Transcript The Periodic Table and Periodicity

The Periodic Table and
Periodicity
Areas of Interest
• Mendeleev’s brilliant organizational skills
• The modern table – groups, families and series
and trends
Dimitri Mendeleev
The father of the periodic table.
In the 19th century elements were being
discovered rapidly, a way was need to
organize them.
He arranged the atoms according to
increasing atomic weight.
Ok so what?
The brilliance of his arrangement came
from the atoms he left off the table . . .
Those that had not yet been discovered.
Mendeleev arranged his table in rows and
columns that not only addressed increasing
atomic mass but was able to predict
undiscovered elements based on
properties.
For years chemists had known about elements sharing similar properties, in
1869 Dimitri Mendeleev provided an organized arrangement.
His most famous omission he named eka-silicon. He predicted an element
that had a greater mass than silicon, a smaller mass than tin but shared
similar properties with both elements.
Property
Ekasilicon
Germanium
atomic mass
72
72.59
density (g/cm³)
5.5
5.35
melting point (°C)
high
947
color
gray
gray
oxide type
refractory dioxide
refractory dioxide
oxide density (g/cm³)
4.7
4.7
oxide activity
feebly basic
feebly basic
chloride boiling point
under 100°C
86°C (GeCl4)
chloride density (g/cm³)
1.9
1.9
Henry Moseley determined the atomic number
of each element and placed them in order of
increasing atomic number, rather than atomic
weight.
This atomic number arrangement also followed
Mendeleev’s trends in properties. Elements
with similar properties were grouped together.
The modern period table was born (1914).
Arranged in rows and columns.
 A row is called a period
 A column is called a group or family
Some of the groups (or families) have special names
that help us identify them as a collective.
Famous families if you will . . .
The include all of Group 1: Li, Na,
K, Rb, Cs and Fr*.
Soft shiny metals that react
violently with water to produce H2
gas.
Electron configurations of ns1. (ie
Li is 1s22s1)
Readily form +1 cations (ie Na
loses and electron to form Na+)
*Francium only exists for microseconds so it cannot be studied in quantity.
The include all of Group 2: Be,
Mg, Ca, Sr and Ba.
These metals are soft but not
quite as much as those of
Group 1.
They are stable in air (unlike
the Alkalai Metals)
Electron configurations of ns2.
(ie Be is 1s22s2)
Readily form +2 cations (ie Mg
loses 2 electrons to form Mg2+)
Electron-rich elements that most resemble what we think of when we
talk about metals:
 they’re malleable and ductile
 they conduct electricity
 the free flow of electrons yields many colorful solutions
 they’re shiny
 they conduct heat
Many of these elements are synthetic, they’re made in particle accelerators
and used for research or highly specific purposes.
They are metals but they are very dense and many are quite rare.
The metalloids
Elements include: B, Si, Ge, As, Sb, Te and
At
They’re not quite metals but they’re not quite
non-metals.
They’re semi-conductors (they can selectively
conduct electricity).
Si, the
semiconductor
the computer
industry is built
upon
The Halogens
Elements include: F, Cl, Br, I and At
Readily form -1 anions (ie Cl gains an
electron to form Cl-)
React well with metals from Groups 1 and
2.
Behave as other non-metals (nonconductive, not shiny etc.)
These are the elements found in Group 18, the
farthest to the right on the periodic table.
They are all gases and are VERY stable (they
do not readily undergo reaction).
The have full energy levels and sub-shells. For
example Ar has electron configuration
1s22s22p63s23p6.
When we pass a high-voltage current through
any of these gases we get extremely bright
light.
Metals
Lanthanides & Actinides
Transition Metals
Metalloids and Non
Metals
Use your textbook to define the following terms
1. Atomic radius – half the distance between 2 nuclei in a covalent bond
(an estimate of the radius of a single atom)
2. 1st Ionization Energy – the energy required to remove the 1st electron
from an atom in the gaseous state (generating a +1 cation)
3. 2nd Ionization Energy – the energy required to remove the 2nd
electron from a gaseous atom (generating a +2 cation)
4. Electron Affinity – the energy released when a gaseous atom
captures an electron (generating a -1 anion)
5. Electronegativity – the tendency of an atom to attract electrons to
itself (higher electronegativity means the atom “wants” electrons)
6. Activity – a trend that dictates replace-ability in single replacement
reactions (decreases down the halogens)
7. Valence Electron – an electron in the highest occupied energy level of
an atom (these are the ones involved in bonding)
8. Effective Nuclear Charge – the “pull” on the electrons from the
protons in the nucleus (decreases as atoms get larger)
Atomic Radius (Atomic Size)
Why this pattern?
Decreasing effective
nuclear charge (the
“pull” the electrons
feel from the protons
in the nucleus)
Atomic Radius – Half the distance of a pure covalent bond
Atomic Radius (Atomic Size)
1st Ionization Energy
- the energy required to remove 1 electron from a neutral atom
Ionization energy
decreases down a
Group.
Ionization increase from
left to right in a period.
Increasing
Electronegativity
Increasing
Effective Nuclear
Charge
Your Assignment
1. This element is larger than chlorine but smaller than iodine.
2. Of the following pairs, which is MORE metallic:
a. Si or Ge
b. As of Ge
c. Ba or Cs
d. Be or B
e. Kr or Xe
3. The electron configurations of six neutral atoms are shown:
i. 1s2 2s2 2p6 3s1
ii. 1s2
iii. 1s2 2s2 2p6 3s2
iv. 1s2 2s2 2p6
v. 1s2 2s1
vi. 1s2 2s2 2p3
a. Which has the highest first ionization energy?
b. Which has the lowest first ionization energy
c. Which would have the lowest second ionization energy?
d. Which would likely have the lowest third ionization energy?