Transcript TREND

THE PERIODIC TABLE
Objectives:
Examine the progression of periodicity
Alkali metals
Transition metals
alkaline earth
metals
Metalloids
(semimetals)
“s” group
“d” block
Noble
gases
Inner transition
metals
“f” group
Nonmetals
“p” block
Halogens
Periodic Patterns

ALKALI METALS (part of the “s” group of elements)
~ all are in shiny solid form but are quite soft
~ form the 1st group of metals on the periodic table
~ highly reactive elements based upon their electron
configurations (ns1)
Other Characteristics:
 malleable and ductile; low density and melting points
 good conductors of electricity;
very soluble as
comps.
Alkaline Earth Metals
PART OF THE “S” GROUP JUST
LIKE ALKALI METALS
 Belong
to the second group
of metals on the periodic
table.
 Harder, more dense, and
stronger than there group 1
counterparts
 Not as reactive as the Alkali
metals due to the metals in
this group having 2 electrons
in their valence shell.
 This gives them the
configuration of “ns2” for
these metals.
Be
Mg
Ca
Sr
Ba
Ra
Transition Metals
The “d” - block elements
 Transition metals begin in
the 4th period after the
alkaline earth metals.
 Metallic elements with
varying properties.
 Not nearly as reactive as
group 1 and 2 elements.
 Fill their sublevels
differently than do the
Main group elements.
Valuable as
structurally useful
materials!
Lanthanoids & Actinoids
The ‘f’-group is broken into two classifications
LANTHANOIDS

Composed of the elements
with atomic numbers 58
through 71

Electrons are being added to
the “ 4f “ sublevel

Shiny reactive metals with
practical uses
ie. dots in TV tubes
ACTINOIDS



Composed of the elements
with the atomic numbers 90
through 103
They fill the “ 5f “ sublevel
All are radioactive with an
unstable nucleus
Nonmetals & Metalloids (semi-metals)
NONMETALS!
 Generally are gases at room temperature (or brittle solids)
 Poor conductors of heat and electricity
 Have more electrons in their outer level than metals
METALLOIDS!
 Properties of both metals and nonmetals
 Will give up (electron donor) electron(s) when reacted with a
nonmetal, and will accept (electron acceptor) electron(s) when
reacted with a metal
 In general, more like nonmetals than metals
 Considered semiconductors
Periodic Trends
Trends (we will study) – atomic radius
(ionic radius), ionization energy,
electronegativity, electron affinity
 Trends are looked at from top to bottom of a
column and from left to right in a period
(row)
 Trends show patterns of atoms properties
(relationships among elements)

Trend Number 1
Li
Na
ATOMIC RADIUS
 Atomic radius is the half the
distance between the nuclei of two
like atoms.
K
Rb
TREND:
the trend for atomic radius
shows us the size of the atom will
increase as we move down a column
WHY: more levels and orbitals,
greater distances from the nucleus
Cs
Fr
ATOMIC RADIUS
TREND:

the atomic radii will decrease from the left to the
right in a period
WHY:
Effective Nuclear Charge (also applies to what takes
place from top to bottom of a column)

positive charge felt by the outermost electrons of
an atom

atomic # - # of inner complete level electrons

The larger the ENC, the greater the attraction of
electrons to the nucleus
Shielding
- the ability of other electrons,especially inner
electrons, to lessen the nuclear charge of the outer
electron(s)
THE TREND!
 The trend shows the
increase of radii
down a group and
decrease of radii
across a period.
NUMBER TWO!
IONIZATION ENERGY
IONIC BOND
 bond formed between two ions by the
transfer of electrons
Ions: How do they form?

In certain types of bonding, the atom
will “lose” or “gain” an electron(s)

When an atom loses or gains
electrons, it is called an ion
Magnesium
Magnesium

Atoms that lose electrons
have a positive charge

Atoms that gain electrons
have a negative charge
BOINK!
BOINK!
CHLORINE

For the most part, the metals
will lose electrons and the
nonmetals will accept the
electrons

The atoms gain or lose
electrons to reach outer shell
(valence) stability
Electron from magnesium
Ionic Bonds: One Big Greedy Thief Dog!
Ion Sizes
+
Li,152 pm
3e and 3p
Li + , 60 pm
2e and 3 p
Does the size go
up or down
when losing an
electron to form
a cation?
Ion Sizes
+
Li,152 pm
3e and 3p


