Download Ionization energy

Chapter 6:
The Periodic Table
General Chemistry
http://www.ccsdualsnap.com/miscellan.htm
Objectives
• Periodicity of physical and chemical
properties relates to atomic structure and led
to the development of the periodic table.
• The periodic table displays the elements in
order of increasing atomic number.
• Explain the relationship of an element’s
position on the periodic table to its atomic
number and mass.
Objectives
• Use the periodic table to identify metals,
nonmetals, metalloids, families (groups),
periods, valence electrons, and reactivity
with other elements in the table.
• Relate the position of an element on the
periodic table to its electron
configuration.
• Identify trends on the periodic table
(ionization energy, electronegativity,
electron affinity, and relative size of
atoms and ions).
Review/Link to Previous
Learning
• In Chapter 4, we learned about electrons
configurations of elements.
• Discovered there is a pattern of electron
configurations on the Periodic Table.
• Are there other patterns on the Periodic
Table? (yes)
• In Chapter 5 we will learn how the Periodic
Table is organized.
Collections
Do you like to play cards?
Do you have a stamp, baseball
card, or comic book collection?
How do you organize your collection?
Attempts at Organizing Elements
• Early scientists knew about some properties
of elements.
• Is there a characteristic of elements that can
organize them?
JOHAN DOBEREINER(1780-1849)
Dobereiner’s Triads
THE LAW OF TRIADS:
The atomic mass of the middle element of the triad
is equal to the mean of the atomic masses of the
other two elements.
EXAMPLE:
Lithium
Sodium
Potassium
Atomic Mass of 7
Atomic Mass of 23
Atomic Mass of 39
According to Dobereiner’s Law, the atomic mass of sodium
Should equal the arithmetic mean of lithium and potassium.
(7+39)/2 = 23, which is the mass of sodium.
Problems with Dobereiner’s
Law of Triads.
1) All the elements known at that time could not be
arranged in triads.
2) The law did not work for very low or very high
massed elements such as F, Cl, and Br.
3) As techniques improved for measuring atomic
masses accurately, the law became obsolete.
Dobereiner’s research made chemists look at groups of
elements with similar chemical and physical properties.
JOHN A.R. NEWLANDS (1837-1898)
Newland’s Law of Octaves
When placed in increasing order of their atomic
masses, every eighth element showed similar physical
and chemical properties.
Li Be B C N
Na Mg Al Si P
K Ca
O
S
F
Cl
Problems with
Newland’s Law of Octaves
1) It was not valid for elements that had atomic
masses higher than Ca.
2) When more elements were discovered
(Noble gases) they could not be accommodated in his
table.
However, the modern periodic table does draw from the
concept of periods of eight.
DMITRI MENDELEEV (1834-1907)
Julius Lothar Meyer (1830-1895)
Mendeleev and Meyer
• Published nearly identical schemes for classifying
elements
• Arranged elements by increasing atomic mass
• Mendeleev generally given more credit
– Published first
– More successful at demonstrating value of table
– Predicted discovery of new elements, properties
of new elements
Properties of Some Elements Predicted By Mendeleev
Mendeleev’s Table: the first periodic table of the elements.
He arranged the table so that elements in the same column
have similar properties.
Problems with Mendeleev’s Table:
1) The positions of isotopes could not be
accommodated within the table.
2) In order to make the elements fit the requirements,
Mendeleev was forced to put an element of slightly
higher atomic weight ahead of one of slightly lower
atomic weight.
Henry Moseley (1887-1915)
• Developed concept of atomic number
– amount of positive charge in the nucleus
• Later determined that arranging periodic
table according to increasing atomic number
eliminated problems seen in Mendeleev’s
table
Why is it the “periodic” table?
• Periodic Law: when elements are
arranged in order of increasing
atomic number, their physical and
chemical properties show a
periodic pattern
Study Buddy Review
• Describe the contribution each person
below made to the development of the
periodic table:
–
–
–
–
–
Johan Dobereiner
John Newland
Dmitri Mendeleev
Julius Meyer
Henry Moseley
Parts of the Periodic Table
Parts of Periodic Table
• Groups/families: vertical columns
–
–
–
–
–
–
–
Alkali metals: 1A
Alkali earth metals: 2A
Boron, carbon families
chalcogens (oxygen family).
pnictogens (nitrogen family)
Halogens (fluorine family): 7A
Noble gasses: 8A/0
• Horizontal rows are called periods
• There are 7 periods
• The elements in the A groups are
8A
1A
0
called the representative elements
2A
3A 4A 5A 6A 7A
outer s or p filling
Parts of Periodic Table
• Metals: left of staircase
– Luster, malleable, conduct, ductile
• Nonmetals: right of staircase
– Dull in appearance, nonconductor, brittle
• Metalloids: elements adjacent to staircase
(except Al, Po)
– Some properties of both metals and nonmetals
The group B are called the
transition metals
 These
are called the inner
transition metals and they
belong here
Study Buddy Review
• Identify the follow parts of the periodic table:
–
–
–
–
–
–
–
–
Halogens
family
Alkali metals
Metals
Inner transition metals
Noble gases
Metalloids
Period
Periodic Properties of Elements
Periodic Trends
•Atomic Radius
•Ionic Radius
•Ionization Energy
•Electron Affinity
•Electronegativity
Atomic Radius
}
Radius
•Atomic Radius = half the distance between two
nuclei of a diatomic molecule.
