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Transcript
Ch. 5 “The Periodic Table”
Why is the Periodic Table
important to me?
• The periodic table is
the most useful tool
to a chemist.
• It organizes lots of
information about all
the known elements.
• You get to use it on
every test.
Pre-Periodic Table Chemistry …
• …was a mess!!!
• No organization of
elements.
• Imagine going to a
grocery store with no
organization!!
• Difficult to find what you
need.
• Chemistry didn’t make
sense.
(CHAOS)
During the nineteenth century,
chemists began to categorize the
elements according to similarities
in their physical and chemical
properties. The end result of
these studies was our modern
periodic table.
How did chemists begin to
organize the known elements?
As the number of elements increased,
chemists inevitably began to find
patterns in their properties.
Chemists used the properties
of elements to sort them
into groups
Ex. Chlorine, bromine, and
iodine have very similar
chemical properties.
Model of Triads
In 1817, Johann Dobereiner classified some
elements into groups of three, which he called
triads.
The elements in a triad had similar chemical
and physical properties.
1780 - 1849
Law of Octaves
In 1865, John Newlands suggested that
elements be arranged in “octaves”
because he noticed when he arranged the
elements in order of increasing atomic
mass certain properties repeated every
8th element.
1838 - 1898
Lightest to heaviest.
A Li
B Be
C B
D C
E N
F O
G F
A Na
B Mg
C Al
D Si
E P
S
F
G Cl
A K
B Ca
C ?
D ?
E As
F Se
G Br
Li
Be
B
C
N
O
F
Na
Mg
Al
Si
P
S
Cl
K
Ca
?
?
As
Se
Br
He called this the
“Law of Octaves”
because of its
similarity to
musical octaves
John Newlands
Law of Octaves
Newlands' claim to see a repeating pattern was met
with savage ridicule on its announcement. His
classification of the elements, he was told, was as
arbitrary as putting them in alphabetical order and
his paper was rejected for publication by the
Chemical Society.
The Modern Periodic Table
In 1869 Dmitri Mendeleev
published a table of the
elements organized by
increasing atomic mass.
He was trying to organize
elements so his students
could learn them more
easily!
1834 - 1907
A. Mendeleev and Chemical
Periodicity
• Mendeleev placed known information of
elements on cards (atomic mass,
density, etc…). He arranged them in
order of increasing atomic masses,
certain similarities in their chemical
properties appeared at regular
intervals. Such a repeating pattern is
referred to as periodic.
At the same time, Lothar Meyer
published his own table of the elements
organized by increasing atomic mass.
1830 - 1895
• Both Mendeleev and Meyer arranged
the elements in order of increasing
atomic mass.
• Both left vacant spaces where unknown
elements should fit.
So why is Mendeleev called the “Father
of the Periodic Table” and not Meyer, or
both?
Could it be his dashing good looks?!
Mendeleev published first!
• Mendeleev left blank
spaces in his table when
the properties of the
elements above and
below did not seem to
match. The existence of
unknown elements was
predicted by Mendeleev
on the basis of the blank
spaces. When the
unknown elements were
discovered, it was found
that Mendeleev had
closely predicted the
properties of the
elements as well as
their discovery.
and the elusive element 32…
Predicted
Properties
Observed
Properties
Atomic weight
72
72.61
Density
5.5 g/cm3
5.32 g/cm3
Melting point
825 C
938 C
Oxide formula
RO2
GeO2
Density of oxide
4.7 g/cm3
4.70 g/cm3
Dates (predicted and
found)
1871
1886
Color
Dark gray
Gray-white
Gallium
Germanium
The Father of the
Periodic Table
After the discovery of these unknown
elements between 1874 and 1885, and the
fact that Mendeleev’s predictions were
amazingly close to the actual values, his
table was generally accepted.
However, in spite of Mendeleev’s great
achievement, problems arose when new
elements were discovered and more accurate
atomic weights were determined.
