The Periodic Table
Guiding Questions
Why is the periodic table so important?
Why is the periodic table shaped the way it's shaped?
Why do elements combine? Why do elements react?
What other patterns are there in the world and how do
they help us?
Development of Periodic
Table
J.W. Döbereiner (1829)
Law of Triads
Elements could be classified into groups of three, or triads.
Trends in physical properties such as density, melting point,
and atomic mass were observed.
J.A.R. Newlands (1864)
Law of Octaves
Arranged the 62 known elements into groups of seven
according to increasing atomic mass.
He proposed that an eighth element would then repeat the
properties of the first element in the previous group.
Mendeleev
• “Father of Periodic Table”
• organized elements based on increasing
atomic mass. Found similarities in chemical
properties and published his first periodic
table in 1869.
Moseley
• discovered while working with elements that
they fit better into pattern when arranged by
nuclear charge (number of protons-also
known as atomic number).
• Periodic Law states that physical and
chemical properties of the elements are
periodic functions of their atomic numbers.
• This means that when elements are
arranged in order of increasing atomic
number, elements with similar properties
appear at regular intervals
Organization of the Periodic
Table
• Periods – rows across, horizontal
– Indicates number of energy levels
(Principle Quantum Number)
• Groups/Families – down, columns
– 18
– Indicate number of e- in outer most energy
level (1-2, 13-18 main group elements)
– Share chemical properities
Three Categories of Elements
• Metals
• Non-Metals
• Metalloids
Three Categories of Elements
• Metals
- Groups 1-12 (except H) and under stair-step groups 13-15
- Form ionic and metallic bonds
- Luster (shiny, reflects light)
- Malleable - can be flattened into sheets
- Ductile - can be drawn into thin wires
- Good Conductors - heat and electricity can flow
throughbecause the outer electrons are not held tightly to
the nucleus and move freely
- Most are solid at room temperature except Mercury
The Periodic Table
H
He
Li
Be
B
C
N
O
F
Ne
Na
Mg
Al
Si
P
S
Cl
Ar
K
Ca
Sc
Ti
V
Cr
Mn
Fe
Co
Ni
Cu
Zn
Ga
Ge
As
Se
Br
Kr
Rb
Sr
Y
Zr
Nb
Mo
Tc
Ru
Rh
Pd
Ag
Cd
In
Sn
Sb
Te
I
Xe
Cs
Ba
La
Hf
Ta
W
Re
Os
Ir
Pt
Au
Hg
Ti
Pb
Bi
Po
At
Rn
Fr
Ra
Ac
Rf
Db
Sg
Bh
Hs
Mt
Ds
Rg
The Metals are represented in the Periodic Table in blue.
Three Categories of Elements
• Non-Metals
- Dull
- Not Malleable/Ductile
- Poor Conductors
- P block
- Form ionic and covalent bonds
- Most do not conduct heat or electricity
-All, except H, are found on right of
periodic table
- At room temperature, most are solid or gaseous
The Periodic Table
H
He
Li
Be
B
C
N
O
F
Ne
Na
Mg
Al
Si
P
S
Cl
Ar
K
Ca
Sc
Ti
V
Cr
Mn
Fe
Co
Ni
Cu
Zn
Ga
Ge
As
Se
Br
Kr
Rb
Sr
Y
Zr
Nb
Mo
Tc
Ru
Rh
Pd
Ag
Cd
In
Sn
Sb
Te
I
Xe
Cs
Ba
La
Hf
Ta
W
Re
Os
Ir
Pt
Au
Hg
Ti
Pb
Bi
Po
At
Rn
Fr
Ra
Ac
Rf
Db
Sg
Bh
Hs
Mt
Ds
Rg
The Non-Metals are represented in the Periodic Table in yellow.
Three Categories of Elements
• Non-Metals
Some Examples of Non-Metals
- Oxygen (O)
- Helium (He)
- Sulfur (S)
- Chlorine (Cl)
- Neon (Ne)
- Nitrogen (N)
Three Categories of Elements
• Metalloids
- Have properties of both metals and non-metals.
- Generally not shiny
- Not malleable and ductile
- Form ionic and covalent bonds
- They conduct heat and electricity better than nonmetals,
but less than metals.
