Download The Periodic Table

The Periodic Table
Pre-Periodic Table Chemistry …
 …was a mess!!!
 No organization of
elements.
 Imagine going to a
grocery store with no
organization!!
 Difficult to find
information.
 Chemistry didn’t
make sense.
Dmitri Mendeleev (1834-1907)
• Presented the first valid periodic table of
elements to the Academy of Sciences in 1869
• Grouped elements in ascending order by
atomic mass and by similarities in properties
• left gaps for elements not yet discovered
• Element Mendelevium (element 101) was
named to honor him
Dmitri Mendeleev: Father of the
Periodic Table
HOW HIS WORKED…
 Put elements in rows
by increasing atomic
weight.
 Put elements in
columns by the way
they reacted.
SOME PROBLEMS…
 He left blank spaces
for what he said
were undiscovered
elements. (Turned
out he was right!)
 He broke the pattern
of increasing atomic
weight to keep
similar reacting
elements together.
The Current Periodic Table
• Mendeleev wasn’t too far off.
• Now elements are put in rows by increasing
ATOMIC NUMBER!!
• The horizontal rows are called periods and are
labeled from 1 to 7.
• The vertical columns called groups are
chemically similar and labeled from 1 to 18.
Across the Periodic Table
•
•
Periods: Are arranged horizontally across the periodic table
(rows 1-7)
•
These elements have the same number of valence shells (outer energy level).
1
IA
1
18
VIIIA
2
IIA
13
IIIA
2nd Period
2
3
3
IIIB
4
IVB
5
VB
6
VIB
7
VIIB
4
5
6
7
6th Period
8
9
VIIIB
10
11
IB
12
IIB
14
IVA
15
VA
16
VIA
17
VIIA
Quick quiz
• 1. Name the two properties that Mendeleev
used to arrange his Periodic Table?
• 2. When was Mendeleev working on his
Periodic Table?
• 3. How did the answer of #1 change as we
moved to the Modern Periodic Table?
Periodic Table
Electron configuration from the Periodic Table
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
Valence Electrons
• The valence electrons are the electrons
in the last shell or outer energy level of
an atom. They show a repeating or
periodic pattern. The valence electrons
increase in number as you go across a
period.
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 and bonds.
Families on the Periodic Table
 Columns are also
grouped into families.
 Families may be one
column, or several
columns put together.
 Families have names
rather than numbers.
(Just like your family
has a common last
name.)
Down the Periodic Table
•Family: Are arranged vertically down the periodic table
•(columns or group, 1- 18 or 1-8 A,B)
•These elements have the same number electrons in the outer most shells, the valence shell.
1
IA
1
18
VIIIA
Alkali Family:
1 e- in the valence shell
2
IIA
13
IIIA
14
IVA
15
VA
16
VIA
2
3
3
IIIB
4
IVB
5
VB
6
VIB
7
VIIB
8
9
VIIIB
10
11
IB
12
IIB
4
5
6
7
Halogen Family:
7 e- in the valence shell
17
VIIA
Infamous Families of the Periodic Table
• Notable families of the Periodic Table
Halogen
Noble Gas
Alkali
Alkaline
(earth)
1
IA
1
18
VIIIA
2
IIA
Transition Metals
2
3
4
5
6
7
3
IIIB
4
IVB
5
VB
6
VIB
7
VIIB
8
9
VIIIB
10
11
IB
12
IIB
13
IIIA
14
IVA
15
VA
16
VIA
17
VIIA
Metals vs Non-Metals
• Metals are found to the
left of the staircase
• Lose e-’s
• Good conductors of
heat and electricity
• Malleable and ductile
• Shiny & hard
• High MP/BP
• Non-metals are right of
the staircase
• Gain e-’s
• Poor conductors of heat
and electricity
• Brittle
• Not shiny & softer
• Low MP/BP
Periodic Table: Metallic arrangement
• Layout of the Periodic Table: Metals vs. nonmetals
1
IA
1
18
VIIIA
2
IIA
13
IIIA
14
IVA
15
VA
16
VIA
17
VIIA
2
3
4
5
6
7
3
IIIB
4
IVB
5
VB
6
VIB
7
VIIB
8
9
VIIIB
Metals
10
11
IB
12
IIB
Nonmetals
Reading the Periodic Table: Classification
•
Nonmetals, Metals, Metalloids, Noble gases
Hydrogen






Hydrogen belongs to a family of its
own.