By looking at our modern periodic table, can
you identify what problems might have caused
chemists a headache?
Ar and K
Co and Ni
Te and I
Th and Pa
Remember This…?!
In 1913, through his work with X-rays,
Henry Moseley 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.”
1887 - 1915
Henry Moseley
His research was halted when the British
government sent him to serve as a foot
soldier in WWI.
He was killed in the
fighting in Gallipoli by a sniper’s bullet, at
the age of 28. Because of this loss, the
British government later restricted its
scientists to noncombatant duties during
WWII.
Periodic Law
When elements are arranged in order of
increasing atomic number, there is a
periodic repetition of their physical and
chemical properties.
The Current Periodic Table
• Mendeleev wasn’t too far off.
• Now the elements are put in rows by
increasing
ATOMIC NUMBER!!
• The vertical columns are called groups or
families and are labeled from 1 to 18 (modern)
• or in A & B Groups (with Roman numerals)
Groups…Here’s Where the
Periodic Table Gets Useful!!
• Elements in the
same group
have similar
chemical and
physical
properties!!
• (Mendeleev did that on
purpose.)
Why??
• They have the
same number of
valence
electrons.
• They will form
the same kinds
of ions.
Groups in the Periodic Table
Elements in groups react in similar ways!
Periodsrows
in the
Periodic
Table
The horizontal
are called
periods
and
are labeled from 1 to 7.
All elements in a period have the same number
of energy levels (= to period #)
11
1
2
3
4
5
6
7
Energy Levels
n=1
n=2
n=3
n=4
In addition to Group Labels,
many of the groups have Family Names
Group 1A: Alkali Metals
lithium
Cutting sodium metal
potassium
Alkali Metals
• They are the most
reactive metals.
• They react violently
with water.
• Alkali metals are never
found as free elements
in nature - they are
always in compounds
with other elements.
• Only 1 valence electron
• Soft metals
• Must be stored under
mineral oil, etc.
Group 2A: Alkaline Earth Metals
Magnesium
Group 2A: Alkaline Earth Metals
Only 2 valence electrons
Too reactive to be uncombined in nature.
calcium
strontium
barium
Group 7A: The Halogens
7 valence electrons
All non-metals
Very reactive
All physical states represented
Colored gases (always poisonous!)
Occur as diatomic molecules when pure
fluorine
F2
Cl2
chlorine
Br2
bromine
Iodine
I2
The Noble Gases
Noble Gases
• Noble Gases are colorless gases that are
extremely un-reactive.(inert)
• They are inactive because their outermost
energy level is full. (8 valence electrons –
except He which has 2)
• Having 8 valence electrons is low in energy
and, therefore, very stable.
Hydrogen
• The hydrogen square sits atop Family
IA, but it is not a member of that
family. Hydrogen is in a class of its
own. (An orphan?)
• Like the Alkali metals, it only needs to
lose one electron to be stable. (but it is
not a metal!)
• Sometimes it’s shown above 7A.
• Like the Halogens, it only needs to gain
one electron to have the stable Noble
Gas electron configuration. (but it is
not a Halogen!)
Hey Cameron, why are those elements
by themselves on the bottom of the
Periodic Table?!
I’ll handle this one, Cam!
If they weren’t put on
the bottom, the Periodic
Table wouldn’t fit very
nicely on a page!
In fact, the table would
look like this.
In fact, we have Glen
Seaborg to thank for
the fact that my
Periodic Table doesn’t
stick out of my
notebook in a truly
tasteless manner!
Glenn T. Seaborg
After co-discovering 10 new elements, in
1944 he moved 14 elements out of the
main body of the periodic table to their
current location below the Lanthanide
series. These became known
as the Actinide series.
1912 - 1999
“I was warned at the time that it was professional
suicide to promote this idea, which has since been
called one of the most significant changes in the
periodic table since Mendeleev’s 19th century
design. Luckily, I stuck to my guns and have seen
the actinide concept become the foundation for
many significant discoveries in heavy element
research.”
