Transcript The Periodic Table

UNIT 3 Part 3:
The Periodic Table
1 Development of the Periodic Table
2 Reading the Periodic Table
3 Periodic Trends
1 Development of the Periodic
Table

Forerunners of the Periodic Table:

As Chemists and other scientists discovered
more and more Elements they needed a way
to keep track of them all. They need a method
of organization.
Döbereiner

In the early 1800’s J.W.
Döbereiner began classifying
elements into groups of three
(TRIADS ). The elements within the
triads share similar chemical
properties.

Ex:, lithium, sodium and potassium;
calcium, strontium and barium; AND
chlorine, bromine, and iodine are three
sets of triads
Newlands

Then in 1865, J.A.R Newlands
noticed that when the known
elements were arranged in
increasing atomic mass, regular
patterns would occur, AND they
would repeat every 8 elements
(law of octaves).
C
Mendeleev & Meyer

In 1869, Dmitri Mendeleev and
Lothar Meyer independently
published practically identical
schemes for classifying elements.
Mendeleev also noticed regular
patterns in the elements when they
were arranged by increasing atomic
mass. He produced the first periodic
table of elements and actually left
spaced to be filled in by
undiscovered elements. He was
able to predict there would be more
and about where they would go!
Mendeleev's Periodic Table of the
Elements
Periodic Law & Moseley

The periodic table Mendeleev
came up with was a good start
but it wasn’t perfect. It wasn’t
until 1913 when H.G.J. Moseley
determined that the correct way
to organize the periodic table is
by ATOMIC NUMBER. The
PERIODIC LAW- when elements
are arranged in order of
increasing atomic number, their
physical and chemical properties
show a periodic pattern.
Increases in atomic #
2 Reading the Periodic Table

Organizing the Squares:

When you look at the modern Periodic Table of
Elements (PTE) you will notice it is arranged in to
columns (vertically) and rows (horizontally). The
columns are called GROUPS or FAMILIES (18).
The rows are called PERIODS (7).
Periods
G
r
o
u
p
s
Labeling & Naming Groups

There are 3 methods for numbering the columns
(the American, the European and the IUPAC) The only 2 your
book uses is the American and the IUPAC.



AMERICAN: 1A, 2A, 3B-8B, 11B, 12B, 3A-8A
IUPAC: 1-18
The families are given names:

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1A- ALKALI METALS (except hydrogen)
2A-ALKALINE EARTH METALS
7A- HALOGENS
8A- NOBLE GASES aka: inert or inactive gases
Metals, Nonmetals and
Semimetals



Metals comprise the majority of the elements on the
periodic table. Commonly they are solids (except Hg),
malleable, conductors of heat and electricity, and are
ductile. Also they have other specific physical properties
such as luster or shine.
Nonmetals comprise the second largest type of
elements on the PTE. Many are also gases at room
temperature (except bromine), which is a liquid.
Semimetals (metalloids) have properties which are like
metals but also like nonmetals. Theses comprise the
smallest group on the PTE.
Electronic Configuration & the
PTE:

The electrons in the outermost energy
level are called the VALENCE electrons.


These are the electrons which will allow one
atom to interact with another atom and are
responsible for the elements chemical
properties.
The members in each group/family have
the same # of valence electron and that
is why they share similar chemical
properties.
The s-, p-, d- and f-block
Elements:


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S-BLOCK- groups 1&2
P-BLOCK- groups 13-18
D-BLOCK- groups 3-12 AKA: transition metals
F-BLOCK- the lanthanide & actinide series AKA: the inner transition
metals
p
S
d
f
3 Periodic Trends

As you move across the periodic table you
begin to see trends/predictable patterns in
the chemical and physical properties of the
elements. This phenomenon is known as
PERIODIC TRENDS.
Atomic Radius

Atoms have a radius called the ATOMIC
RADIUS- the distance from the center of the atom to
the farthest most/ highest energy level. This distance is
not exact. We estimate.

Contrary to popular opinion… atoms with more
electrons do not always have a larger atomic
radius. The trend is:


As you move from TOP to BOTTOM down the PTE
the radius grows larger
As you move from RIGHT to LEFT across the PTE
the radius grows larger
ATOMIC RADII
Largest

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The first trend has a simple explanation: as you
move down through a group/family you are
actually adding additional principle quantum
numbers (ex group 1: 1s, 2s, 3s, 4s, 5s, 6s, and 7s) with
corresponding increases in the number of
electrons from 1, 3, 11, 19, 37, 55, and 87. Ergo
the radius must increase to accommodate these
increases in electrons.
The second trend is more complex. As you
move across a period you are increasing both
the # of electrons as well as the number of
protons. This allows the protons to pull more
strongly on the electrons in the same energy
level thus shrinking the radius.
Ionic Size

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We know atoms can gain or lose electron to
form Ions. Depending whether they gain (becoming
negatively charged) or lose (becoming positively charged) will
depend on whether ionic size will shrink or grow.
When you gain electrons- the ionic size grows
As you lose electrons- the ionic size shrinks (more
protons pulling on fewer electrons= shrinking).

The trend is complex- Figure 5-17 p.176
Ionization Energy


The energy needed to strip an atom of an
electron is called IONIZATION ENERGY.
Typically measured in J/atom you can think of it
as a measure of how strongly the atom holds on
to its valence electrons.
Its trend is exactly opposite the atomic radius
trend:


As you move from BOTTOM to TOP the ionization
energy increases
As you move from LEFT to RIGHT the ionization
energy increases
Ionization Energy
Hardest
Easiest
Successive Ionization Energies
What if you want to pull more than one electron from
an atom? What happens to the ionization energy
for the 2nd, 3rd, or even 4th electron?
 Successive ionization energies refer to the energy
required to remove the 2nd, 3rd and more
electrons. The energy required increases, partially
because of electron-electron repulsion, however it
is not a linear increase. Examine figure 5-20 p.180

*note* there is always a dramatic increase in ionization
energy when you start extracting electrons from the inner
noble gas core.
Electron Affinity

If you can pull an electron from an atom… you
can ADD one to an atom. When an atom gains
an electron energy is REQUIRED! They are
reported in J/mol (NOT ATOMS). Not all energy
used in the addition of an electron is gained by
the atom that the electron is being added to. It is
possible to have + and – electron affinities. If
the value is + then the atom gaining the electron
is also gaining energy. If the value is – then the
atom gaining the electron is releasing energy.

General rule for the electron affinity trendNonmetals have more NEGATIVE electron
affinities than the metals.
 The exception is the noble gases.

 **
OCTET RULE** atoms tend to gain, lose or
even share electrons in order to have a full set of
valence electrons (8)- to have a full s and p orbital.
More Negative
Electronegativity
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This is the ability of an atom to attract an
electron in a chemical bond. This is a unit less
value.
The least electronegative elements are in the
LOWER LEFT CORNER of the PTE. The trend
is the same as Ionization energy.
When bonds occur… the atom with the higher
electronegativity would attract the electron more
strongly. Depending on the discrepancy of the
attraction there will form different types of bonds
(Ch7).
More electronegative
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