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The Periodic Table
chapter 6
How’d They Come Up With
That?
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Our current society takes for granted all of the hard
work, research, chance,
and luck that has gone into creating
and discovering the materials that are used in the
products we utilize every day.
For example, who was the first person to set or find
a random black rock (coal) on fire and discover that
it provided a good, constant source of heat?
Who was the first person to discover that a
substance found in some rocks was capable of being
the ultimate explosive (uranium)?
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In nature and in the lab we have discovered
over 100 different elements.
We’ve organized the elements into a table
based on their PHYSICAL and CHEMICAL
PROPERTIES
It took us almost 2000 years to figure out
the properties of the elements currently in
the Periodic Table of Elements and arrange
them.
Developing the Periodic
Table
By the early 1800s, enough
information was known about the
elements that scientists wanted an
easy way to categorize the Earth’s
ingredients.
 Many methods of organization were
tried before scientists found the
most effective way of grouping the
elements

“Mayan”
Periodic
Table,
named for its
similarity to
the Mayan
calendar.
Johann Dobereiner 1780 - 1849

In 1829, he classified some elements into
groups of three, which he called triads.
The elements in a triad had similar
chemical properties and orderly physical
properties.
(ex. Cl, Br, I and
Ca, Sr, Ba)
Model of triads
John Newlands 1838 - 1898

In 1863, he suggested that elements be
arranged in “octaves” because he noticed
(after arranging the elements in order of
increasing atomic mass) that certain
properties repeated every 8th element.
Law of Octaves
Dmitri Mendeleev
(1834 – 1907)
Russian chemist, Dmitri
Mendeleev organized
elements into a table
based on atomic mass and
similar properties.
 Mendeleev stated that the
properties of elements are
a periodic function of their
atomic masses.

Mendeleev’s Periodic Table
Mendeleev’s Prediction
Mendeleev’s table had several missing
elements. When these elements were
discovered, they were almost exactly
as Mendeleev predicted.
 The following is an example of the
element we know as Germanium.
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Germanium is located below silicon.
Mendeleev predicted its properties
based on this location in his table.
Ekasilicon (Es)
Germanium (Ge)
1. Atomic mass: 72
1. Atomic mass: 72.61
2. High melting pt.
2. Melting pt: 945° C
3. Density: 5.5g/cm3
3. Density: 5.323g/cm3
4. Dark gray metal
4. Gray metal
5. Will obtain from
K2EsF6
5. Obtain from K2GeF6
6. Will form EsO2
6. Forms oxide (GeO2)
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However, in spite of Mendeleev’s great
achievement, problems arose when new
elements were discovered and more
accurate atomic weights determined. By
looking at our modern periodic table, can
you identify what problems might have
caused chemists a headache?
• 18Ar, 39.95 amu and 19K, 39.10 amu
• 27Co, 58.93 amu and 28Ni, 58.69 amu
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Modern Periodic Law
Henry Moseley 1887 - 1915
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In 1913, through his work with X-rays, he
determined the actual nuclear charge
(atomic number) of the elements*. He
rearranged the elements in order of
increasing atomic number.
*“There is in the atom a fundamental quantity which
increases by regular steps as we pass from each
element to the next. This quantity can only be the
charge on the central positive nucleus.”
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Increasing atomic number is the
basis for our current periodic law.
His research was halted when
the British government sent
him to serve as a foot soldier
in WWI. He was killed in the
fighting in Gallipoli by a
sniper’s bullet, at the age of
28. Because of this loss, the
British
government
later
restricted its scientists to
noncombatant duties during
WWII.
Glenn T. Seaborg 1912 - 1999
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After co-discovering 10 new
elements, in 1944 he moved
14 elements out of the main
body of the periodic table to
their current location below
the Lanthanide series. These
became known as the Actinide
series.
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He is the only person to
have an element named
after him while still alive.
106Sg- Seaborgium
"This is the greatest honor
ever bestowed upon me even better, I think, than
winning the Nobel Prize."
Periodic Table
Periodic Table Review:
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Rows on the
periodic table are
called PERIODS
Columns on the
periodic table are
called GROUPS or
FAMILIES
Periodic Table Review
There are 7 periods and 18 groups.
 Electron arrangements are repeated
in periods.
 Elements with similar econfigurations are placed in the same
group.
 Elements in groups are also listed in
order of their increasing principal
quantum numbers.
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Electron Configuration
Sublevel /
s
p
d
f
e- capacity
2
6
10
14
S – block (sublevel)
Contains elements in Group 1,
Group 2, and He from Group 18.
 Electrons are added to the s –
orbitals.
 EX:
H = 1s1
He = 1s2
Li = 1s22s1
Be = 1s22s2
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P – block (sublevel)
Contains elements in Group 13, Group
14, Group 15, Group 16, Group 17, and
the remaining elements from Group 18
(except He)
 Electrons are added to the p – orbitals.
 Ex:
B = 1s22s22p1
C = 1s22s22p2
N = 1s22s22p3
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D – block (sublevel)
Contains elements from the center of
the periodic table.
 These elements are called transition
metals.
 Electrons are added to the d –
orbitals of the transitions metals as
well as La and Ac of the inner
transition elements (rare earth).