Seaborgium
Glenn T. Seaborg
He is the only person to have an element
named after him while still alive.
"This is the greatest honor ever bestowed
upon me - even better, I think, than
winning the Nobel Prize."
106
Sg
1912 - 1999
Seaborgium
271
There are many ways that
we can break the Periodic
Table up into sections!
Metals
Metals
Metals
Metals are good conductors of heat. That's
why a branding iron is made from metal.
The heat transfers quickly to the animal's hide.
Metals also conduct electricity.
Notice that the Tesla coil sparks seek out metallic
objects because they conduct electricity better than the
nonmetallic materials such as wood or soil.
Metals are also malleable and can be bent or hammered
into various shapes.
Metals
•
•
•
•
What comes to mind?
Most elements are metals
Loosely held valence e-’s
Properties of metals:
1. Good conductors of heat
and electricity (p)
2. High density (p)
3. High melting points (p)
4. Luster (p)
5. Malleable (p)
6. Ductile (p)
7. 1, 2, or 3 valence electrons
Nonmetals
Nonmetals do
not conduct
heat well. The
insulating
tiles from the
Space
Shuttle are
made
from fibers of
silicon
and oxygen
(silica=sand).
Nonmetals
•
•
Opposite of metals
Properties of nonmetals:
1. Dull (no luster)
2. Do not conduct heat/elec.
3. Not ductile
4. Not malleable
5. All phases
6. Have 5, 6, or 7 valence
electrons
• Form many compounds with
metals
Metalloids
(Semi-Metals)
• Means “metal-like”
• Dividing line between
metals and nonmetals
• Al is the exception
• Properties of both
metals and nonmetals
Four Main Categories of Elements
• Noble Gases- group 18 or 0 or 8A
– s & p sublevels filled
– 8 valence __s2…__p6
– Inert- not reactive- because of full outer
shell of electrons
• Representative Elements also called
main group elements- Groups 1A-7A
– s & p partly filled
– Includes alkali metals, alkaline earth
metals, and halogens
1A
 The
2A
elements in the A groups 8A
0
are called the representative
3A 4A 5A 6A 7A
elements
Four Main Categories of Elements
• Transition Metals –
– Unfilled inner shells
– outermost s & inner d sublevels contain
electrons
– Hard & brittle
• Inner transition metals– outermost s & nearby f sublevel contain
electrons
– Lanthanides (4f) and actinides (5f)
Representative
Representative
Noble Gases
Inner Transition Elements
Using the Diagonal Rule is just
so bothersome! I wish there
was an easier way to figure
out electron configurations!
Oh, but there
is! Watch this!
1
2
H
1 1s1
Li 1s22s1
3
22s22p63s1
1s
Na
3
11
K
4
19 1s22s22p63s23p64s1
Rb
22s22p63s23p63d104s24p65s1
5
1s
37
Cs 1s22s22p63s23p63d104s24p64d105s2 5p66s1
6 55
22s22p63s23p63d104s24p64d104f145s25p65d1
1s
Fr
06s26p67s1
7 87
Group 2A
Be
1s22s2
Mg
1s22s22p63s2
Ca
1s22s22p63s23p64s2
Sr
1s22s22p63s23p64s23d104p65s2
s-block
1
H
2
He
2
3
Li
4
Be
3
11
Na
12
Mg
Always for row you are on! 4
19
K
20
Ca
5
37
Rb
38
Sr
6
55
Cs
56
Ba
7
87
Fr
88
Ra
Group 3A
B
Al
Ga
1s22s22p1
1s22s22p63s23p1
1s22s22p63s23p64s23d104p1
Always for row you are on!
The P-block
p 1 p2
p3
p4
p5
p6
6.2
Electron Configurations in Groups
• In atoms of the Group 1A elements below,
there is only one electron in the highest
occupied energy level.
