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1. The Periodic Table

2. • Introduction
– The periodic table is made up of rows of elements
and columns.
– An element is identified by its chemical symbol.
– The number above the symbol is the atomic number
– The number below the symbol is the rounded
atomic weight of the element.
– A row is called a period
– A column is called a group

3. Organizing the Elements
• Chemists used the properties of
elements to sort them into groups.
• JW. Dobreiner grouped elements into
triads.
• A triad is a set of three elements with
similar properties.

4. Mendeleev’s Periodic Table
• In 1869, a Russian
chemist and
teacher published a
table of the
elements.
• Mendeleev arranged
the elements in the
periodic table in
order of increasing
atomic mass.

5. Henry Moseley
1887 - 1915
In 1913, through his work with X-rays, he
determined the actual nuclear charge
(atomic number) of the elements*. He
rearranged the elements in order of
increasing atomic number.
*“There is in the atom a fundamental
quantity which increases by regular
steps as we pass from each element to
the next. This quantity can only be the
charge on the central positive nucleus.”

6. The Periodic Law
In the modern periodic
table elements are
arranged in order of
increasing atom ic
num ber.
Periodic Law states:
When elements are
arranged in order of
increasing atomic
number, there is a
periodic repetition
of their physical and
chemical properties.

7. • The elements can be grouped into
three broad classes based on their
general properties.
• Three classes of elements are Metals,
Nonmetals, and Metalloids.
• Across a period, the properties of
elements become less metallic and
more nonmetallic.

8. Properties of Metals
• Metals are good conductors
of heat and electricity.
• Metals are shiny.
• Metals are ductile (can be
stretched into thin wires).
• Metals are malleable (can be
pounded into thin sheets).
• A chemical property of metal
is its reaction with water
which results in corrosion.
• Solid at room temperature
except Hg.

9. Properties of Non-Metals
• Non-metals are poor
conductors of heat and
electricity.
• Non-metals are not ductile
or malleable.
• Solid non-metals are
brittle and break easily.
• They are dull.
• Many non-metals are
gases.
Sulfur

10. Properties of Metalloids
• Metalloids (metal-like) have
properties of both metals and
non-metals.
• They are solids that can be
shiny or dull.
• They conduct heat and
electricity better than non-
metals but not as well as
metals.
• They are ductile and
malleable.
Silicon

11. Groups Periods
Groups Periods
 Columns of elements are
Columns of elements are
called groups or families.
called groups or families.
 Elements in each group
Elements in each group
have similar but not
have similar but not
identical properties.
identical properties.
 For example, lithium (Li),
For example, lithium (Li),
sodium (Na), potassium
sodium (Na), potassium
(K), and other members of
(K), and other members of
group IA are all soft, white,
group IA are all soft, white,
shiny metals.
shiny metals.
 All elements in a group
All elements in a group
have the same number of
have the same number of
valence electrons.
valence electrons.
 Each horizontal row of
Each horizontal row of
elements is called a period.
elements is called a period.
 The elements in a period
The elements in a period
are not alike in properties.
are not alike in properties.
 In fact, the properties
In fact, the properties
change greatly across even
change greatly across even
given row.
given row.
 The first element in a period
The first element in a period
is always an extremely
is always an extremely
active solid. The last
active solid. The last
element in a period, is
element in a period, is
always an inactive gas.
always an inactive gas.

12. Hydrogen
Hydrogen
 The hydrogen square sits atop group AI, but
The hydrogen square sits atop group AI, but
it is not a member of that group. Hydrogen is
it is not a member of that group. Hydrogen is
in a class of its own.
in a class of its own.
 It’s a gas at room temperature.
It’s a gas at room temperature.
 It has one proton and one electron in its one
It has one proton and one electron in its one
and only energy level.
and only energy level.
 Hydrogen only needs 2 electrons to fill up its
Hydrogen only needs 2 electrons to fill up its
valence shell.
valence shell.

13. 6.2 Classifying the Elements
6.2 Classifying the Elements
The periodic table
The periodic table
displays the symbols
displays the symbols
and names of the
and names of the
elements along with
elements along with
information about the
information about the
structure of their
structure of their
atoms.
atoms.

14.  Four chemical groups
Four chemical groups
of the periodic table:
of the periodic table:
2.
2. alkali metals (IA)
alkali metals (IA)
3.
3. alkaline earth metals
alkaline earth metals
(IIA),
(IIA),
4.
4. Halogens (VII),
Halogens (VII),
5.
5. Noble
Noble gases
gases (VIIIA).
(VIIIA).

