development-of-the-PTE.ppt

DEVELOPMENT AND FEATURES OF
THE MODERN PERIODIC TABLE OF
ELEMENTS
REPORTER: MS. MYRAFE M. RODELLAS

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

Main Idea
The periodic table evolved
over time as scientists
discovered more useful
ways to compare and
organize the elements.

Lavoisier (1700’s)
 compiled a list of all the known elements of the time.

Johann Wolfgang Dobereiner
(1817)
 A German chemist who
formed the triads of
elements with similar
properties like the triad of
calcium, barium, and
strontium.

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).

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.

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.

FEATURES OF THE
PERIODIC TABLE OF
Elements

The periodic table is divided into:
 PERIODS or SERIES - horizontal rows
 GROUPS or FAMILIES - vertical columns

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.

Alkali Metals  the most reactive
metals.
 react with water to
release hydrogen
gas.
 never found as
free elements in
nature.

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.

Boron Family
 named after the first
element in the family.
 found in Group IIIA or
13.
 Atoms in this family
have 3 valence
electrons.

Carbon Family
 Atoms of this family
have 4 valence
electrons.
 They found in Group
IVA or 14.

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.

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.

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.

Transition Metals
 Transition
Elements include
those elements in
the B families.
 They are good
conductors of
heat and
electricity.

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.

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.

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.

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

Trends among
elements in the
periodic table include
their sizes and their
abilities to lose or
attract electrons.
Big Idea

2-8-18-32-18-8-1
Fr
7
2-8-18-18-8-1
Cs
6
2-8-18-8-1
Rb
5
2-8-8-1
K
4
2-8-1
Na
3
2-1
Li
2
1
H
1
Going down tocolumn 1:
increasing # energy levels as we go down

Increasing number of energy levels

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

Increasing
number
of
energy
levels
Increasing AtomicRadius

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

Increasing
number
of
energy
levels
Increasing
Atomic
Radius
Decreasing
AtomicRadius
Why does
this happen?

• As you go from left to right, you gain
more protons (the atomic number
increases)
• You have greater “proton pulling
power”

as go across row size tends to decrease a bit
because of greater PPP “proton pulling power”
previous | index | next
,

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-)

previous | index | next
Calculate “effective nuclear charge”
# protons minus # inner electrons

What the inner electrons do….
They shield the charge felt by the valence electrons.

H and He: only
elements whose
valence
electrons feel
full nuclear
charge (pull)
NOTHING TO
SHIELD
THEM

Increasing
number
of
energy
levels
Increasing
Atomic
Radius
Decreasing Atomic Radius
Increased Electron Shielding

Atomic Radii – Practice I
 Rank the following atoms in increasing
atomic radius.
Fluorine < Carbon < Beryllium < Lithium
• Carbon
• Fluorine
• Beryllium
• Lithium

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

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+

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.

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

Anions – Bigger than the
neutral atom
 Why?
 The addition of an electron increases electrostatic
repulsion.

 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

Ionic Radii – Practice
 Arrange the following ions in order of
increasing ionic radius:
 Na+, Al3+, Mg2+
Al3 + Mg2+ Na+

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

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)

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

Increasing
number
of
energy
levels
Increasing
Atomic
Radius
Increased
Electron
Shielding
Decreased Ionization Energy
(easier to remove an electron)
Increased Ionization Energy (harder to remove an electron)

Ionization Energy - Practice
 Arrange the following elements in order of
decreasing Ionization Energy.
 Al, Mg, Na, Si
Si Al Mg Na

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

Increasing
number
of
energy
levels
Increasing
Atomic
Radius
Increased
Electron
Shielding
Decreased
Ionization
Energy
(easier
to
remove
an
electron) Increased Ionization Energy (harder to remove an electron)
Decreased Electronegativity
IncreasedElectronegativity

Electronegativity - Practice
 Arrange the following in increasing order of
electronegativity:
 Na, Li, K
 Ca, Br, Se
K Na Li
Ca Se Br

4. Reactivity ofMetals
• judge reactivity of metals by how they
easily give up electrons (they’re losers)
• Low I.E = High Reactivity.

Reactivity of Non-metals
 judge reactivity of non-metals by how
they easily gain electrons (they are
winners)

Increasing
number
of
energy
levels
Increasing
Atomic
Radius
Increased
Electron
Shielding
Decreased
Ionization
Energy
(easier
to
remove
an
electron)
Decreased
Electronegativity
Increased Electronegativity
More metallic
Most Reactive
Nonmetal
= F
Nonreactiv
e
BACK
Most reactive metal = Fr
(the most metallic)

PRACTICE:
Arrange the following elements in
order of increasing reactivity of
metals:
Cs, K, Rb
K Cs Rb

PRACTICE:
order of decreasing reactivity of
non-metals:
Se, O, S
Se S O

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).

Increasing
number
of
energy
levels
Increasing
Atomic
Radius
Increased
Electron
Shielding
Decreased
Ionization
Energy
Decreased
Electronegativity
Increased Electronegativity
Most Reactive
Nonmetal
= F
Nonreactiv
e
BACK
Metallic characterdecreases
Metallic characterincreases

PRACTICE:
order of decreasing metallic
character:
Pt, Rh, Fe
Pt Rh Fe

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.

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.

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

Effective Nuclear Charge
 Effective nuclear charge is the net
positive charge experienced by valence electrons.

Interesting IE Pattern
 The ionization at which the large increase in energy
occurs is related to the number of valence electrons.

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.

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.

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)

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.

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