It’s always s1 for
the row it’s on!
• In atoms of the Group 4A elements below,
there are four electrons in the highest occupied
energy level. Always s2p2 for the row they’re on!
6.2
Electron Configurations in
Groups
– The Noble Gases
• In atoms of the Group 8A elements below, there
are eight electrons in the highest occupied energy
level.
Except for He, always s2p6 for the row they are on!
Chemical elements in d-block
Always 1 level in (back) from the
row you’re on!
Group
→
3
4
5
6
7
8
9
10
11
12
4
21
Sc
22
Ti
23
V
24
Cr
25
Mn
26
Fe
27
Co
28
Ni
29
Cu
30
Zn
3d
5
39
Y
40
Zr
41
Nb
42
Mo
43
Tc
44
Ru
45
Rh
46
Pd
47
Ag
48
Cd
4d
6
71
Lu
72
Hf
73
Ta
74
W
75
Re
76
Os
77
Ir
78
Pt
79
Au
80
Hg
5d
7
103 104 105 106 107 108 109 110 111 112
Lr
Rf Db Sg Bh Hs Mt Ds Rg Uub
6d
↓ Period
F - block
 inner
transition elements
Always go back 2 energy levels
from the row you’re on!
f1 f2 f3 f4 f5 f6 f7 f8 f9 f10 f11 f12 f13 f14
4f
5f
Blocks of Elements
s-block
d-block
p-block
f-block
Periodic Table
e- configuration from the periodic
periodic table
(To be covered in future chapters)
1
IA
18
VIIIA
2
IIA
1
H
1s1
2
Li Be
2s1 2s2
Na Mg
3s1 3s2
3
4
5
6
7
13
IIIA
3
IIIB
4
IVB
Sc
3d1
Rb
5s1
Ca
4s2
Sr
5s2
Y
4d1
V
Ti
Cr Mn Fe Co
3d2 3d3 4s13d5 3d5 3d6 3d7
Zr Nb Mo Tc Ru Rh
4d2 4d3 5s14d5 4d5 4d6 4d7
Cs
6s1
Ba
6s2
La
5d1
Hf Ta W Re Os
5d2 5d3 6s15d5 5d5 5d6
Fr
7s1
Ra
7s2
Ac Rf
6d1 6d2
K
4s1
5
VB
6
VIB
7
VIIB
Db Sg Bh
6d3 7s16d5 6d5
8
9
VIIIB
14
IVA
15
VA
16
VIA
17
VIIA
B
2p1
•B
C
N
O
1
2
3
•2p
2p 2p 2p4
F
2p5
Ne
2p6
Al
3p1
Si
3p2
Cl
3p5
Ar
3p6
He
1s2
10
11
IB
12
IIB
Ni
3d8
Cu
4s13d10
Ni
4d8
5s14d10
Zn Ga Ge
3d10 4p1 4p2
Cd
In Sn
10
4d
5p1 5p2
As Se Be
4p3 4p4 4p5
I
Sb Te
5p3 5p4 5p5
Kr
4p6
Xe
5p6
Hg
Tl Pb
5d10 6p1 6p2
Bi Po At
6p3 6p4 6p5
Rn
6p6
Ir
Ni
7
5d 5d8
Hs Mt
6d6 6d7
Ag
Au
6s15d10
S
P
3
3p 3p4
Periodicity
When one looks at the chemical properties of
elements, one notices a repeating pattern of
properties when the elements are in order of
increasing atomic number.
Explaining Periodic Trends
Why a property is higher/lower, bigger/smaller, etc.!
1. Nuclear charge- the number of protons
in the nucleus.
More protons = increased nuclear charge
so increased attraction between the
nucleus and electrons.
Think of the nucleus as a magnet – each extra
proton makes the magnet more powerful at
attracting electrons & holding them tight!
