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2. DEVELOPMENT AND FEATURES OF
THE MODERN PERIODIC TABLE OF
ELEMENTS
REPORTER: MS. MYRAFE M. RODELLAS
3. Essential Questions & Vocabulary
How was the periodic table developed?
What are the key features of the periodic table?
Vocabulary
Period law
Group
Period
Representative element
Transition element
Alkali metal
Alkaline earth metal
Transitional metal
Inner transition metal
Lanthanide series
Actinide series
Nonmetal
Halogen
Noble gas
metalloid
4. Main Idea
The periodic table evolved
over time as scientists
discovered more useful
ways to compare and
organize the elements.
9. Johann Wolfgang Dobereiner
(1817)
A German chemist who
formed the triads of
elements with similar
properties like the triad of
calcium, barium, and
strontium.
10. John Newlands (1864)
proposed an arrangement where
elements were ordered by increasing
atomic mass.
He noticed that when the elements
were arranged by increasing atomic
mass, their properties repeated every
eighth element (law of octaves).
11. John Lothar Meyer and Dmitri
Mendeleev (1869)
They both demonstrated a
connection between
atomic mass and
elemental properties.
arranged elements in
order of increasing atomic
mass into columns with
similar properties.
13. Henry Moseley - 1914
An English physicist who observed that
the order of the x-ray frequencies
emitted by elements follows the ordering
of elements by atomic number.
His observation led to the development
of the modern periodic law which states
that the properties of elements vary
periodically with atomic number.
26. AlkaliMetals (Li, Na,K,Rb, Cs, Fr)
Atoms of the alkali
metals have a single
electron in their
outermost level, in other
words, 1 valence
electron.
27. Alkali Metals the most reactive
metals.
react with water to
release hydrogen
gas.
never found as
free elements in
nature.
28. Alkaline Earth Metals
(Be, Mg, Ca, Sr, Ba, Ra)
They are found in Group
IIA or 2
They have two valence
electrons.
They are very reactive,
but not as much as the
alkali metals.
29. Boron Family
named after the first
element in the family.
found in Group IIIA or
13.
Atoms in this family
have 3 valence
electrons.
30. Carbon Family
Atoms of this family
have 4 valence
electrons.
They found in Group
IVA or 14.
31. Nitrogen Family
The nitrogen family
is named after the
element that makes
up 78% of our
atmosphere.
They are found in
Group VA or 15
Atoms in the
nitrogen family have
5 valence electrons.
They tend to share
electrons when they
bond.
32. Oxygen Family
Atoms of this family
have 6 valence
electrons.
They belong to
Group VIA or 16.
Most elements in
this family share
electrons when
forming compounds.
33. Halogen Family (F, Cl, Br, I, At
They can be
found in Group
VIIA or 17.
Halogens have 7
valence
electrons. They react with alkali metals
to form salts.
35. Transition Metals
The compounds of
transition metals are
usually brightly
colored and are often
used to color paints.
Transition elements
have 1 or 2 valence
electrons, which they
lose when they form
bonds with other
atoms.
36. Noble Gases
colorless gases that are extremely un-
reactive.
One important property is their
inactivity. They are inactive because
their outermost energy level is full.
Because they do not readily combine
with other elements to form compounds,
the noble gases are called inert.
37. Rare Earth Elements
The thirty rare earth
elements are
composed of the
lanthanide and
actinide series.
One element of the
lanthanide series
and most of the
elements in the
actinide series are
called trans-
uranium, which
means synthetic or
man-made.
40. Essential Questions & Vocabulary
What are the period and group trends of different
properties?
How are period and groups trends in atomic radii related
to electron configuration?
Vocabulary
Energy level of an atom
Electron Shielding
Ions (cation and anion)
Effective Nuclear Charge
Atomic & Ionic Radii
Ionization energy
Electronegativity
reactivity
41. Trends among
elements in the
periodic table include
their sizes and their
abilities to lose or
attract electrons.
Big Idea
44. 1. Atomic Radius
• Atomic radius: defined as ½ distance
between neighboring nuclei in molecule or
crystal
• It is affected by:
1. # of energy
levels
2. Proton Pulling
Power
47. 2-8
Ne
VIIIA or 18
2-7
F
VIIA or 17
2-6
O
VIA or 16
2-5
N
VA or 15
2-4
C
IVA or 14
2-3
B
IIIA or 13
2-2
Be
IIA or 2
2-1
Li
IA or 1
Configuration
Element
Family
49. • As you go from left to right, you gain
more protons (the atomic number
increases)
• You have greater “proton pulling
power”
50. as go across row size tends to decrease a bit
because of greater PPP “proton pulling power”
previous | index | next
,
51. Wecan“measure” theProtonPullingPower
by determining theEffective nuclearcharge
• It is the charge actually felt by valence electrons
• The equation
Nuclear charge - # inner shell electrons
(doesn’t include valence e-)
52. previous | index | next
Calculate “effective nuclear charge”
# protons minus # inner electrons
53. What the inner electrons do….
They shield the charge felt by the valence electrons.
54. H and He: only
elements whose
valence
electrons feel
full nuclear
charge (pull)
NOTHING TO
SHIELD
THEM
57. Atomic Radii – Practice I
Rank the following atoms in increasing
atomic radius.