F – block (sublevel)
Contains elements from the inner
transition metals (rare earth elements)
 Electrons are added to the f – orbitals.
 Ex:
Ce  Lu
Th  Lr
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Octet Rule
Atoms with full outer levels are stable
(less reactive)
 For elements (except He) this stable
configuration would have eight e-.
(two in the outer s sublevels and six in
the outer p sublevels)
 These outer eight e- (valence
electrons) are called an octet.
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Octet Rule
Eight electrons in an outer level
render an atom unreactive.
 This is referred to as the Octet Rule.
 When atoms react with one another,
they do so to obtain a stable config.
 Some atoms gain or lose e- (ions)
and some share e- (molecules).
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Organizing Information on
the Periodic Table
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Use a pen to label the following:
Group 1
Alkali metals
Group 2
Alkaline earth metals
Group 16
Chalcogens
Group 17
Halogens
Group 18
Noble gases
Sc – Uub
Transition metals
La – Lu
Lanthanoids
Ac – Lr
Actinoids
Organizing Information on
the Periodic Table
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Draw a stair step dark line starting
between B and Al.
Label the right side: metals
Label the left side: nonmetals
Write METALLOID along stair step line.
Label the valence e- (outer electrons).
Use colored pencils to shade each group
or category a different color.
Basic Properties of Metals,
Nonmetals, and Metalloids
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Metals:
1. Dense and shiny (luster).
2. Conduct heat and electricity well.
3. Have high melting/boiling points
(high densities).
4.Malleable and ductile.
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Nonmetals:
1. Generally gases or brittle solids.
2. If solid, dull surface.
3. Good insulators.
4. Have low melting/boiling points
(low densities)
Metalloids:
1. Properties of both
metals and nonmetals.
2. Some semiconductors.
EX: Silicon, for example,
possesses a metallic
luster, yet it is an
inefficient conductor
(semiconductor) and is
brittle.
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Properties of Alkali Metals
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Group 1 metals
Soft silver metals.
Less dense than other
metals and lower melting
points.
Very reactive due to
large size and one
loosely held valence
electron.
Too reactive to be found
free in nature.
Properties of Alkaline
Earth Metals
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Group 2 Metals
Shiny silvery-white
metals
Have 2 valence electrons
Not as reactive as alkali
metals but very reactive
All found in the Earth’s
crust in mineral form
Too reactive to be found
in free element form
Properties of Halogens
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Group 17 nonmetals
All diatomic gases at
room temperature EX:
F2, (Br2 -liquid at room
temp)
Too reactive to be
found as free elements
in nature
Most important group
to be used in industry
Properties of Chalcogens
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Group 16 nonmetals
Diverse group that
includes nonmetals,
metalloids, and
metals
Properties of Noble Gases
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Group 18
nonmetals
Complete octet of
valence electrons
s2p6
Largely
unreactive
Monotomic gases
Periodic Trends
Using the Periodic Table to
Predict Properties of Elements
The basis of the periodic table is the
atomic structures of the elements.
 Position on the table and properties
of these elements arise from the econfigurations of the atoms.
 Properties such as density, atomic
radius, oxidation numbers,
ionization energy, and
electronegativity can be predicted.
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Trends in Oxidation
Numbers
Our knowledge of e- configurations
and the stability of noble gases
allows us to predict oxidation
numbers for elements.
 Oxidation numbers represent the
charge an ion obtains after losing or
gaining valence electrons.
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1+
2+
Tend to have more
than one oxidation
number
3+
3+ or 4+
3+
2+
or
4+
3-
2-
1-
0
Two hydrogen atoms are walking
down the road. One said, “I think I
lost an electron!”.
 “Really”, the other replied, “ Are you
sure?”.
 “Yes, I’m positive”.

