Transcript What makes a group of elements

Chapter 6
The Periodic Table
What makes a group of
elements?
What makes a group of elements




Some elements are gases, some are
liquids, and others are solids.
Some are colored and some are
colorless.
Though some elements look
dissimilar, they react similarly.
Samples of the compounds NaCl,
NaBr, and NaI show similar physical
characteristics.
Periodic Pattern


In 1865, the English chemist John
Newlands arranged the first 16
elements known at the time in order
of increasing atomic mass.
When Newlands placed the elements in
two rows, he observed that the elements
in each column had similar chemical and
physical properties.
Li
Na
Be
Mg
B
Al
C
Si
N
P
O
S
F
Cl
Because the chemical and physical
properties repeated with the eight
element, Newlands called this pattern
the law of octaves.



In 1870, the Russian chemist Dmitri
Mendeleev made use of this scheme and
other information to produce the first
orderly arrangement of all 63
elements known at the time.
Mendeleev wrote the symbol for each
element on a card along with its relative
atomic mass. Arranged elements in
order of increasing atomic mass.
He is considered the “father of the
periodic table”



Periodic law- states that the
physical and chemical
properties of the elements are
periodic functions of their
atomic numbers.
The most common periodic table is
based on the periodic law.
The periodic table contains a wealth
of information about individual
elements.
The modern periodic table



The rows of the periodic table
are called periods.
Elements in a period have similar
electron configurations.
For example, the first period has
two elements, hydrogen and
helium, and the electrons of these
elements occupy the 1s orbital.



The columns of the periodic
table are called groups.
The elements within a group
have properties in common but
with some gradation.
For example, all Group I elements
are solid at room temp., and are
good conductors of electricity.
METALS


It is convenient to divide the
periodic table into two distinct
regions, metals and nonmetals.
The metals include all members
of Groups 1 through 12 as well
as some elements of Group 13
through 16.



All metals are good conductors of
electricity, and their conductivity
increases as temperature decreases.
Except for Mercury, all metals are
solid at room temperature.
Elements in Group 3 through 12
including the two long rows below
the main table, are called transition
metals.


As you move from left to right
across the transition metals, you
will see that the electrons are
usually added to d orbitals.
For this reason the transition metals
are sometimes referred to as the dblock elements.
Metals that are good conductors of
electricity are also good conductors of
heat.



In general, poor electrical
conductors are poor heat
conductors.
Think about it.
This means that a mechanism by
which electricity is conducted must
also be closely connected with the
mechanism for conduction of heat.


You don’t have to look that far for
the agent that causes conductivity
in metals.
Ever since the discovery of
electrons it has been known that
electrons are responsible for the
conduction of electricity by
metals.




They are responsible for the
conduction of heat as well.
However, metals are the only elemental
substances that are good conductors of
electricity.
This must mean that at least some of
the electrons in metals must be in a
different configuration than in
nonmetals.
Those electrons are free to move
through the metal in all directions.



Metals can have extremely high
melting points, and some can
have very low melting points.
Metals can be extremely
reactive while others do not
react at all.
Metals can be strong and
durable.

Metals can also be ductile (wire)
and malleable (metal sheets).
NONMETALS

The second region of the
periodic table contains the
nonmetals.


The nonmetals include all of
Groups 17 and 18 as well as
some members of Groups 14
through 16.
The characteristics shared by all
nonmetals is that they are poor
conductors of electricity.



Nonmetals may be gases,
liquids, or solids at room
temperature.
Along the stair-step line
separating metals from
nonmetals are the elements
known as semiconductors, or
metalloids.
Metalloids are solids at room temp.
Main-group elements

Groups 1, 2, and 13 through 18 are
referred to as the main-group
elements.


The electron configuration of elements
within each group is quite consistent.
For example, all of the elements in Group
14 have four electrons in their outermost
shell.


The main-group elements
include gases, liquids, solids,
metals, and nonmetals.
The main group elements silicon
and oxygen account for four of
every five atoms found on or near
the Earth’s surface.


Four groups within the main-group
elements have special names.
These are alkali metals (Group 1),
the alkaline-earth metals (Group
2), the halogens (Group 17), and
the noble gases (Group 18).
Group 1 (Alkali Metals)


Alkali metals- react with water
to produce alkaline solutions
and because they have metallic
properties.
The term alkali dates back to
ancient times, when people
discovered that wood ashes mixed
with water produces a slippery
solution that can remove grease.


The alkali metals are so soft
that they can be cut with a
knife. The freshly cut surface of an
alkali metal is shiny, but it dulls
quickly as the metal reacts with
oxygen and water in the air.
All of the alkali metals are excellent
conductors of electricity.
Group 2 (Alkaline-Earth Metals)



The elements of Group 2 are called the
alkaline-earth metals.
Compared with the alkali metals, the
alkaline-earth metals are harder,
denser, stronger, and have higher
melting points.
The best known alkaline-earth metal is
calcium. Calcium compounds such as
those in limestone and marble, are
common in the Earth’s crust.
Group 17 (Halogens)



The elements of Group 17 are the
halogens.
Fluorine, chlorine, bromine, iodine,
and astatine.
The halogens combine with
most metals to produce the
compounds known as salts.