Li + , 78 pm
2e and 3 p
Forming
a cation.
CATIONS are SMALLER than the atoms from
which they come.
The electron/proton attraction has gone UP and
so size DECREASES.
Ion Sizes
Does the size go up or down when
gaining an electron to form an
anion?
Ion Sizes
F, 71 pm
9e and 9p


F- , 133 pm
10 e and 9 p
Forming
an anion.
ANIONS are LARGER than the atoms from which they
come.
The electron/proton attraction has gone DOWN and so
size INCREASES.
Trends in Ion Sizes
Figure 8.13
IONIZATION ENERGY
 the energy required to remove the most
loosely held electron from an atom
 ionization energy decreases as the size of
the atom increases
(top to bottom of a column)
“Y”? Because the outer most electron is
farther from the nucleus and the electrical
attraction to the protons.
More Details!
Energy
is absorbed by the atom to free the
electron(s)
Ionization is endothermic, meaning that the atom or
molecule increases its internal energy ( takes energy
from an outside source)
A + energy
A+ + e-
Ionization Energy is affected by
three factors:
1. Effective Nuclear Charge
2. Number of Energy Levels
3. Shielding
Ionization Energies

The first ionization energy, I1, is the energy needed to
remove the first electron from the atom:

Mg  Mg+ + 1e-

The second ionization energy, I2, is the energy needed to
remove the next (i.e. the second) electron from the atom

Mg+  Mg2+ + 1e-
•The higher the value of the
ionization energy, the more
difficult it is to remove the
electron
Ionization Energies in kJ/mol
1st IE
2nd IE
3rd IE
4th IE
5th IE
6th IE
7th IE
Na
496
4,560
Mg
738
1,450
7,730
Al
577
1,816
2,881
11,600
Si
786
1,577
3,228
4,354
16,100
P
1,060
1,890
2,905
4,950
6,270
21,200
S
999.6
2,260
3,375
4,565
6,950
8,490
27,107
Cl
1,256
2,295
3,850
5,160
6,560
9,360
11,000
Ar
1,520
2,665
3,945
5,770
7,230
8,780
12,000
• Within each period ( row) the ionization
energy increases with atomic number.
•Y?
-Electrons are being added to the same energy
level (ENC)
- increasing valence electrons as approaching
the nonmetals
Na
Mg
Al
Si
P
S
Cl
Ar
The Trend
Electronegativity

The tendency for an atom to attract electrons to itself when
in combination with another atom

Defined differences in electronegativity determine the
bonding character of a compound
• Ionic or Covalent bonds
Linus Pauling scale is used to determine electronegativity
differences
COVALENT BOND
 bond formed by the
sharing of electron clouds
• Between nonmetallic elements of similar
electronegativity.
•Formed by sharing electron pairs
NONPOLAR
COVALENT BONDS
when electron clouds are shared
equally
H2 or Cl2
2. Covalent bonds-
Two atoms share one or more pairs of outer-shell
electrons.
Oxygen Atom
Oxygen Atom
Oxygen Molecule (O2)
POLAR COVALENT BONDS
•
when electron clouds are shared but
shared unequally
H2O
Polar Covalent Bonds: Unevenly
matched, but willing to share.
- water is a polar molecule because oxygen is more
electronegative than hydrogen, and therefore electrons
are pulled closer to oxygen.
Electronegativity Differences and Bond Type
0
• If the electronegativity difference is less
than 0.2 then the bond is a nonpolar
covalent
nonpolar
covalent
0.2
polar
covalent
1.6
?
• If the difference is between 0.2 and 1.6,
the bond is polar covalent
•If the difference is greater than 2, the
bond is ionic
2
ionic
4
?  between 1.6 and 2, if
a metal is involved, the
bond is ionic. If only
nonmetals are involved the
bond is polar covalent
Trend of EN
decrease
increase
Electron Affinity
elements GAIN electrons to form anions.
Electron affinity is the energy change when an
electron is added:
A(g) + e- ---> A-(g)
E.A. = ∆E
Electron Affinity of Oxygen
O atom [He] 
 

+ electron
O- ion [He] 
 
EA = - 141 kJ

∆E is EXOthermic
because O has an
affinity for an e-.
Trends in Electron Affinity


