JW Dobereiner
Early 1800s
Classified elements into triads based on similar properties
Ca, Sr, Ba
Li, Na, K
Cl, Br, I
J.A.R. Newlands
Law of Octaves – every 8 elements begin to repeat in
patterns
Wasn’t taken seriously because he likened chemistry to music
Dmitri Mendeleev
Father of modern periodic table
Arranged the elements according to increasing atomic
mass
Left blanks for elements that hadn’t been discovered
but that he predicted existed
Made some exceptions (Te and I) because he knew
that iodine was more closely related to Br based on its
properties
Henry Moseley
Assigned atomic numbers to elements (number of
protons)
Arranged the elements in order of increasing atomic
number
No need to make exceptions like Mendeleev had to
Periodic Law
There is a periodic repetition of chemical and physical
properties of elements when they are arranged by
increasing atomic number.
Groups – Vertical families
Group 1 or 1A-Alkali Metals
Group 2 or 2A-Alkaline Earth Metals
Group 13 or 3A-Boron Family
Group 14 or 4A-Carbon Family
Group 15 or 5A-Nitrogen Family
Group 16 or 6A-Oxygen Family
Group 17 or 7A-Halogen Family
Group 18 or 8A-Noble Gas Family
Groups
Note: All of the Group A elements are called the “Representative Elements.”
Group A # = # of valence electrons
The other elements (Group B elements) are known as the transition and inner transition elements.
The Octet Rule- atoms tend to gain, lose, or share electrons to acquire a full set of eight valence electrons.
Some classifications
Metals – grey, typically lustrous or shiny, good
conductors of heat and electricity, and are generally
solids at R.T
Nonmetals – pink, have opposite properties. They can
be solid, liquid or gas at R.T
Metalloids/Semi-metals – blue, along the zigzag
Periodic Trends
Electron Shielding – when the presence of core
electrons diminishes the effect the nuclear charge has
on the valence electrons
Atomic Radii – 1/2 the distance between the nuclei of
two atoms of the same element when the atoms are
joined
Periodic Trends
Atomic Radii
Down a group – atomic radii tends to increase
Why?
Electrons enter higher principle energy levels
Across a period – tends to decrease
Why?
# of protons increase, electrons enter same
principal energy level, no increase in shielding
but increase in nuclear charge
Periodic Trends
Ionization energy – the energy required to remove an electron from a gaseous atom
Down a group – tends to decrease
Why?
Electrons are farther from nucleus therefore easier to pull off
Across a period – tends to increase
Why?
More difficult to remove an electron as you get closer to noble gas configuration
First Io
niz
ation
ene
rgy
Atomic number
He
He has a greater IE than H.
same shielding
greater nuclear charge
H
First Io
niz
ation
ene
rgy
Atomic number
H
He
Li has lower IE than
H
more shielding
further away
outweighs greater
nuclear charge
Li
First Io
niz
ation
ene
rgy
Atomic number
H
He
Be has higher IE
than Li
same shielding
greater nuclear
charge
Li
Be
First Io
niz
ation
ene
rgy
Atomic number
H
He B has lower IE than
Be
same shielding
greater nuclear
charge
By removing an
electron we make s
orbital full Li
Be
B
First Io
niz
ation
ene
rgy
Atomic number
H
He
Li
Be
B
C
N
O
Breaks the pattern because
removing an electron gets to
1/2 filled p orbital
First Io
niz
ation
ene
rgy
Atomic number
H
He
Li
Be
B
C
N
O
F
Ne Ne has a lower IE than He
Both are full,
Ne has more shielding
Greater distance
First Io
niz
ation
ene
rgy
Atomic number
H
He
Li
Be
B
C
N
O
F
Ne
Na has a lower
IE than Li
Both are s1
Na has more
shielding
Greater distance
Na
Periodic Trends
Electronegativity – the relative ability of an atom to attract an electron when the atom is in a compound
Down a group – tends to decrease
Why?
The shielding effect is greater
Across a period – tends to increase (excluding noble gases)
Why?
Greater ‘desire’ to achieve noble gas configuration
Periodic Trends
Ionic Size
metals - Cations
smaller than their respective atoms
Why?
# of electrons is less # of protons
Example: Na+ is smaller than Na
Periodic Trends
Ionic Size
Non-metals - Anions
larger than their respective atoms
Why?
# of electrons is larger than # of protons
Example: Cl- is bigger than Cl
Comparing ions to ions regarding size
Down a group – ions tend to increase in size
Across a period – tend to decrease
Size of Isoelectronic ions Iso - same
Iso electronic ions have the same # of electrons
Al+3 Mg+2 Na+1 Ne F-1 O-2 and N-3
all have 10 electrons
all have the configuration 1s22s22p6