2. Shielding- lessens the attractive force of the
nucleus for the valence electrons
– caused by electrons in energy levels between
the nucleus and the valence electrons
Shielding increases as you go down a group
because there are more energy levels (more
core electrons).
Shielding stays the same as you move across a
period because the number of energy levels is
staying the same.
Which atom has more shielding? (A) K or Ca
(B) Na or K
Which atom is smaller? (A) N or P
(B) Li or K
Metallic Character
Atomic Radius
• Atomic Radius- half the distance
between the nuclei of two atoms of the
same element in a diatomic molecule
Atomic Radius
• Trend for atomic size– Down a group, size increases
• Occurs because # of energy levels
increases
*Makes a BIG difference in size!!
• shielding also increases.
– Across a period, size decreases
• # of protons increases (nuclear charge
increases), pulling electrons closer
shielding doesn’t change because
electrons are added to the same energy
Atomic Radius
Ionization Energy
• Ionization Energy- energy needed to remove
an electron from an atom.
• Outer shell electrons are easier to remove than
‘core’ electrons so it takes less energy to
remove them!
Highest toward upper right corner of PT since
these atoms are smaller & their valence
electrons are closer to the nucleus
-so held more tightly
Trends in Ionization Energy
Periodic Trends
• Ionization energy
– Down a group- decreases
– because electrons are held more loosely
due to increased # of energy levels &
increased shielding
– Across a period- increases because
electrons are held more tightly due to
increased nuclear charge (increased # of
protons in the nucleus)
Ionization Energy
Ionization Energy
• There are big jumps in ionization
energy whenever you try to remove
an electron from an inner energy
level!
Ionization Energy
Electron Affinity
• Electron affinity of an element is the
energy given off when an atom (in the gas
phase) gains an electron to form an ion
– Example: F(g) + e-  F-(g)
– Ho (ENERGY) = -328.0 kJ/mol
Trends in Electron Affinity
• It decreases down a group, because
electron shielding blocks some of the
attraction from the nucleus
• It increases across a period, because
nuclear charge increases, attracting
electrons more strongly.
Periodic Trends
• Electronegativity- tendency for the
atoms of the element to attract
electrons when the atoms are part of a
compound
• Noble gases- no electronegativity valuesdon’t form compounds
• In general, metals have low EN and
nonmetals have high EN.
• The actual amount of EN an atom has is
indicated by a number on the Pauling
Electronegativity Scale that goes from 0 to 4.
• Dr. Linus Pauling set up this scale and gave
the element having the greatest EN an arbitrary
number of 4, and he assigned numbers to the
others relative to this element.
• Flourine is the most electronegative
element at 4. (3.98) and Francium is the least
electronegative at 0.7.
Periodic Trends
• Electronegativity Trends– Down a group – decreases- electron shielding
results in less attraction for electrons by the
nucleus
– Across a period- increases- higher atomic
number and consistent electron shielding
result in more attraction for electrons
• Electronegativity allows you to predict
bond type: covalent- polar vs. nonpolar
and ionic
General Trends in the Periodic Table:
Atomic & Ionic Radii
Ionization Energy (IE)
Electron Affinity (EA)
Electronegativity (c)
IE, EA, and c are useful concepts used to characterize
different types of bonding and estimate bond energies.
6.3
Summary of Trends
Increases
Decreases
Ionization
Ionicof
size
energy
Size
Size
Electronegativity
Atomic
Nuclear
Shielding
of
anions
cations
Size
Charge
Increases
DecreasesConstant
IONS
• Remember – Atoms are neutral
• But…atoms can gain or lose electrons
(*# of protons NEVER changes during
reactions!)
• IONS are atoms or groups of atoms with a
charge.
•
To tell the difference between an atom and an
ion, look to see if there is a charge in the
superscript!