Affinity for electron increases
across a period (EA becomes
more negative).
Affinity decreases down a group
(EA becomes less negative).
Atom EA
F
-328 kJ
Cl -349 kJ
Br -325 kJ
I
-295 kJ
Trends in Electron Affinity
Practice with Comparing
Ionization Energies
For each of the following sets of atoms, decide
which has the highest and lowest ionization
energies and why.
a. Mg, Si, S
b. Mg, Ca, Ba
c. F, Cl, Br
d. Ba, Cu, Ne
e. Si, P, N
Answers to Comparing Ionization Energies
Here are answers to the exercises above.
a. Mg, Si, S
All are in the same period and use the same number of
energy levels. Mg has the lowest I.E. because it has the
lowest effective nuclear charge. S has the highest I.E.
because it has the highest effective nuclear charge.
b. Mg, Ca, Ba
All are in the same group and have the same effective
nuclear charge. Mg has the highest I.E. because it uses
the smallest number of energy levels. Ba has the lowest
I.E. because it uses the largest number of energy
levels.
c. F, Cl, Br
All are in the same group and have the same
effective nuclear charge. F has the highest I.E.
because it uses the smallest number of energy
levels. Br has the lowest I.E. because it uses
the largest number of energy levels.
d. Ba, Cu, Ne
All are in different groups and periods, so both
factors must be considered. Fortunately both
factors reinforce one another. Ba has the
lowest I.E. because it has the lowest effective
nuclear charge and uses the highest number of
energy levels. Ne has the highest I.E. because it
has the highest effective nuclear charge and
uses the lowest number of energy levels.
e. Si, P, N
Si has the lowest I.E. because it has
the lowest effective nuclear charge
and is tied (with P) for using the
most energy levels. N has the
highest I.E. because it uses the
fewest energy levels and is tied
(with P) for having the highest
effective nuclear charge.
BECAUSE...
The relative stability of an atom can be
predicted by its electron configuration
Rule of Thumb
•
As a general rule,
elements with three or fewer
electrons in their outer level
are considered to be metals.
Lets Review!
1. What is the periodic Law?
2. How is an element’s outer electron configuration related to its
position in the periodic table?
3 Indicate which element in each of the following pairs has the greater
atomic radius.
a. sodium & lithium
b. strontium & magnesium
c. carbon & germanium
d. selenium & oxygen
4. In general, would you expect metals or nonmetals to have higher
ionization energies?
Test
Friday
More review!
5. Arrange the following elements in order of increasing ionization
energies.
a. Be, Mg, Sr
b. Bi, Cs, Ba
c. Na, Al, S
6. How does the ionic radius of a typical metallic atom compare to its
atomic radius?
7. Explain why it takes more energy to remove a 4s electron from an
atom of zinc than from than from an atom of calcium.
8. Give the symbol of the element found at each of the following
locations in the periodic table.
a. group 1, period 4
b. group 13, period 3
c. group 2, period 6
d. group 10, period 2
Even more review!
9. What was Newland’s Law of Octaves all about?
10. How was Mendeleev’s periodic table of elements better than the
previous attempts by others?
11. What property do the noble gases share? How does this property
relate to the electron configuration of the noble gases?
12. How do the electron configurations of the transition metals differ
form the electron configurations of the metals in groups 1 and 2?
13. What group numbers make up the main-block elements?
This test will be a bear if you forget to study!
Are you kidding me, more good stuff!
14. Define what ionization energy and electron affinity are.
15. What periodic trends exist for ionization energy? How about for
electron affinity? What about atomic radius and its trend?
16. Why does the first period of the periodic table contain only two
elements while all the other periods have eight or more element in
them?
17 What feature of electron configuration is unique to actinoids and
lanthanoids?
That should just about cover it!
TO REINFORCE YOUR ALREADY EXTENSIVE
KNOWLEDGE OF THE PERIODIC TABLE, YOU
CAN READ THROUGH THE PAGES IN YOUR
BOOK OF CHAPTER 14.
Lastly, if you need to
any last minute problems
you can show up at 7:00 in
room 224 for some last
minute brushing up.
out
Transition cont.
IT’S ALL ABOUT ELECTRONS!
Transition
elements fill
their sublevels differently
than do the Main group
elements.
For
the most part, there
are a few exceptions,
these “d” block metals will
place the 2 electrons into
a higher s-sublevel before
the electrons go into a
“d” energy sublevel.
Inner Transition Metals
the “ f ”-group

The ‘f’-group is broken into two classifications
~lanthanides
~actinides