15. Alkali Metals
Alkali Metals
 The alkali family is found in
The alkali family is found in
the first column of the
the first column of the
periodic table.
periodic table.
 Atoms of the alkali metals
Atoms of the alkali metals
have a single electron in their
have a single electron in their
outermost level, in other
outermost level, in other
words, 1 valence electron.
words, 1 valence electron.
 They are shiny, have the
They are shiny, have the
consistency of clay, and are
consistency of clay, and are
easily cut with a knife.
easily cut with a knife.

16. Alkali Metals
Alkali Metals
 They are the most
They are the most
reactive metals.
reactive metals.
 They react violently
They react violently
with water.
with water.
 Alkali metals are
Alkali metals are
never found as free
never found as free
elements in nature.
elements in nature.
They are always
They are always
bonded with another
bonded with another
element.
element.

17. Alkaline Earth Metals
Alkaline Earth Metals
 They are never found uncombined in nature.
They are never found uncombined in nature.
 They have two valence electrons.
They have two valence electrons.
 Alkaline earth metals include magnesium and
Alkaline earth metals include magnesium and
calcium, among others.
calcium, among others.

18. Transition Metals
Transition Metals
 Transition Elements
Transition Elements
include those elements in
include those elements in
the B groups.
the B groups.
 These are the metals you
These are the metals you
are probably most
are probably most
familiar: copper, tin, zinc,
familiar: copper, tin, zinc,
iron, nickel, gold, and
iron, nickel, gold, and
silver.
silver.
 They are good conductors
They are good conductors
of heat and electricity.
of heat and electricity.

19. Transition Metals
Transition Metals
 The compounds of transition metals are usually brightly
The compounds of transition metals are usually brightly
colored and are often used to color paints.
colored and are often used to color paints.
 Transition elements have 1 or 2 valence electrons, which
Transition elements have 1 or 2 valence electrons, which
they lose when they form bonds with other atoms. Some
they lose when they form bonds with other atoms. Some
transition elements can lose electrons in their next-to-
transition elements can lose electrons in their next-to-
outermost level.
outermost level.

20. Transition Elements
Transition Elements
 Transition elements
Transition elements have properties
have properties
similar to one another and to other metals,
similar to one another and to other metals,
but their properties do not fit in with those
but their properties do not fit in with those
of any other group.
of any other group.
 Many transition metals combine
Many transition metals combine
chemically with oxygen to form
chemically with oxygen to form
compounds called oxides.
compounds called oxides.

21. Representative Elements
Representative Elements
 Groups 1A – 7A.
Groups 1A – 7A.
 Elements are refered to as representative
Elements are refered to as representative
elements because they display a wide
elements because they display a wide
range of physical and chemical properties.
range of physical and chemical properties.
 For any representative element, its group
For any representative element, its group
number equals the number of electrons in
number equals the number of electrons in
the highest occupied energy level.
the highest occupied energy level.

22. Trends in the periodic
Trends in the periodic
table:
table:
Ionization Energy
Ionization Energy
Atomic Radius
Atomic Radius
Electron Affinity
Electron Affinity
Electronegativity
Electronegativity

23. Sizes of Atoms
Sizes of Atoms
The bonding atomic
The bonding atomic
radius is defined as
radius is defined as
one-half of the
one-half of the
distance between
distance between
covalently bonded
covalently bonded
nuclei.
nuclei.

24. Atomic Radius Trend
Atomic Radius Trend
 Group Trend – As you go
Group Trend – As you go down a column
down a column,
,
atomic radius increases.
atomic radius increases.
As you go down, e
As you go down, e-
-
are filled into orbitals that are
are filled into orbitals that are
farther away from the nucleus (attraction not
farther away from the nucleus (attraction not
as strong).
as strong).
 Periodic Trend – As you go
Periodic Trend – As you go across a period
across a period (L
(L
to R),
to R), atomic radius decreases.
atomic radius decreases.
As you go L to R, e
As you go L to R, e-
-
are put into the same orbital,
are put into the same orbital,
but more p
but more p+
+
and e
and e-
-
total (more attraction =
total (more attraction =
smaller size).
smaller size).

25. Atomic Radius
Atomic Radius

27. Ionic Radius Trend
Ionic Radius Trend
 Metals
Metals – lose e
– lose e-
-
, which means more p
, which means more p+
+
than e
than e-
-
(more attraction) SO…
(more attraction) SO…
Ionic Radius
Ionic Radius <
< Neutral Atomic Radius
Neutral Atomic Radius
 Nonmetals
Nonmetals – gain e
– gain e-
-
, which means more e
, which means more e-
-
than p
than p+
+
(not as much attraction) SO…
(not as much attraction) SO…
Ionic Radius
Ionic Radius >
> Neutral Atomic Radius
Neutral Atomic Radius

28. Sizes of Ions
Sizes of Ions
 Ionic size depends
Ionic size depends
upon:
upon:
 Nuclear charge.
Nuclear charge.
 Number of
Number of
electrons.
electrons.
 Orbitals in which
Orbitals in which
electrons reside.
electrons reside.

