Transcript Ch 6 Notes

Russian chemist Dmitri Mendeleev
placed the known elements in order of
increasing atomic mass.
 When he did this he noticed that the
elements’ properties repeated in a
regular pattern, or a periodic pattern.
 Mendeleev placed the known elements
in a table, where he arranged elements
into columns with similar properties.

Mendeleev predicted the properties of
several elements that were unknown at
the time.
 Many of his predictions were correct and
were well accepted by the scientific
community.


The only changes to Mendeleev’s
periodic table were the addition of
newly found elements and that the table
was organized by atomic number rather
than atomic mass.
The horizontal rows on the periodic table
are called periods. This is because the
properties begin to repeat in each new
row.
 The columns on the periodic table are
called groups (or families) and the
elements in the groups have similar
properties.
 Groups are designated with a number
and the letter A or B.

The main group elements can be found
in Groups 1A through 8A.
 The main group elements (Group A
elements) are also called representative
elements.


The group B elements are called the
transition elements.
Elements can be divided into three main
classes.
 Those classes are:

› Metals
› Metalloids
› Nonmetals

Metals have the following properties.
They are generally:
› Shiny solids
› Good conductors of heat and electricity
Group 1A has the name alkali metals
 Group 2A has the name alkaline earth
metals.

Elements to the right of the heavy stairstep line on the periodic table are called
nonmetals.
 Nonmetals generally have the properties
that they are generally gases or brittle
solids at room temperature.
 Group 7A are called halogens.
 Group 8A are called the noble gases.

Many of the elements that border the
stair step are metalloids.
 Metalloids share properties in between
those of metals and nonmetals.


1.
2.
3.
4.
5.
6.
7.
Column A
Strontium
Chromium
Iodine
Nitrogen
Argon
Rubidium
Silicon

a.
b.
c.
d.
e.
f.
g.
Column B
Halogen
Noble gas
Alkaline earth metal
Metalloid
Alkali metal
Representative
element
Transition Element
Scientists now understand that the
repetition of properties of elements
occurs because the electron
configurations of atoms repeat.
 The arrangement of elements in the
periodic table reflects the electron
structures of atoms.

For representative elements, the Group
Number in front of the A tells the number
of valence electrons in the atoms in the
column.
 Also the period number (or row number)
of a representative element tells the
energy level of the valence electrons.






The periodic table is divided into blocks that
correspond to the energy sublevel being
filled as you move across a period.
Groups 1A and 2A are the s block.
Groups 3A- 8a are the p block.
The B elements represent the d block;
however remember to go down one for the
first quantum number.
The rows that are removed from the table
and placed at the bottom represent the f
block. Go down 2 numbers for the row
number.

1.
2.
3.
4.
Without using the periodic table,
determine the group, period, and block
in which an element with each of the
following electron configurations is
found.
[He] 2s2 2p5
[Ar] 4s2
[Kr]5s2 4d10 5p3
[Ar]4s2 3d3
Section 6.3
Periodic Trends
The electron structure of an atom
determines many of its chemical and
physical properties.
 There are several trends that can be
observed using the periodic table.

The atomic radius is a measure of the size of
an atom. The larger the radius, the larger
the atom.
 As you move across a period the atom
decreases in size.
 This is due to the increasing positive charge
of the nucleus while electrons are being
added, but the orbitals are close in energy.
 The increased nuclear charge pulls the
outermost electrons closer to the nucleus,
making the atom smaller.

As you move down a group, atomic radii
increases.
 This is due to the addition of a larger
energy level each time.
 Electrons in higher levels are located
farther from the nucleus than those in the
lower energy levels.

Trend

1.
2.
3.
4.
5.
From each of the following pairs, predict
which atom is larger.
Mg, Sr
Sr, Sn
Ge, Sn
Ge, Br
Cr, W
Ion Formation
The driving force that makes reactions
happen is ion formation.
 Ions form when atoms gain or lose
electrons.
 When electrons are gained or lost the
resultant ion has a positive or negative
charge.

Cations
When atoms lose electrons and form
positively charged ions, they become
smaller, and are called cations.
 They become smaller because the loss
of the electrons means that the number
of protons is greater than the number of
electrons.
 Therefore, the electrons will be pulled
more tightly to the nucleus, and the
outer electrons will feel the pull of the
nucleus more strongly than before.

Anions
When atoms gain electrons they form
negatively charged ions, called anions.
 The added electron makes the ionic
radius increase, because of repulsion
and the weaker pull of the nucleus on
each electron.

Trend in Ionic Radius
As you move from left to right across the
periodic table, your positively forming
ions, Groups 1A-4A get smaller, between
4A and 5A the ions get bigger but
Groups 5A- 8A the ions get smaller again.
 As you move from top to bottom the ions
get larger in each column because of
the addition of energy levels.

Ionization Energy
To form a positive ion, an electron must
be removed from a neutral atom.
 Removing the electron requires energy.
That energy must overcome the
attraction between the positive charge
in the nucleus and the negative charge
of the electron.
 This energy, known as ionization energy is
defined as the energy required to
remove an electron from an atom in the
gaseous state.

First Ionization Energy
The first ionization energy is the amount
of energy required to remove the first
electron from the outer shell of the atom.
 Remember that as you move across a
period, you increase atomic number
and therefore add more positive charge
to the atom.
 That addition makes it more difficult to
remove the electron and therefore more
energy is required.

Ionization Energy
Trend in Ionization Energy
For the first ionization energy, as you
move across a period, from left to right,
the ionization energy increases.
 As you move down a group, the
ionization energy generally decreases.
Why?

Trend
Additional Ionization Energies
It is also possible to remove electrons
after removing the first, or furthermost
electron from the electron cloud.
 However, once you remove the first
electron, much more energy is required
to remove the second, and third, and
fourth, etc.

Octet Rule
Remember, when Newland tried to
design his periodic table he came up
with a law of octaves, which wasn’t
accepted.
 We see using ionization energies that his
predictions were every correct.
 When atoms/ions have eight electrons in
their outer shell, they are much more
stable.

Octet Rule
The octet rule has come to be one of the
most important principles in chemistry.
 It says that atoms tend to gain, lose or
share electrons in order to acquire a full
set of eight valence electrons.
 Use this to predict what kinds of ions will
form: elements on the left will form
positive ions and elements on the right
negative ions.

Electronegativity
The electronegativity of an element tells
about its ability to attract electrons in a
chemical bond.
 The more electronegative an element,
the more electron loving it is.
 Noble gases form few compounds and
so virtually lack electronegativity.

Trend
As you move from top to bottom down a
group, electronegativity generally
decreases.
 As you move from left to right,
electronegativity generally increases.
 So what is the most electronegative
element?

Trend