons than
protons
Therefore, the protons have a weaker “pull” on each of the electrons.
Ionic Radii
Period Trend: 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
What is Reactivity?
How chemically active an atom is to
another
Metal Reactivity Trends
Period Trend Decreases
Group
Trend
Increases
METAL Reactivity Trend: WHY?
Why does METAL reactivity decrease across a period or increase
down a group?
Move across the
periodic table
e
e
n
p
n p
n p
e
Lithium atom
e
e
Radius decreases
IE increases
e
pn
p p
p
n nn
e
Beryllium atom
The most reactive metals have the lowest ionization energy.
NonMetal Reactivity Trends
Period Trend Increases
Group
Trend
decreases
NON Metal Reactivity Trend: WHY?
Why does nonmetal reactivity increase across a period or
decrease down a group?
Move across the
periodic table
e
e
n
p
n p
n p
e
Lithium atom
e
e
EN increases
IE increases
e
pn
p p
p
n nn
e
Beryllium atom
The most reactive nonmetals have the highest Ionization energy &
electronegativity!
Section 4.6—Light
Light is Electromagnetic Radiation
• Electromagnetic radiation is all energy that
travels through space in a wave-like manner
• Examples of Electromagnetic
Radiation…visible, microwaves, infared,
ultraviolet, radio waves, gamma, x-rays
• ALL Electromagnetic energy travels at the
speed of light “c” which equals (3.0 × 108 m/s)
ElectroMagnetic Radiation Spectrum
Wave Properties—Wavelength
• Wavelength () is the distance from trough to
trough of a wave (measured in meters “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
Relationship between wave properties
• Notice the shorter the wavelength, the higher
the frequency. This is an INVERSE
relationship.
• 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 infared and violet is near ultra violet.
Reference Sheet
Examples Using Reference Sheet
• 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.
• 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
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
What type of electromagnetic
radiation is represented by a
wavelength of 1870 nm?
a) infrared
b) visible light
c) ultraviolet
A
d) x-ray
Examples Using Reference Sheet
•
C
What type of electromagnetic radiation is represented by a wavelength
of 4.7x10-1m?
a) gamma rays
b)infrared
c) microwaves
d) visible light
Reference Sheet
Lower frequency
Lower Energy
Higher frequency
Higher Energy
Section 4.7—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 “step-wise” 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
c) 1282 nm
d) 1282 m
b. Classify the type of electromagnetic radiation this wavelength represents:
a) infrared
b) visible light
c) ultraviolet
d) x-ray
c. Is this energy emitted by the atom or absorbed by the atom?
__________________
D
A
emitted
Flame Tests
• Metalscan be identified by the wavelength of
light they emit. When metals absorb energy
from a flame, theelectronsabsorb 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 thevisiblelight 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
Ways of producing light
 Incandescence: Energy is put in from heat and given
off as visible light
 Chemiluminescence: Energy released during a
chemical reaction is absorbed to cause excitation.
Relaxation produces visible light
 Biolouminescence: Chemiluminescence that occurs
in a biological organism.
 Triboluminescence: Physical pressure or torque
provides energy for excitation. Relaxation produces
visible light.