Fluorine < Carbon < Beryllium < Lithium
• Carbon
• Fluorine
• Beryllium
• Lithium
58. Ions
An ion is an atom or bonded group of atoms with a positive or
negative charge.
Cations: atoms lose electrons and become positively charged
Anions: atoms gain electrons and become negatively charged
59. Cation Formation
11p+
Na atom
1 valence
electron
Valence e-
lost in ion
formation
Effective
nuclear charge
on remaining
electrons
increases.
Remaining e-
are pulled in
closer to the
nucleus. Ionic
size decreases.
Result: a smaller
sodium cation, Na+
60. Anion Formation
17
p+
Chlorine
atom with 7
valence e-
One e- is
added to the
outer shell.
Effective nuclear charge is reduced and
the e- cloud expands.
A chloride ion
is produced. It
is larger than
the original
atom.
61. Cations are smaller than the
neutral atom
Smaller than the neutral atom
1. The loss of a valence electron can leave an empty outer orbital,
resulting in a smaller radius.
2. Electrostatic repulsion decreases allowing the electrons to be
pulled closer to the nucleus
62. Anions – Bigger than the
neutral atom
Why?
The addition of an electron increases electrostatic
repulsion.
63. The ionic radii positive ions (cations) generally decrease
from left to right.
The ionic radii of negative ions (anions) generally decrease
from left to right, beginning with group 15 or 16.
Both positive and negative ions increase in size moving
down a group.
Ionic Radius
65. Ionic Radii – Practice
Arrange the following ions in order of
increasing ionic radius:
Na+, Al3+, Mg2+
Al3 + Mg2+ Na+
66. Atomic & Ionic Radii-
Mixed Practice
A. If the figure represents the atoms helium, krypton, and
radon, match the letter to the correct atom.
A. If the figure represents a cation, an anion, and a neutral
atom from the same period, match the letter to correct term.
A B C
A – Radon B – Krypton C - Helium
A – Anion B – Atom C - Cation
67. 2. IonizationEnergy
= amount of energy required to remove a
valence electron from an atom in gas phase
1st ionization energy = energy required to
remove the most loosely held valence
electron (e- farthest from nucleus)
68. Cs valence electron is a lot farther away from nucleus than
Li
•electrostatic attraction is much weaker so it becomes
easier to steal electron away from Cs
•THEREFORE, Li has a higher Ionization energy than
Cs
70. Ionization Energy - Practice
Arrange the following elements in order of
decreasing Ionization Energy.
Al, Mg, Na, Si
Si Al Mg Na
71. 3. Electronegativity
• ability of atom to attract electrons in bond
• noble gases tend not to form bonds, so they
don’t have electronegativity values
• Unit = Pauling
• Fluorine: most electronegative element
= 4.0 Paulings
79. 5. Metallic Character
Properties of a Metal –
Easy to shape (malleable); many are ductile
(can be pulled into wires)
Conduct electricity and heat
Shiny
Group Trend – As you go down a column, metallic character
increases (because ionization energy decreases).
Periodic Trend – As you go across a period (L to R), metallic
character decreases (because ionization energy decreases)
(L to R, you are going from metals to non-metals).
83. Atomic Radius
Atomic size is a periodic trend
influenced by electron
configuration.
For metals, atomic radius is
half the distance between
adjacent nuclei in a crystal of
the element.
84. Atomic Radius
For elements that occur as
molecules, the atomic radius
is half the distance between
nuclei of identical atoms
that are chemically bonded
together.
86. Shielding/Screening
Electrons have an attraction or pull towards the nucleus of the
atom (opposite charges attract)
Electrons are also repelled away from the inner electrons (like
charges repel)
Shielding/ Screening: the attraction of valence (outer-shell)
electrons is counterbalanced by the repulsion of the inner-shell
electrons.
The inner-shell electrons “screen” or “shield” the outer-shell
electrons from full attraction
87. Effective Nuclear Charge
Effective nuclear charge is the net
positive charge experienced by valence electrons.
88. Interesting IE Pattern
The ionization at which the large increase in energy
occurs is related to the number of valence electrons.
89. Octet Rule
Octet rule - states that atoms tend to gain, lose or share
electrons in order to acquire a full set of eight valence electrons.
The octet rule is useful for predicting what types of ions an
element is likely to form.
90. Electronegativity
Ability for an atom to attract electrons
When it is chemically combined with another
atom.
Elements with high electronegativities
(nonmetals) often gain electrons to form
anions.
Elements with low electronegativities (metals)
often lose electrons to form cations.
91. Electronegativity
Increases from left to right
Decreases from top to bottom
Fluorine has the
highest
electronegativity
Opposite trend of Atomic Radius. Smaller radii – higher
electronegativity (closer electron can get to the nucleus)
92. Electronegativity Visual
Which visual representation best describes
electronegativity?
The ability of a nucleus of one atom to attract an
electron from another atom in a chemical bond.
Look at all the shielding Francium's one valence electron has. It barely feels the proton pull from the nucleus. No wonder it will lose it’s one electron the easiest. No wonder it’s the most reactive metal
Atoms become charged by either gaining or losing electrons.
Increasing from left to right because as the atom increases its no. of proton, more energy is required to remove its valence electron.