Hydrogen is a diatomic, reactive
gas.
Hydrogen was involved in the
explosion of the Hindenburg in
1937.
Hydrogen is promising as an
alternative fuel source for the
future
1 valence ehttp://www.bing.com/videos/sear
ch?q=hindenburg+disaster&qpvt=
hindenburg+disaster&FORM=VDRE
#view=detail&mid=41F914E72CBA
1C33237741F914E72CBA1C33237
7
Alkali Metals
 1st column on the periodic
table (Group 1) not including
Hydrogen.
 Very reactive metals, always
combined with something else
in nature (like in salt, NaCl).
Never found as individual
atoms.
 Metals explode in water
http://www.sciencekids.co.nz/
videos/experiments/metalexpl
osion.html
 Soft enough to cut with a
butter knife. Low melting
point and boiling point
 1 valence e-
Quick Quiz
• 1. Why is Hydrogen not included in the Alkali Family?
– Although hydrogen has 1 valence electron, it is a non-metal
• 2. Are Alkali Metals known for their lack of chemical activity?
– No, alkali metals are highly reactive
• 3. Are Alkali Metals stored underwater so they cannot come in
contact with the Oxygen in the air and begin to oxidize?
– No, alkali metals explode in water!
• 4. If Hydrogen becomes a substitute for petroleum fuels, will cars
become much less safe than today’s vehicles?
– No, hydrogen fuel does not produce hot ash or radiant heat as
gasoline, and hydrogen leaks disperse rapidly in the atmosphere
resulting in less time to burn; Issues with tank material, strength, and
weight
Alkaline Earth Metals
 Second column on the
periodic table. (Group 2)
 Reactive metals that are
always combined with
nonmetals in nature.
 Harder, denser, and stronger
than alkali metals
 Several of these elements are
important mineral nutrients
(such as Mg in bones and Ca
in teeth and bones.
 2 valence e-’s
Transition Metals
 Elements in groups 3-12
 Less reactive harder metals
 Gold and platinum among
the least reactive of all
elements
 High luster; strong but
deformable
 Good electric and heat
conductors
 Includes metals used in
jewelry and construction
 Metals used “as metal”
 Similar properties due to
same number of electrons in
outer energy level
Boron Family
• Elements in group 13
• Aluminum metal was once rare
and expensive as Gold. Not a
“disposable metal.”
• Has now been discovered in a very
common ore called Bauxite
• Family has 3 valence e-’s
Rare Earth Elements
•
•
•
•
•
•
Lanthanide Series
Atomic Numbers 57-71
Not periodic
Unpredictable
Shiny metals
Similar reactivity to
Group 2 Alkaline Earth
Metals
• Actinide Series
• Atomic Numbers 89103
• Uranium is heaviest
natural element
• Heavier elements are
man-made ONLY
• All element above 92
are Radioactive
Carbon Family
• Elements in group 14
• Contains CARBON which is
important to all living things
• Carbon is the basis for Organic
Chemistry.
• Silicon and Germanium are
important in computers as
semiconductors.
• Form the most, strongest bonds
• 4 valence e-’s
Nitrogen Family
 Elements in group 15
 Nitrogen makes up 78% of the
atmosphere.
 Nitrogen and phosphorus are
both important in living things
(proteins and amino acids).
 Most of the world’s nitrogen is
not available to living things.
 Nitrogen is found in explosives
(TNT) and fertilizers
 The red stuff on match heads is
Phosphorus
 5 valence e-
Oxygen Family
 Elements in group 16
 Oxygen is necessary
for respiration and
combustion.
 Many things that
stink, contain sulfur
(rotten eggs, garlic,
skunks,etc.)