Three Categories of Elements
• Metalloids
Here are the Metalloids
- Boron (B)
- Arsenic (As)
- Tellurium (Te)
- Antimony (Sb)
- Polonium (Po)
- Silicon (Si)
- Germanium (Ge)
The Periodic Table
H
He
Li
Be
B
C
N
O
F
Ne
Na
Mg
Al
Si
P
S
Cl
Ar
K
Ca
Sc
Ti
V
Cr
Mn
Fe
Co
Ni
Cu
Zn
Ga
Ge
As
Se
Br
Kr
Rb
Sr
Y
Zr
Nb
Mo
Tc
Ru
Rh
Pd
Ag
Cd
In
Sn
Sb
Te
I
Xe
Cs
Ba
La
Hf
Ta
W
Re
Os
Ir
Pt
Au
Hg
Ti
Pb
Bi
Po
At
Rn
Fr
Ra
Ac
Rf
Db
Sg
Bh
Hs
Mt
Ds
Rg
The Metalloids are represented in the Periodic Table in green.
The Periodic Table
H
He
Li
Be
B
C
N
O
F
Ne
Na
Mg
Al
Si
P
S
Cl
Ar
K
Ca
Sc
Ti
V
Cr
Mn
Fe
Co
Ni
Cu
Zn
Ga
Ge
As
Se
Br
Kr
Rb
Sr
Y
Zr
Nb
Mo
Tc
Ru
Rh
Pd
Ag
Cd
In
Sn
Sb
Te
I
Xe
Cs
Ba
La
Hf
Ta
W
Re
Os
Ir
Pt
Au
Hg
Ti
Pb
Bi
Po
At
Rn
Fr
Ra
Ac
Rf
Db
Sg
Bh
Hs
Mt
Ds
Rg
What do you notice about the way the element groups are
arranged in the Periodic Table?
Hydrogen
• In a class by itself
• Most common element in the universe (3/4)
• Behaves unlike any other element because it
consists of one proton and one electron.
Main Group Elements
• S and p blocks
• Groups 1,2, and 13-18
• Four groups have special names:
Alkali metals (Group 1)
Alkaline-earth metals (Group 2)
Halogens (Group 17)
Noble gases (Group 18)
Group 1 – Alkali Metals
Characteristics
• Soft, silver metals
• Low melting points and densities
• Highly reactive (esp. w/ water)
• Do not occur in nature in elemental form
• Stored in kerosene or mineral oil
• Have one Valence electron
Group 2 – Alkaline-earth
Metals
Characteristics
• Gray, metallic solids
• Reactive; but, less than alkali
• Not found as free elements in nature
• Bright fireworks, aircraft
• Harder, denser, and stronger than alkali
metals
• Higher melting points than alkali
Have 2 valence electrons
Transition Elements
• Groups 3-12 (elements in transition) D
block
• Vary in reactivity and can be found as free
elements
• Metal characteristics
• Form colored compounds
• Often occur in nature as uncombined
elements
• Hg – mercury – liquid metal
Valence Electrons for
Transition Elements
–
–
–
–
–
–
–
–
–
–
Group 3: 3 valence electrons
Group 4: 2 to 4 valence electrons
Group 5: 2 to 5 valence electrons
Group 6: 2 to 6 valence electrons
Group 7: 2 to 7 valence electrons
Group 8: 2 or 3 valence electrons
Group 9: 2 or 3 valence electrons
Group 10: 2 or 3 valence electrons
Group 11: 1 or 2 valence electrons
Group 12: 2 valence electrons
Group 13 – Boron Group or
Icosagens
• Have 3 valence electrons
• Boron is the only metalloid in the family
• The rest are poor metals.
Group 14 – Carbon group or
Crystallogens
• Have 4 valence electrons
• Unique feature is that the elements can form
different anions and cations
Ex: C: 4Si and Ge: 4+
Sn and Pb: 2+
Group 15 – Nitrogen Group
or Pnictogens
• Have 5 valence electrons
• Able to form double and triple bonds
Group 16 – Oxygen group or
Chalcogens
• Have 6 valence electrons
• From -2 ions
• Physical properities vary dramatically
Ex: oxygen – colorless gas
Sulfur – yellow solid
Tellurium – silver metalloid
Selenium - black
Group 17 - Halogens
•
•
•
•
•
•
•
•
Most reactive non-metal
Irritating odor
Interact with alkali metals to form salts
7 electrons in outer energy level
Easily pickup one electron
Bromine – only liquid nonmetal
Flourine and chlorine – gases at room temp
Iodine and astatine – solids at room temp
Group 18 – Noble Gases
•
•
•
•
No color or odor
Exist as individual gas atoms (monatomic)
Full outer energy level
Relatively un-reactive (inert gases)
Inner Transitional Metals
• f-block, n-2
– ALL are radioactive and unstable
– Includes lanthanides (atomic number 5871)
• Shiny metals, similar in reactivity to
alkaline earth metals
– and actinides (atomic number 90-103)
• First 4 are found on earth, the remaining
are synthetic
Diatomic Molecules
Elements That Exist as Diatomic Molecules in Their Elemental Forms
Element Present
hydrogen
nitrogen
oxygen
fluorine
chlorine
bromine
iodine
Elemental State at 25 oC
colorless gas
colorless gas
pale blue gas
pale yellow gas
pale green gas
reddish-brown liquid
lustrous, dark purple solid
Molecule
H2
N2
O2
F2
Cl2
Br2
I2
How do we use an elements location
on the periodic table to determine its
ionic charge?
