The Periodic Table

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Transcript The Periodic Table

The Periodic Table
Chemistry
Chapter 6
Organizing the Elements

Chemists first used the properties of
elements to sort them into groups.
First classification system (1829) involved
grouping know elements into triads. Didn’t
work for all known elements.
 In the mid-1800s another system was
proposed, by a Russian Chemist Dmitri
Mendeleev.

Mendeleev’s Periodic table
He arranged elements in a periodic table.
 Elements in a periodic table are arranged
into groups based on a set of repeating
properties.
Mendeleev arranged the elements in his
table in order of increasing atomic
mass.


He left spaces in his table for yet
undiscovered elements and was even able to
predict their properties.
Problem with Mendeleev’s Table

What was wrong with Mendeleev’s table?


50 years after Mendeleev, another scientist
named Henry Mosley corrected the order of
the elements.
Mosley determined the atomic numbers of
all the elements and arranged them
based on their atomic numbers.

This allowed elements with similar physical
and chemical characteristics to fall into place
without exception.
Periodic Law
(Mendleev’s Law)
When elements are arranged in order of
increasing atomic number, there is a periodic
repetition of their physical and chemical
properties.
Elements
in a group have
similar properties!!!!!!! Why??
Same
# of Valence Electrons!
Metals, Nonmetals, & Metalloids

Elements can be classified into three
broad classes based on their general
properties.
Metals – good conductors, high luster,
ductile, and malleable. Most elements are
metals.
 Nonmetals – great variation in physical
properties. In general have properties
opposite of the metals.
 Metalloids – has properties of both metals &
nonmetals.

Atomic
number
6
C
carbon
12.011
2
Electrons in each
4
Energy level
Element
symbol
Element name
Atomic mass
Electron Configurations & Groups

Elements can be sorted into noble gases,
representative elements, transition
metals and inner transition metals based
on their electron configurations.

With the exception of the transition metals
(Group B), the group # lets you know how
many electrons are in the outer energy
level (# of valence electrons).
The Nobel Gases

The noble gases (inert gases) are the
elements in Group 8A of the periodic table.
S & P sublevels are completely filled with
electrons.
 Due to filled outer energy level, noble gases do
not readily react with other elements.

 Octet
Rule – atoms are most stable when they
have 8 electrons in their out level.
The Representative Elements

Elements in groups 1A through 7A are often
referred to as representative elements
because they display a wide range of physical
& chemical properties.

In atoms of rep. elements, the s & p sublevels of the
highest occupied energy level are not filled.
 Its
group number is equals the number of electrons in
the highest occupied energy level.
Transition Elements

Elements in the B groups, which provide a
connection between the two sets of rep.
elements, are referred to as transition elements;
two types of transition elements.
Transition metals – the highest occupied s
sublevel & nearby d sublevel contain electrons.
These elements are characterized by electrons in the d
orbitals.
 Inner transition metals (rare earth elements) –
The highest occupied s sublevel and nearby f
sublevel generally contain electrons. Characterized
by f orbitals that contain electrons.

Periodic Trends

Trends in Atomic Size

Atomic radius is one half of the distance
between the nuclei of two atoms of the same
element when the atoms are joined.
In general, atomic size increases from
top to bottom within a group and
decreases from left to right across a
period.
Atomic radius
•Atomic size increases in groups as atomic number increases.
•Atomic size decreases across a period from left to right.
Ions

Some compounds are composed of particles
called ions.
 An ion is an atom or group of atoms that has a +
or – charge. How does this occur?
or – ions form when electrons are transferred
between atoms.
+

An ion with a + charge is called a cation.
 always

smaller than the atoms from which they form
An ion with a – charge is called an anion.
 always
larger than the atoms from which they form.
Trends in Ionization Energy

If an atom is subject to enough energy
electrons may be removed; this is referred to as
ionization energy.

The energy required to remove the 1st electron from
an atom is called 1st ionization energy.
First ionization energy tends to decrease
from top to bottom with in a group and
increase from left to right across a
period.
Trends in Electronegativity

Electronegativity is the ability of an atom of an
element to attract electrons when the atom is in a
compound (electron affinity).

Used to predict the type of bond that will form during a
chemical reaction.
In general, the electronegativity values
decrease from top to bottom within a group.
For representative elements, the values tend to
increase from left to right across a period.
Summary of Trends
Atomic Radius trend: Largest toward SW corner of PT
Ionization Energy trend: Largest toward NE of PT
Electronegativity : Most favorable NE of PT