 6 valence e-
Halogens
 Elements in group 17
F, Cl, Br, I
 Very reactive, volatile,
diatomic, nonmetals
 Very TOXIC
 Always found
combined with other
elements in nature .
 Used to kill bacteria
and to strengthen
teeth.
 7 valence e-’s
Quick Quiz
• 1. What element is most often associated with explosives?
– Nitrogen
• 2. What element is frequently associated with bad smelling
products?
– Sulfur
• 3. If your toothpaste has Sodium Fluoride, what is it doing for
you?
– Kill bacteria and strengthen teeth
• 4. What properties of Carbon cause it to be the “Basis of
Life”?
– Four valence electrons for bonding; can form long chains
of stable molecules; single double and triple bonds
The Noble Gases
Helium, Neon, Argon, Krypton, Xenon,
Radon and Ununoctium
The Noble Gases
 Elements in group 18
 VERY UNREACTIVE
monatomic gases
 Used in lighted “neon”
signs
 Helium is used to fix
the Hindenburg
problem and balloons
 Have a full valence
shell with usually 8
valence e-’s.
Metalloids
•
•
•
•
B, Si, Ge, As, Sb, Te
Semiconductors
Mostly brittle solids
Some properties of metals and some of nonmetals
Summary
• Periodic Table: Map of the Building block of matter
• Type: Metal, Metalloid and Nonmetal
• Groupings: Representative or main, Transition and
Lanthanides/Actanides
• Family: Elements in the same column have similar
chemical property because of similar valence electrons
•
Alkali, Alkaline, Halogens, Noble gases
• Period: Elements in the same row have valence
electrons in the same shell (energy level)
ALL Periodic Table Trends
• Influenced by three factors:
1. Energy Level (shell)
–Higher energy levels are further away
from the nucleus.
2. Charge on nucleus (# protons)
–More charge pulls electrons in closer.
(+ and – attract each other)
• 3. Shielding effect (blocking effect)
What do they influence?
Energy levels and Shielding have
an effect on elements within a
GROUP
Nuclear charge has an effect on
elements within a PERIOD
Atomic Size
}
Radius
• Measure the Atomic Radius - this is half the distance
between the two nuclei of a diatomic molecule.
Atomic Radius - Group trends
• As we increase the
atomic number (or
go down a group). . .
• each atom has
another energy level,
• so the atoms get
H
Li
Na
K
bigger.
Rb
Atomic Radius - Period Trends
• Going from left to right across a period, the size
gets smaller.
• Electrons are in the same energy level.
• But, there is more nuclear charge.
• Outermost electrons are pulled closer.
Na
Mg
Al
Si
P
S Cl Ar
Trend in Atomic Radius
Ions
• Some compounds are composed of
particles called “ions”
– An ion is an atom (or group of atoms) that has a
positive or negative charge
• Atoms are neutral because the number of
protons equals electrons
– Positive and negative ions are formed when
electrons are transferred (lost or gained)
between atoms
Ions
• Metals tend to LOSE electrons, from
their outer energy level
–Sodium loses one: there are now more
protons (11) than electrons (10), and
thus a positively charged particle is
formed = “cation”
–The charge is written as a number
followed by a plus sign: Na1+
–Now named a “sodium ion”
Ions
• Nonmetals tend to GAIN one or more
electrons
–Chlorine will gain one electron
–Protons (17) no longer equals the
electrons (18), so a charge of -1
–Cl1- is re-named a “chloride ion”
–Negative ions are called “anions”
Trends in Ionic Size: Cations
• Cations form by losing electrons.
• Cations are smaller than the atom they
came from – not only do they lose
electrons, they lose an entire energy
level.
• Metals form cations.
• Cations of representative elements have
the noble gas configuration before them.
Ionic size: Anions
• Anions form by gaining electrons.
• Anions are bigger than the atom they
came from – have the same energy level,
but a greater area the nuclear charge
needs to cover
• Nonmetals form anions.
• Anions of representative elements have
the noble gas configuration after them.