Atomic Radius

Simply put, this is a measurement
of the size of an atom

(it’s determined by finding ½ the bond
distance between two atoms of the
same element).

#1. Group trends
As we increase the atomic
number (or go down a group). .
 each atom has another energy
level,
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so the atoms get
bigger
 #2
- Period Trends
Going from left to right across a period,
the size gets smaller.
 Electrons are in the same energy
level.
 But, there is more nuclear charge.
 Outermost electrons are pulled closer
which reduces the volume of the
electron cloud.
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Predicting Atomic Radius
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General rule: atomic size increases
as you move diagonally from top
right corner to bottom left corner.
When graphed, atomic radii
demonstrates a periodic trend
Radii of ions: Ions are atoms that
have gained or lost e- from the outer
orbitals.
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Cations: (+)
Become smaller
1. Positive charged
nucleus attracting
fewer e- so pulls
electron cloud in
tighter.
2. Reduced the
number of energy
levels.
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Na+1
11p+
10e-
Sodium
atom is
much
larger
than the
positive
sodium
ion.
Anions: (-)
 Become larger
1. Positive charged
nucleus attracting
more e- expands
electron cloud.
2. Add more energy
levels.
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S-2
16p+
18e-
The pull
on each
electron is
reduced
expanding
the
electron
cloud.
Ionization Energy
The energy required to remove an efrom an atom.
 The larger the atom, the less energy is
required because the e- are farther from
the positive center.
 As atoms get larger ionization energy
decreases because of the shielding effect
(which says that the farther an electron
is from the nucleus, the less tightly the
positive nucleus grabs it).

Remove the most loosely held e- is
first ionization energy.
 Measured in kilojoules per mole
kJ/mol
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Ionization energy increases
diagonally from bottom left corner to
top right corner.
Classification based on First
Ionization Energy
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METAL
1. Low 1st
ionization energy.
2. Located on left
side of Periodic
Table.
3. Form positive
ions.
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NONMETAL
1. High 1st
ionization energy.
2. Located on the
right side of
Periodic Table.
3. Form negative
ions.
Multiple Ionization Energies
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Additional e- can be lost from an atom
and the ionization energies can be
measured.
IONIZATION ENERGIES (kilojoules per mole)
Element
1st
2nd
3rd
4th
H
1312.0
He
2372.3
Li
520.2
7300
11750
Be
899.5
1760
14850
20900
B
800.6
2420
3660
25020
5th
5220
32660
Electronegativity
Electronegativity is the ability of an
atom to capture an electron.
 The smaller the atom the stronger
its ability to take electrons from
other atoms.
 Electronegativity is a unitless value.
Fluorine is highest at 3.98
Francium is the lowest at 0.7
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It increases from bottom left to top
right corners.
Electron Affinity
e- affinity is a measure of an atom’s
attraction for an e-.
 Metals have low e- affinities.
 Nonmetals have high e- affinities.
 Chemical reactions occur between
atoms with high e- affinity and those
with low e- affinity.
 EX: Al + Br  Al2Br3
(low) (high) (more stable)
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Review
Review

Based on our trends:
The most reactive metal element
would be
Francium
The most reactive nonmetal element
would be
Fluorine
In Summary
Periodic table is a chart of elements
in which the elements are arranged
based on their e- configurations
which dictates their properties.
 Moving down a group in the periodic
table, atomic radii becomes larger
because more energy levels are
needed for more e-.

In Summary
As the size becomes larger, the eare located farther away from the
positive center.
 This decreases the affinity of that
atom to hold on to these outer e-,
thus decreasing e- affinity.
 Ionization energy is low because it is
easy for the atom to lose these
outer e-.

In Summary
Moving across a period in the periodic
table, atomic radii becomes smaller
because the energy levels of periods
are the same but the positive centers of
atoms increase. This pulls the e- cloud
closer to the nucleus, making the atom
smaller.
 Ionization energy and e- affinity
increases for these smaller atoms.

THE END