29. Sizes of Ions
Sizes of Ions
 Cations are
Cations are
smaller than their
smaller than their
parent atoms.
parent atoms.
 The outermost
The outermost
electron is
electron is
removed and
removed and
repulsions are
repulsions are
reduced.
reduced.

30. Sizes of Ions
Sizes of Ions
 Anions are larger
Anions are larger
than their parent
than their parent
atoms.
atoms.
 Electrons are
Electrons are
added and
added and
repulsions are
repulsions are
increased.
increased.

31. Sizes of Ions
Sizes of Ions
 Ions increase in size
Ions increase in size
as you go down a
as you go down a
column.
column.
 Due to increasing
Due to increasing
value of
value of n
n.
.

32. Metals versus Nonmetals
Metals versus Nonmetals
 Metals tend to form cations.
Metals tend to form cations.
 Nonmetals tend to form anions.
Nonmetals tend to form anions.

33. Background
Background
 Electrons can jump between shells (Bohr’s
Electrons can jump between shells (Bohr’s
model supported by line spectra)
model supported by line spectra)
 The electrons can be pushed so far that
The electrons can be pushed so far that
they escape the attraction of the nucleus
they escape the attraction of the nucleus
 Losing an electron is called ionization
Losing an electron is called ionization
 An ion is an atom that has either a net
An ion is an atom that has either a net
positive or net negative charge
positive or net negative charge
 Q: what would the charge be on an atom
Q: what would the charge be on an atom
that lost an electron? Gained two electrons?
that lost an electron? Gained two electrons?
 A: +1 (because your
A: +1 (because your losing
losing a -ve electron)
a -ve electron)
 A: -2 (because you gain 2 -ve electrons)
A: -2 (because you gain 2 -ve electrons)

34. Ionization Energy
Ionization Energy
 Amount of energy required to remove an
Amount of energy required to remove an
electron from the ground state of a
electron from the ground state of a
gaseous atom or ion.
gaseous atom or ion.
 First ionization energy is that energy required
First ionization energy is that energy required
to remove first electron.
to remove first electron.
 Second ionization energy is that energy
Second ionization energy is that energy
required to remove second electron, etc.
required to remove second electron, etc.

35. Ionization Energy
Ionization Energy
 Group Trend – As you go
Group Trend – As you go down a column
down a column,
,
ionization energy decreases.
ionization energy decreases.
As you go down, atomic size is increasing (less
As you go down, atomic size is increasing (less
attraction), so easier to remove an e
attraction), so easier to remove an e-
-
.
.
 Periodic Trend – As you go
Periodic Trend – As you go across a period
across a period (L to
(L to
R),
R), ionization energy increases.
ionization energy increases.
As you go L to R, atomic size is decreasing (more
As you go L to R, atomic size is decreasing (more
attraction), so more difficult to remove an e
attraction), so more difficult to remove an e-
-
(also, metals want to lose e
(also, metals want to lose e-
-
, but nonmetals do
, but nonmetals do
not).
not).

36. Ionization Energy
Ionization Energy
 It requires more energy to remove each
It requires more energy to remove each
successive electron.
successive electron.
 When all valence electrons have been removed,
When all valence electrons have been removed,
the ionization energy takes a quantum leap.
the ionization energy takes a quantum leap.

37. Trends in First Ionization
Trends in First Ionization
Energies
Energies
 As one goes down a
As one goes down a
column, less energy
column, less energy
is required to remove
is required to remove
the first electron.
the first electron.
 For atoms in the same
For atoms in the same
group,
group, Z
Zeff
eff is essentially
is essentially
the same, but the
the same, but the
valence electrons are
valence electrons are
farther from the
farther from the
nucleus.
nucleus.

38. Electronegativity
Electronegativity
 Electronegativity-
Electronegativity-
tendency of an
tendency of an
atom to attract e
atom to attract e-
-
.
.

39. Electronegativity Trend
Electronegativity Trend
 Group Trend – As you go
Group Trend – As you go down a column
down a column,
,
electronegativity decreases.
electronegativity decreases.
As you go down, atomic size is increasing, so less
As you go down, atomic size is increasing, so less
attraction to its own e
attraction to its own e-
-
and other atom’s e
and other atom’s e-
-
.
.
 Periodic Trend – As you go
Periodic Trend – As you go across a period
across a period (L to R),
(L to R),
electronegativity increases.
electronegativity increases.
As you go L to R, atomic size is decreasing, so there is
As you go L to R, atomic size is decreasing, so there is
more attraction to its own e
more attraction to its own e-
-
and other atom’s e
and other atom’s e-
-
.
.

40. Electronegativity
Electronegativity