• Examples:
Na Ca I O
•
Na+ Ca+2 I- O-2
•When an atom loses an electron it gets a
positive charge (because it now has more
protons than electrons)
Mg --> Mg+2 + 2 eWhen an atom gains an electron it forms a
negative ion (because it now has more
electrons than protons)
F + e- --> F-
Atom versus Ion
Forming Cations & Anions
A CATION forms
when an atom
loses one or
more electrons.
Mg --> Mg2+ + 2 eNow has 12 protons
& 10 electrons
An ANION forms
when an atom
gains one or
more electrons
F + e- --> FNow has 9 protons
& 10 electrons
– Metals have 1, 2, or 3 valence electrons
so tend to lose electrons (to get an
octet)- forming cations. (+ charge)
– Non-Metals have 5, 6, or 7 valence
electrons so tend to gain electrons (to
get an octet) - forming anions. (charge)
Periodic Trends
• Ionic Radii Trends
– Cations- smaller than neutral atom because
fewer electrons result in greater attraction by
nuclei
– Anions- larger than neutral atom because
more electrons result in less attraction by
nuclei
– Within period- size decreases
– Down a group – size increases
Forming cations
Forming anions
EXAMPLE
What would the charge be on a sodium ion?
Since sodium in in Group IA it has 1 valence eand so it would LOSE an electron
It goes from 11 protons & 11 electrons to
11 protons & 10 electrons
So it gets a charge of +1
Remember an electron is negatively charged.
When an atom loses electrons it forms
positively charged ions.
When electrons are gained negatively charged ions
form
EXAMPLE
How would you write the symbol for the sodium CATION?
Na
+1
How many outer electrons does sodium have before it
loses one?
It has 1…remember the group number!
5
6.3 Section Quiz
– 1.
Which of the following sequences is
correct for atomic size?
•
•
•
•
Mg > Al > S
Li > Na > K
F>N>B
F > Cl > Br
6.3 Section Quiz
– 2.
•
•
•
•
Metals tend to
gain electrons to form cations.
gain electrons to form anions.
lose electrons to form anions.
lose electrons to form cations.
6.3 Section Quiz
– 3.
Which of the following is the most
electronegative?
•
•
•
•
Cl
Se
Na
I
The Periodic Table
Summary of Trend
• Periodic Table and Periodic Trends
• 1. Electron Configuration
3. Ionization Energy: Largest toward NE of PT
4. Electron Affinity: Most favorable NE of PT
2. Atomic Radius: Largest toward SW corner of PT
Periodic Table: electron
behavior
West (South)
Mid-plains
East (North)
METALS
Alkali
Alkaline
Transition
These elements
tend to give up
e - and form
CATIONS
METALLOID
These elements
will give up e- or
accept e-
NON-METALS
Noble gas
Halogens
Calcogens
These elements
tend to accept
e - and form
ANIONS
• The periodic table can be classified by the behavior of their electrons
1
IA
1
2
IIA
13
IIIA
2
3
4
5
6
7
3
IIIB
4
IVB
5
VB
6
VIB
7
VIIB
8
9
VIIIB
10
11
IB
12
IIB
14
IVA
15
VA
16
VIA
17
VIIA
18
VIIIA
ELEMENTS THAT EXIST
AS DIATOMIC MOLECULES
Remember:
BrINClHOF
These elements
only exist as
PAIRS. Note that
when they
combine to make
compounds, they
are no longer
elements so they
are no longer in
pairs!
Valence electrons
1
2
3
4
5
6
7
8
Select an element
(
= Internet link )
Configuration of Ions
Ions of representative elements have noble gas
configuration
Na is 1s22s22p63s1
Forms a 1+ ion - 1s22s22p6
Same configuration as neon
Metals form ions with the configuration of the noble
gas before them - they lose electrons
This can explain why metals are shiny. This is the surface of copper
at a ridiculously high magnification. The surface
shows a lake of electrons along with ripples.
The two islands are imperfections on the surface.
Most likely a couple of atoms that aren't copper.