Transcript Alkaline Earth Metals

The Periodic Table
Concepts to Master
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Who was important in the development of
the periodic table and why?
What is the difference between a chemical
and physical property?
What are elements arranged on the periodic
table horizontally?
What are elements arranged on the periodic
table vertically?
How many electrons can each orbital
contain?
Why can the periodic table be subdivided
into s,p,d, and f blocks?
What are the other names for the s, p, d,
and f blocks?
Is atomic radius a chemical or physical
property?
Is electronegativity a chemical or physical
property?
Why are the noble gases inert?
Why do the alkali metals react so readily
with the halogens?
Why do the transition metals have multiple
oxidation numbers?
What is the trend in metallic character as
you go down the periodic table?
What is the common theme as you go down
the periodic table?
What is the diagonal rule?
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Where are electrons located?
What are the characteristics of metals?
What are the characteristics of nonmetals?
What are the characteristics of metalloids?
Explain periodicity?
How are sublevels and PELs related?
How are orbitals and orbits related?
What’s the formula for the maximum number
of electrons allowed in a certain PEL?
List the elements that exist as diatomics.
What are the trends in atomic radius as you
go down and across the periodic table?
Why?
What are the trends in electronegativity as
you go down and across the periodic table?
Why?
What are the trends in ionization energy as
you go down and across the periodic table?
Why?
What is the most reactive metal and the
most reactive non-metal?
What is the common theme as you go across the
periodic table?
Vocab
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Alkali Metals
Alkaline Earth Metals
Anion
Atom
Atomic Radius
Boiling point
Cation
Density
Diatomic
Diatomic
Ductile
Electronegativity
Groups
Halogens
Inert Gases
Ionization Energy
Kernel
Lustrous
Malleable
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Melting point
Metalloids
Metals
non metals
Modern Periodic Law
Monatomic
Orbitals
Orbits
Periods
Periodic Law
Polyatomic
Principle energy level
Principle quantum number
Reactivity
Reactivity
Sublevels
Transition Metals
Valence electrons
Valence shell
Labs: Graphing Trends and Constructing a Table
Cool Websites
• http://www.colorado.edu/physics/2000/app
lets/a2.html
• http://www.youtube.com/watch?v=DYW50
F42ss8
– Element song
People and Periodic Table
• Dmitri Mendeleev
– Credited with organization of FIRST
periodic table
– Mendeleev’s greatest achievement
was recognizing the fundamental rule
that the chemical elements show an
approximate repetition in their
properties.
• Elements were arranged by increasing
atomic mass.
• Elements were listed in columns so that
those with similar properties were side by
side.
• He predicted the existence and properties
of new elements (blank spaces in the first
periodic table).
People and Periodic Table
• Henry Moseley
– With the discovery of isotopes of
the elements, it became apparent
that atomic mass was not the
significant player in the periodic
law as Mendeleev, had proposed.
– Moseley used X-rays to determine
the atomic number of the known
elements and then arranged them
according to increasing atomic
number.
– Because of Moseley's work, the
modern periodic table is based on
the atomic numbers of the
elements.
pg186
Mosely
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Student in Rutherford ‘s lab at the univ of Manchester
Remember Rutherford was playing with radioactive sources to do gold-foil
exp so he had access to xrays
Found a mathematical relationship between the amt of energy in the xray
beam (wavelength) and the # of protons in nucleus.
This was huge because at the time Rutherford’s idea of a nucleus wasn’t
“proven” (not enough scientists had repeated it) so it was too tentative for
other scientists to accept.
Mosely’s work confirmed Rutherfords nucleus conclusion .
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His work could be repeated by anyone
Linked the order of the elements with a physical characteristic based on the atoms structure
Now scientists knew what to look for when searching for new elements
He died in battle during WWI – he was 27
Periodic Law
• Mendeleev - The properties of the elements are a periodic function
of their atomic masses.
• Moseley - The properties of the elements are a periodic function of
their atomic numbers.
• Modern Periodic Law states that many of the physical and chemical
properties of the elements tend to recur in a systematic manner with
increasing atomic number.
– Periods are the horizontal rows in the table.
– Progressing from the lightest to the heaviest atoms, certain properties of
the elements approximate those of precursors at regular intervals of 2,
8, 18, and 32 (periodicity).
