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Transcript
1
About 70 elements had been
discovered by the mid-1800’s, but no
one had found a way to relate the
elements in a systematic, logical way.
2
He develop the 1st
periodic table of the
elements.
Arranged elements
in order of increasing
atomic mass and
created columns
with elements having
similar properties.
3
Mendeleev left
blank spaces in
the table
because there
were no known
elements with
the appropriate
properties and
masses.
4
Mendeleev and others were able
to predict the physical and
chemical properties of the missing
elements. Eventually these
elements were discovered and
were found to have properties
similar to those predicted. There
were many exceptions in his table,
however.
5
In 1913, arranged
elements in order
of increasing
atomic number
thus reversing the
order of the
elements and
correcting the
drawbacks found
in Mendeleev’s
table.
6
7
Neat-o Animation
8
Periodic law states the physical and
chemical properties of the elements
are periodic functions of their atomic
number. In other words, when the
elements are arranged in order of
atomic number, elements with similar
properties appear at regular intervals.
9
A Group or Family is a column on the
periodic table. Elements in the same
column have similar chemical
properties.
10
2 conventions for
numbering:
1-18
A/B elements
11
12
Main Group Elements
Transition Metals
Inner Transition
Metals
C. Johannesson
Group A elements all have electrons
in the outer s, or s and p orbitals.
These are known as representative
elements. The group number
indicates the number of valence, or
outer shell, electrons except with
helium which has 2.
Examples:
IIA - Ca (20) 1s22s22p63s23p64s2
VIA – S (16) 1s22s22p63s23p4
14
Group 18 (VIIIA)
elements are the
noble gases with
8 valence
electrons, except
helium which
has 2. Noble
gases are inert
(nonreactive) in
nature. They do
not form ions.
15
Full s & p orbitals in
the highest principal
energy level
Electron configuration
very stable, making
them inert
When other atoms of
other elements gain or
lose electrons in
reactions, they
achieve electron
configuration of noble
gases
16
Group B elements or transition
elements (d block) have electrons in
their outer d orbitals. The have
varying number of valence electrons.
Example: Zn (30)
1s22s22p63s23p64s23d10
17
Lanthanides and Actinides elements (f-block)
have electrons in their outer f orbitals.
These elements have varying numbers of
valence electrons.
Example: Nd (60)
2
2
6
2
6
2
10
6
2
1s 2s 2p 3s 3p 4s 3d 4p 5s
10
6
2
1
3
4d 5p 6s 5d 4f
18
19
Hydrogen
No group number
Only element in family
Most common element in the
universe
Very reactive
Compounds of H very commonH2O
Found in proteins, carbs, and fats
with C and O
20
 When elements are arranged in order of
increasing atomic #, elements with similar
properties appear at regular intervals.
Atomic Radius (pm)
250
200
150
100
50
0
0
5
10
Atomic Number
C. Johannesson
15
20
Trends on table occur vertically and
horizontally
Group 1
Li
increasing reactivity with H2O
Na
K
Knowing the trends enables you to
predict chemical behavior.
22
You know from the quantum
mechanical model that an atom does
not have a sharply defined boundary
that sets the limit of its size.
Therefore, the radius of an atom cannot
be measured directly. There are,
however, several ways to estimate the
relative size of atoms.
The atomic radius is one-half of the
distance from center to center of 2 like
atoms.
23
Group Trends
Atomic size generally increases as
you move down a group of the periodic
table.
Why?
Li 2s1 adding principal energy
levels
Na 3s1
K 4s1
Atoms getting larger with more energy
levels
Electrons getting further away from +
charged nucleus
24
shielding effect- the reduction of the
attractive forces between a nucleus and
its outer electrons due to the blocking
effect of inner electrons
• The shielding of the nucleus by
electrons also increases with the
additional occupied orbitals between
the outermost orbital and the nucleus.
25
Periodic Trends
Atomic size generally decreases as you move
from left to right across a period.
Why?
As you go across a period, the principal
energy level remains the same. Each element
has one more proton and one more electron
than the preceding element.
The electrons are added to the same energy
level, causing the increasing positive charge
to pull them in closer.
Realize at some point this effect is less
pronounced more electrons, more reaction
between them repulsion and that force is
greater than positive attraction of nucleus
26
27
When an atom gains or loses an electron, it
becomes an ion.
The energy required to overcome the attraction of
the nuclear charge and remove an electron from a
atom is called the ionization energy.
Removing one electron results in the formation of a
positive ion with a 1+ charge.
Na(g)
Na+(g) + e-
28
The energy required to remove this first
outermost electron is called the first
ionization energy.
To remove the outermost electron from
the 1+ ion requires an amount of
energy called the second ionization
energy, and so forth.
Nifty swell animation
29
Group Trends
Ionization energy generally decreases
as you move down a group of the periodic
table.
This is because the size of the atoms
increases as you descend, so the outermost
electron is farther from the nucleus.
The outermost electron should be
more easily removed, and the element
should have a lower ionization energy.
30
Periodic Trends
For the representative elements, ionization energy
generally increases as you move from left to right
across a period.
The atomic number and therefore positive charge
increases and the shielding effect is constant as
you move across. A greater attraction of the
nucleus for the electron leads to the increase in
ionization energy.
Also, electron configuration/ noble gas
configuration harder to remove electron to get to
a more stable electron configuration- easier to
gain an electron
31
32
The electronegativity of an element is the
tendency for the atoms of the element to
attract electrons when they are chemically
combined with atoms of another element.
Electronegativity generally decreases as you
move down a group.
As you go across a period from left to right,
the electronegativity of the representative
elements increases.
33
The metallic elements at the far left of the periodic
table have low electronegativities. By contrast, the
nonmetallic elements at the far right (excluding the
noble gases), have high electronegativities.
The electronegativity of cesium, a metal, the least
electronegative element, is 0.7; the electronegativity
of fluorine, a nonmetal, the most electronegative
element, is 4.0.
Because fluorine has such a strong tendency to
attract electrons, when it is chemically combined to
any other element it either attracts the shared
electrons or forms a negative ion.
In contrast, cesium has the least tendency to attract
electrons.
Groovy animation
34
35