Trends in Atomic Radius
• Influenced by three factors:
1. Charge on nucleus
– More charge pulls electrons in closer.
2. Energy Level
– Higher E level is further away from nucleus
3. Shielding effect
– The number of electrons between electrons and
nucleus affects the pull felt by the outer
electrons
Atomic Radius
Group trends
• As we go down a
group...
• each atom has
another energy
level
• so the atoms get
bigger.
H
Li
Na
K
Rb
Atomic Radius Periodic Trends
• As you go across a period, the radius gets
smaller.
• The increasing number of protons in the
nucleus pulls the electrons in more tightly
Na
Mg
Al
Si
P
S Cl Ar
Atomic Radius
0.250
Atomic Radii
Atomic radius, nm
0.200
0.150
0.100
0.050
Atomic Radii
0.000
3
4
5
6
7
8
9
10
11
12
13
atomic number, Z
14
15
16
17
18
19
20
Ionic Size
• Ion: electrically charged atom
• Cation: positively charged ion
• Anion: negatively charged ion
• Ions aren't the same size as the neutral
atoms they come from.
– Compare the sizes of sodium and chloride ions
with the sizes of sodium and chlorine atoms.
• Positive ions are
smaller than the
atoms they come
from.
• The sodium ion
loses a whole layer
of electrons, and the
remaining 10
electrons are being
pulled in by the full
force of 11 protons.
• Negative ions are
bigger than the atoms
they come from.
• Although the
electrons are still all in
the 3-level, the extra
repulsion produced by
the incoming electron
causes the atom to
expand. There are still
only 17 protons, but
they are now having to
hold 18 electrons.
Study Buddy Review-A.R., I.R.
• Describe the pattern for atomic radius
– As you move across a period
– As you move down a column
• What charge does a cation have?
• What charge does an anion have?
• Which is larger than its parent atom, a
cation or an anion?
First Ionization Energy
• Ionization energy is the energy required to remove
the first electron from an atom of an element
• Elements want to have the e- configuration like
that of a noble gas (filled)
– Column 1A elements have need to LOSE one electron
to have noble gas configuration so it is EASY to
remove electron
– Column 7A element need to GAIN one electron to have
noble gas configuration, so it is HARD to remove
electron
First Ionization Energy vs.
Atomic Number
Ionization Energy
Ionization Energy
• As you move down

a group ionization
energy decreases…
• The electrons that
are further away
from the nucleus
are easier to
remove and thus
require less energy
to remove
• As you move
across a period
ionization energy
increases…
• Elements on left of
table want to lose
electrons to have
full energy level
(requires low
energy to remove
electron)
Successive Ionization Energies
• more than one electron can be removed
from atoms
• Second Ionization energy: when a second
electron is removed from an atom that has
already lost one electron
• Third Ionization energy: when a third
electron is removed from an atom that has
already lost two electrons
Sucessive Ionization Energies
Symbol First
H
He
Li
Be
B
C
N
O
F
Ne
1312
2731
520
900
800
1086
1402
1314
1681
2080
Second
Third
5247
7297
1757
2430
2352
2857
3391
3375
3963
11810
14840
3569
4619
4577
5301
6045
6276
Relationship Between Common
Charge and I.E.
•
•
•
•
Consider Beryllium:
Electron config: [He] 2s2
Low energy to remove 1st and 2nd electrons
MUCH higher energy to remove 3rd
electron because it would be removed form
a noble gas configuration
Study Buddy Review-I.E.
• What is ionization energy?
• Describe the pattern for ionization energy as
you
– Move down a family
– Move across a row
• What does “first” ionization energy mean?
Electron Affinity
Electron affinity is:
• the energy change associated with the
adding an electron to a gaseous atom
• the more attraction for an electron the
energy is released when the atom gains the
electron
– Released energy is negative (-350 kJ)
Electron Affinity
Electron Affinity
General Trend:
• Halogens (s2p5 configurations) are most
negative electron affinities. They are most
likely to want to gain electrons to obtain
noble gas configuration
• As you go down a family, electron affinity
is less negative (harder to gain electrons
with increasing atomic size)
Study Buddy Review-E.A.
• What does it mean when an energy is
negative?
• Which elements generally have a very
negative electron affinity?
Electronegativity
Electronegativity:
the tendency for an atom
to attract electrons to itself when it is
chemically combined with another element.
• As you move down a
group, electronegativity
decreases
• As you move across a
period, electronegativity
increases
Which element is the MOST electronegative?
Electronegativity
Electronegativity
4
3.5
3
2.5
2
1.5
1
0.5
0
1
2
3
4
5
6
7
8
9
10
11
12
atom ic num ber, Z
13
14
15
16
17
18
19
20
Study Buddy Review-Electroneg
• Define electronegativity.
• Describe the pattern for electronegativity as
you
– Move down a group
– Move across a period
• Which element is the most electronegative?
Resources
• http://www.chemguide.co.uk/atoms/properties/atra
dius.html
• http://wine1.sb.fsu.edu/chm1045/notes/Periodic/A
ffinity/Period05.htm
• http://www.webelements.com/
• http://www.public.asu.edu/~jpbirk/CHM113_BLB/Chpt07/sld017.htm
• Jeanette Boles
• Tina Lula
• Dr. Stephen L. Cotton, Charles Page High School