Transcript Periodic Law Power Point

The Periodic Law
Chapter 5
History of the Periodic Table
Before 1860, there was no method
for accurately determining an
element’s atomic mass.
Different chemists used different
atomic masses for the same
elements, resulting in different
compositions being proposed for
the same compounds.
In 1860 scientists led by Stanislao
Cannizzaro standardized calculations
for atomic mass.
Then in 1869, a Russian chemist,
Dimitri Mendeleev published the first
periodic table.
The table was a list of known elements
arranged by increasing atomic mass.
Mendeleev’s grouped elements
showed repeating patterns of
properties. Repeating patterns are
called periodic functions and the
elements arranged in this way are
said to show periodicity.
The second hand of a watch, for
example, passes over and given
mark at periodic, 60-second
intervals.
Mendeleev’s table had blanks which
showed places where unknown
elements were later placed.
Mendeleev also noted discrepancies
between grouped properties and some
of the element's atomic mass orders.
For example:
Mendeleev placed iodine (127) after
tellurium (128) so tellurium would be in
a group of elements with which it
shared similar properties.
But the question remained—why could
most of the elements be arranged in
the order of increasing atomic mass,
but a few could not?
In 1911, English scientist, Henry
Mosely noticed a better pattern. He
made a periodic table ordered by
increasing atomic number.
This modification helped correct the
discrepancies between properties
and order noted by Mendeleev.
Modern Periodic Table
Periodic Law
The physical and chemical
properties of the elements are
periodic functions of their atomic
number.
In other words, when the elements
are arranged in order of increasing
atomic number, elements with
similar properties appear at regular
intervals.
Periodic Table
An arrangement of the elements in
order of their atomic numbers so
that elements with similar
properties fall in the same column
or group.
Electron Configuration
and the Periodic Table
Before we continue…
Chemical compounds are formed
because electrons are lost, gained, or
shared between atoms.
The electrons that interact in this
manner are those in the highest
energy level.
These electrons are the most subject
to the influence of nearby atoms.
The electrons available to be lost, gained,
or shared in the formation of chemical
compounds are referred to as valence
electrons.
The stability of the noble gases results
from their special electron configurations.
The highest occupied energy levels
contain stable octets, or 8 valence
electrons.
Horizontal rows or periods
represent primary energy levels.
The vertical columns (groups or
families) are arranged to place
elements with the same outer level
electrons (valence electrons)
together.
Generally the electron configuration of
an atom’s highest occupied energy
level governs the atom’s chemical
properties.
So, if the periodic table is arranged to
properties, it makes sense that you will
see a pattern in the electron
configuration.
Blocks of the Periodic Table
The structure of the table results in
blocks of elements based on the
filling of the energy sublevels.
s-block Elements
Include Group 1 (alkali metals) and the
Group 2 (alkaline–earth metals).
These elements are chemically
reactive metals which do not occur as
free elements in nature.
s-block elements have 1 or 2, valence
electrons.
Hydrogen
Hydrogen is a big exception to the
block structure of the periodic table.
For convenience hydrogen is
usually placed on top of Group 1
although it is not considered an
alkali metal.
It has a 1s1electron configuration
but does not share the properties of
the elements in group 1.
Hydrogen's structure and
properties make it unique—it
doesn’t really fit with any group.
Helium
Helium has a similar electron
configuration as the Group 2
elements, but it is part of Group 18.
Because its highest occupied
energy level is filled by 2 electrons,
helium possesses special chemical
stability like the noble gases.
p-Block Elements
Elements filling the p sublevel.
Includes metals, metalloids and
nonmetals.
Includes the most reactive
nonmetals, Group 17 (halogens).
Includes the least reactive
elements, Group 18 (noble gases).
P-block and s-block elements together
make up the main-group or
representative elements.
All p-block elements have 3-8 valance
electrons with the exception of helium,
which has 2 valance electrons.
The valence electrons are in the outer
most s and p sublevels.
d-Block Elements
Elements filling d sublevels.
All elements in d-block are transition
elements.
All d-block elements are metals with
varying properties caused by interactions
of an unfilled d sublevel interacting with an
unfilled higher primary energy level.
This is where the exceptions to the rules
are found.
f-Block Elements
Elements filling the f sublevel.
f-block elements should appear in the
center of the periodic table, between
groups 3 and 4 in the 6th and 7th
periods.
For convenience scientists place these
two periods underneath the table to
shorten the table.
The top row of the f-block is called the
lanthanide series.
The bottom row of the f-block is called
the actinide series.
Each series is named after the element
that precedes them in the chart
(lanthanum- lanthanide) and (actiniumactinide).
Trends in the Periodic Table
Atomic Radii
Ionization Energy
Ionic Radii
Electron Affinity
Electronegativity
Atomic Radii
Atomic radius is the radius of an atom. It can
be defined as one-half the distance between
the nuclei of identical atoms that are bonded
together.
Atomic radii tend to decrease across the
period due to an increase in positive nuclear
charge.
Atomic radii tend to increase going down a
group due to addition of a primary energy
level.
Ionization Energy
Ionization energy is the energy required to
remove one electron from the neutral atom
of an element.
An ion is an atom or group of bonded
atoms that has a positive or negative
charge.
Any process that results in the formation of
an ion is referred to as ionization.
Ionization energy tends to increase
across a period.
Ionization energy tends to decrease
down a group.
The energy to remove the first electron
is called the first ionization energy.
Removing successive electrons
requires increased energy.
Electron Affinity
Neutral atoms can also acquire
electrons, not just lose them.
The energy change that occurs
when an electron is acquired by a
neutral atom is called the atom’s
electron affinity.
Most atoms release energy when
they acquire electrons.
Energy released is represented as
a negative number, energy
absorbed by a positive number.
Think of energy as a balance in a
bank account…
As you move across a period, it is
generally easier for the atoms to
acquire electrons. The values become
increasingly negative because the
easier it is for the atom to gain an
electron, the more energy will be
released.
As a general rule, electrons add
with greater difficulty going down a
group, but there are exceptions.
The main thing you need to know
is what electron affinity is.
Ionic Radii
Ionic radius is the radius of an ion.
Includes positive charged atoms (cations)
and negatively charged atoms (anions).
Ionic radii tend to decrease across a period
due to increase in positive charge to
negative charge ratio.
Ionic radii tend to increase down a group
due to the addition of a primary energy
level.
Electronegativity
In many compounds, one element will attract
the electrons more strongly than the other.
The uneven concentration of charge has a
significant effect on the chemical properties of
a compound and therefore it is useful to have
a measure of how strongly one atom attracts
the electrons of another atom within a
compound.
Electronegativity is a measure of the ability of an
atom in a chemical compound to attract
electrons.
Electronegativity tends to increase across a
period.
Electronegativity tends to decrease (or remain
the same) down a group.
The element with the highest electronegativity is
F and the lowest electronegativity is Fr.