Ion Group trends
• Each step down a
group is adding an
energy level
• Ions therefore get
bigger as you go
down, because of
the additional energy
level.
Li1+
Na1+
K1+
Rb1+
Cs1+
Ion Period Trends
• Across the period from left to right,
the nuclear charge increases - so they
get smaller.
• Notice the energy level changes
between anions and cations.
Li1+
B3+
Be2+
C4+
N3-
O2-
F1-
Configuration of Ions
• Ions always have noble gas configurations
( = a full outer level)
• Na atom is: 1s22s22p63s1
• Forms a 1+ sodium ion: 1s22s22p6
• Same configuration as neon.
• Metals form ions with the configuration of
the noble gas before them - they lose
electrons.
Configuration of Ions
• Non-metals form ions by gaining
electrons to achieve noble gas
configuration.
• They end up with the
configuration of the noble gas
after them.
Trends in Ionization Energy
• Ionization energy is the amount of
energy required to completely remove an
electron (from a gaseous atom).
• Removing one electron makes a 1+ ion.
• The energy required to remove only the
first electron is called the first ionization
energy.
Ionization Energy
• The second ionization energy is the
energy required to remove the second
electron.
–Always greater than first IE.
• The third IE is the energy required to
remove a third electron.
–Greater than 1st or 2nd IE.
Symbol First
H
He
Li
Be
B
C
N
O
F
Ne
1312
2731
520
900
800
1086
1402
1314
1681
2080
Second
5247
7297
1757
2430
2352
2857
3391
3375
3963
Third
Why did these values
increase so much?
11810
14840
3569
4619
4577
5301
6045
6276
What factors determine IE
• The greater the nuclear charge, the
greater IE.
• Greater distance from nucleus decreases
IE
• Filled and half-filled orbitals have lower
energy, so achieving them is easier, lower
IE.
• Shielding effect
Shielding
• The electron on the
outermost energy level
has to look through all
the other energy levels
to see the nucleus.
• Second electron has
same shielding, if it is in
the same period
Ionization Energy - Period trends
• All the atoms in the same period have
the same energy level.
• Same shielding.
• But, increasing nuclear charge
• So IE generally increases from left to
right.
• Exceptions at full and 1/2 full orbitals.
Driving Forces
• Full Energy Levels require lots of
energy to remove their electrons.
–Noble Gases have full orbitals.
• Atoms behave in ways to try and
achieve a noble gas configuration.
2nd Ionization Energy
• For elements that reach a filled or
half-filled orbital by removing 2
electrons, 2nd IE is lower than
expected.
• True for s2
• Alkaline earth metals form 2+ ions.
3rd IE
• Using the same logic s2p1 atoms
have an low 3rd IE.
• Atoms in the aluminum family
form 3+ ions.
• 2nd IE and 3rd IE are always
higher than 1st IE!!!
Trend in Ionization Energy
Ionization potential:
The energy required to remove the
valence electron from an atom. Largest
toward top right corner of PT since these
atoms hold on to their valence e- the
tightest.
Electronegativity
• Property of atoms in compounds
• Measure of the ability of an atom in a
chemical compound to attract electrons from
another atom in the compound
Electronegativity Period Trend
•
•
•
•
Metals are at the left of the table.
They let their electrons go easily
Thus, low electronegativity
At the right end are the nonmetals.
•
•
•
•
They want more electrons.
Try to take them away from others
High electronegativity.
See p.161 in text
Trend in Electron Affinity
Electron Affinity:
The energy release
when an electron is
added to an atom. Most
favorable toward top
right corner of PT since
these atoms have a
great affinity for e-.
The arrows indicate the trend:
Ionization energy, Electron Affinity, and
Electronegativity INCREASE in these
directions
Atomic size and Ionic size increase
in these directions:
Summary of Trends
1. Electron Configuration
2. Atomic Radii and Ionic Radii: Largest toward SW corner of PT
3. Ionization Energy: Largest toward upper right corner of PT
4. Electron Affinity and Electronegativity: Highest toward upper right of PT