– Examples:
• The 2d element (helium) is similar in its chemical behavior to the 10th
(neon), as well as to the 18th (argon), the 36th (krypton), the 54th (xenon),
and the 86th (radon).
Pg 187
• The chemical family called the halogens, composed of elements 9 (fluorine),
17 (chlorine), 35 (bromine), 53 (iodine), and 85 (astatine), is an extremely
reactive family.
Electron Location
• Kernel of an atom is the nucleus and all the electrons but the
valence electrons.
• The elements are arranged vertically in columns of the periodic table
called GROUPS or FAMILIES.
• Group # indicates the number of valence electrons.
• Because of Periodicity, the elements with the same # of valence
electrons are in the same group.
• These electrons influence the chemical and physical properties of
elements the most.
• Electron Configuration shows the location of all the electrons for
the atom.
Physical and Chemical properties
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Because of Periodicity, the elements with the same # of valence electrons
are in the same group so they share similar chemical and physical
properties.
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Chemical properties of matter describe its "potential" to undergo some
chemical change or reaction by virtue of its composition. What elements,
electrons, and bonding are present to give the potential for chemical
change. The result of the change is the formation of a new substance.
– Toxicity
– Flammability
– Reactivity
• Electronegativity
• Ionization Energy
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Physical properties can be observed or measured without changing the
composition of matter. They are used to observe and describe matter.
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Atomic radius
Density
Melting Point
Boiling Point
Color
Solubility
Odor
Conductivity
Transition Metals
Lanthanide Series
Actinide Series
Noble Gases
or Inert Gases
Halogens
Alkaline Earth Metals
Alkali Metals
Because of periodicity….
• The alkali metals are silvercolored
• Soft solids (Fr and Cs are liquids)
• The first three are biologically
important
• low-density metals
• react readily with halogens
• react readily with water
• one valence electron
– so they want to lose an electron and
achieve a noble configuration (which
is?)
– form +1 cations
1A
NOT
Alkali Metals
Alkali Metals
Alkaline Earth Metals
– so they want to lose 2 electrons and
achieve a noble configuration (an octet)
– form +2 cations
1A
2A
Alkaline Earth Metals
silvery colored
Soft solids
Ca and Mg have biological functions
low-density metals
react readily with halogens
react readily with water - though not as
rapidly as the alkali metals
• Beryllium is an exception: It does not
react with water
• two valence electron
Alkali Metals
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Transition Metals
• They often form colored compounds.
• They are often good catalysts
– lowers activation energy so rxns are
faster
– not used up in the rxn
– enzymes
• They are silvery-blue at room
temperature (except copper and gold)
- lustrous.
• Malleable
• They are solids at room temperature
(except Hg)
• Partly filled d sublevel
B group
Transition Metals
– They can have a variety of different
charged cations
– 4s fills before 3d (clouds are more
apparent and overlapping occurs)
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chromium
Iron
Vanadium
Silver (5s fills before 4f)
• Good conductors of electricity (Why?)
Thus they are transitioning
between the filling of their
outermost orbitals.
Transition Metal Colors
Lanthanide
• Rare Earth Metals
• Silvery-white metals that
tarnish when exposed to
air
• Relatively soft metals
• Very reactive
• Many rare earth
compounds fluoresce
strongly under ultraviolet
light
Lanthanide Series
Actinide
• All are radioactive
• The metals tarnish
readily in air
• Actinides are very
dense metals
• Actinides combine
directly with most
nonmetals
Actinide Series
Metallic Characteristics
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Conduct electricity and heat
Dense
Malleable (bendable to form shapes – jewelry)
Ductile (able to be drawn out into wires)
Lustrous (shiny)
Reactive
High Melting point
Solid at RT
Metalloids
• Share properties with
both metals and
nonmetals
• Solids
• Semi-conductors
(between a conductor
and an insulator)
• Form cations or anions. #
of valence electrons
varies.
• 4 sit on steps and 2 are
beneath.
8A
Less
3A 4A 5A 6A 7A
metallic
metalloid
Poor
Metals
More
metallic
metalloid
NonMetals
8A
5A 6A 7A
Noble gases
4A
Halogens
• poor conductors of heat and
electricity
• in solid form, they are dull and brittle
• usually have lower densities than
metals
• most of the crust, atmosphere and
oceans are made up of nonmetals.
• Bulk tissues of living organisms are
composed almost entirely of
nonmetals.