•
•
•
•
Atomic Radius (Angstrom)
½ the distance from the nuclei to another
From the nucleus to edge of e- cloud
Going down a group, Atomic radius
increases because of the increasing number
of energy levels
• Going across a period (left to right), atomic
radius decreases because of the increase in
positive charge in the nucleus
Atomic Radii
Coulombic attraction
• Attraction of + and – charges
• Two factors determine the strength:
1. Amount of charge
2. Distance between charges
Shielding Effect
• Kernel electrons “shield” valence electrons
from attractive force of the nucleus
• Caused by kernel and valence electrons
repelling each other
• The more electron shells there are, the
greater the shielding effect.
• Explains why valence electrons are more
easily removed
Shielding Effect
Valence
+
nucleus
-
Kernel electrons block
the attractive force of
the nucleus from the
valence electrons
-
Electron
Shield
“kernel”
electrons
Electrons
Why is cesium bigger than
sodium?
• Sodium has the electron configuration 1s22s22p63s1.
There is one valence electron.
• The attraction between this lone valence electron
and the nucleus with 11 protons is shielded by the
other 10 core electrons.
• The electron configuration for cesium is
1s22s22p63s23p64s23d104p65s24d105p66s1.
• While there are more protons in a cesium atom,
there are also more electrons shielding the outer
electron from the nucleus. The outermost electron,
6s1, therefore, is held very loosely. Because of
shielding, the nucleus has less control over this 6s1
electron than it does over a 3s1 electron.
Ionization Energy
• Ionization is a process that results
in the formation of an ion
– An ion is an atom or group of
atoms that have a “+” or “-”
charge
• Change is created by gain or loss of e• Losing an e creates a “+” charge
(cation)
Ionization Energy
• Losing the e- requires energy. Energy
required to remove one e- from a neutral
atom is called
IONIZATION ENERGY
(also known as 1st ionization
energy)
Factors that affect ionization
• Nuclear charge – the larger the nuclear
charge , the greater IE
• Shielding effect – the greater the shielding
effect, the less IE
• Radius – greater the distance between the
nucleus and the outer electrons of an atom,
the less IE.
• Sublevel – an electron from a sublevel that is
more than half-full requires additional energy
to be removed
• Ionization energy decreases as you go down
a group because valence electrons are
farther from the nucleus
• And increases as you go across a period
because of the greater positive charge leads
to greater attraction to electron
Ionization Energy
Electron Affinity
– Gaining an electron results in an
ion with a “-” charge (anion)
• When an atom gains an e- it causes an
energy change. The energy change
when a neutral atom gains an e- is the
ELECTRON AFFINITY
Electron Affinity
• Electron affinity decreases as you move
down a group
• Increases as you move across a period
• Halogens have the highest electron affinities
Electron Affinity
Electronegativity (Pauling)
Ability of an atom to attract (or
remove) an e- from another
atom
• Fluorine is most electronegative
(4)
• Metals have Electronegativity of
less than 2
Electronegativity
• Electronegativity decreases as you go down
a group because the electrons are farther
from the nucleus
• Electronegativity increases as you go to the
right because atoms are more inclined to
gain electrons in order to gain a full shell
Electronegativity
Nuclear charge increases
Shielding increases
Atomic radius increases
Ionic size increases
Ionization energy decreases
Electronegativity decreases
Summary of Periodic Trends
Shielding is constant
Atomic radius decreases
Ionization energy increases
Electronegativity increases
Nuclear charge increases
1A
0
2A
Ionic size (cations)
decreases
3A 4A 5A 6A 7A
Ionic size (anions)
decreases