The word halogen is derived from
Greek and means “salt former”
In common table salt, sodium
chloride, the halogen chlorine has
reacted with the alkali metal sodium
to form sodium and chloride ions.
Group 18 (Noble Gases)


The Group 18 elements, are called
noble gases.
Noble gas atoms are
characterized by an octet of
electrons in the outermost
energy level.
Hydrogen


Hydrogen is in a class by itself
Hydrogen is the most common
element in the universe. It behaves
unlike other elements because it
has just one proton and one
electron. This distinguishes
hydrogen from all of the other
elements.
6-2 , 6-3
What trends are found in the period
table?



You have read that elements are
arranged in the periodic table in
order of increasing atomic number.
The elements are further organized
into groups and periods.
The arrangement of the periodic
table also reveals trends in the
chemical and physical properties of
the elements.



What is a trend?
A trend is a predictable change
in a particular direction.
For example, as you move down
group 1, reactivity increases for
each element.
Periodic Trends in atomic radii


The exact size of the atom is
difficult to determine.
Bond radius-half the distance
from center to center in two like
atoms bonded together.


The Vand der Waals radius is
seldom used to state the size of
atoms, but data are only available
for only a few main group elements.
Van der Waals radius- half the
distance between the nuclei in
adjacent non-bonded molecules.


We will be using bond radius to
determine the size of atoms
because there is more information
available.
Measuring bond radius is a
useful way to compare sizes of
atoms.
Electron Shielding




The electrons in the inner energy
levels are between the nucleus and
the outermost valence electrons
This shields the valence electrons
from the full charge of the nucleus.
This phenomenon is called electron
shielding.
Because the valence electrons are not
subject to the full charge of the nucleus,
they are not held as close to the nucleus.
Atomic radius increases as you
move down a group.


There is a trend toward larger
radii as you proceed down a
group.
This is caused by the addition of
another main energy level as
you move from one period to the
next.
Atomic radius decreases as you
move across a period.

From left to right across a
period, each atom has one more
proton and one more electron
that the element before it.



The additional electrons are going
into the same energy level.
Electrons in an outer energy level
do not screen the other electrons in
that energy level very effectively.
Meanwhile the nuclear charge is
increasing as protons are added and
the electrons are pulled closer to
the nucleus, reducing the size of the
atom.
In other words…



There is a gradual decrease in the
atomic radii across the second
period from Li to Ne
The trend to smaller atoms across a
period is caused by the increasing
positive charge of the nucleus
Increased pull results in a decrease
in atomic radii
Ionization energy, Electron affinity, and
Electronegativity.



Recall that atoms are
electrically neutral.
But if you add enough energy,
the atom may lose an electron
to become a positive ion.
Imagine that you can reach into an
atom and remove on of its valence
electrons, creating an ion.


The energy you used to remove that
electron called ionization energy.
Ionization energy- the amount
of energy needed to remove an
outer electron from a specific
atom or ion in its ground state.
Ionization energy increases across a
period and decreases down a group

Group 1 elements have the lowest first
ionization energy



They lose electrons most easily
Major reason for the high reactivity of the
Group 1
The Group 18 elements, the noble gases,
have the highest ionization energies.

They do not lose their electrons easily; the
low reactivity is partly due to the difficulty to
removed an electron




In general, ionization energies of
the main-group elements increase
across each period
Increase is due to increase in
nuclear charge of the protons
Higher charge attracts electrons in
the same energy level
In general, nonmetals have higher
ionization energies than metals do


As you move down a group, the
number of energy levels
between the nucleus and the
valence electrons increases and
the outermost electrons are
farther from the nucleus.
The nuclear charge stays the same
so the electrons are held less tightly
to the nucleus , and less energy is
required to remove one of them.
Periodic Trends in Electron Affinity




The ability of an atom to attract and
hold an electron is called electron
affinity.
You may wonder why a neutral atom
would attract electrons in the first place.
The answer is that electrons in the
orbitals generally do not shield the
nuclear charge to a full 100%.
An approaching electron may experience
the a net pull because the nuclear charge
is greater than 0.


Across a period, shielding
remains the same, but nuclear
charge increases.
Therefore, the atoms attraction
for extra electrons increases.



Going down a group, both
shielding and nuclear charge
increase.
However, the shielding effect offsets
the increase in nuclear charge.
Therefore, the atom’s attraction
for extra electrons decreases.
Electronegativity



Electronegativity is a measure
of the ability of an atom in a
chemical compound to attract
electrons
The most electronegative
element is Fluorine
The values are assigned 0-4,
Fluorine is 4
Period Trends



Electronegativities tend to
increase across each period,
although there are exceptions
Electronegativities tend to
either decrease down a group or
remain about the same
Noble gases – some do not form
compounds so they cannot be
assigned electronegativities
Positive Ions



A positive ion is known as a
cation
The formation of a cation by loss
of one or more electrons always
leads to a decrease in atomic
radius because of the removal of
the highest-energy-level electrons
results in a smaller electron cloud
The remaining electrons are
drawn closer to the nucleus
Negative Ions

A negative ion is known as an
anion.
The formation of an anion by the
addition of one or more
electrons always leads to an
increase in atomic radius

Electron cloud spreads out

Period Trends cont..



The metals on the left form
cations
Nonmetals on the right tend to
form anions
Cationic radii decreases across a period


The electron cloud shrinks due to the
increasing nuclear charge acting on the
electrons in the same main energy level
Anionic radii decrease across a period
Group Trends



The outer electrons in both cations
and anions are higher in energy
levels as you read going down a
group
Atomic radii gradually increases
going down a group
Ionic radii gradually increases going
down a group