• Many nonmetals (hydrogen,
nitrogen, oxygen, fluorine, chlorine,
bromine, and iodine) are diatomic,
and most of the rest are polyatomic
• Prefer to form anions - gaining
electrons to achieve an octet.
Halogens
• This high reactivity is due to their atoms being one
electron short of a full outer shell.
8A
• Halogens are highly reactive
• They form diatomic molecules (F2, Cl2, Br2, I2)
• All three states of matter are represented
– fluorine and chlorine are gases
– bromine is a liquid
– iodine and astatine are solids
Noble Gases
• Both chlorine and bromine are used as
disinfectants
7A
Halogens
– harmful or lethal to biological organisms in sufficient
quantities.
– Fluorine is the most reactive element in existence
– This high reactivity is due to their atoms being one
electron short of a full outer shell of electrons. They
form -1 anions.
Diatomics
• hydrogen, nitrogen, oxygen, fluorine, chlorine, bromine,
and iodine
• Diatomic elements are nonmetal elements that form a
covalent bond between two atoms.
• As elements they always travel in pairs of atoms and
therefore you must write then as:
• H2, N2, O2, F2, Cl2, Br2, I2
Noble Gases or Inert Gases
– They have the maximum number
of valence electrons their outer
shell can hold.
Noble Inert Gases
• odorless, colorless, monatomic
gases.
• Lighting (Ne), welding and
space technology (processes
performed under Ar so that no
unwanted chem rxns occur) .
• Stable or unreactive
8A
PEL / PQN
• Probable Location according to the wave mechanical
model
• Principle Energy Levels (PEL) or Principle Quantum
Number (PQN) = n
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the total number of orbits around the nucleus
Period # = n
Max number of e in that PEL, PQN = 2n2
When n = 1, max # of e = 2
When n = 2, max # of e = 8
When n = 3, max # of e = 18
When n = 4, max # of e =
When n = 5, max # of e =
When n = 6, max # of e =
When n = 7, max # of e =
Sublevels
• Sublevels exist in each orbit (PEL)
– Shape of the electron cloud that is created by fast moving
electrons.
– spdfg
– The number of sublevels present in each PEL also = n, so PEL
5 contains 5 sublevels. They are 5s, 5p, 5d, 5f, and 5g
Element
(Neutral)
Galium
E shown on
the periodic
table
2
8
18
3
PQN (4th
period)
1
2
3
4
Sublevels
1s
2s
2p
3s
3p
3d
4s
4p
4d
4f
Orbitals
• Each sublevel has orbitals
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The possible orientations of the shapes around the x,y, and z axis
s sublevel has 1 orbital
p sublevel has 3 orbitals
d sublevel has 5 orbitals
f sublevel has 7 orbitals
Element
(Neutral)
Galium
E shown on
the periodic
table
2
8
18
3
PQN
1
2
3
4
Sublevels
1s
2s
2p
3s
3p
3d
4s
4p
4d
4f
Orbitals
1
1
3
1
3
5
1
3
5
7
s sublevel orbitals
p sublevel orbitals
d sublevel orbitals
Electrons in Orbitals
• Each orbital can contain 2 electrons maximum
• 2n2
Element
(Neutral)
Galium
E shown on the
periodic
table
2
8
18
3
PQN
1
2
3
4
Sublevels
1s
2s
2p
3s
3p
3d
4s
4p
4d
4f
Orbitals
1
1
3
1
3
5
1
3
5
7
# of E possible
2
2
6
2
6
10
2
6
10
14
Actual # of E
2
2
6
2
6
10
2
1
E configuration
1s2
2s2
2p6
3s2
3p6
3d10
4s2
4p1
Determining spdf configuration
using THE Table
The number of columns present in the block equals the
number of possible electrons for that sublevel.
f Block location if inserted
Really know the location of your
valence electrons
Using PT to determine detailed
electron configuration
A short cut…?
• Diagonal rule - A
guideline explaining the
order in which electrons
fill the orbital levels.
• Pros
– Easy to use if given to you
• Cons
– Memorization required
– There are exceptions to
this rule when filling the
orbitals of heavier
elements.
7s
7p
7d
7f
6s
6p
6d
6f
5s
5p
5d
5f
4s
4p
4d
4f
3s
3p
3d
2s
2p
1s
Using Diagonal Rule compared to using PT for
Determining detailed electron configuration
Use PT for determination of Cr
Use diagonal rule for determination of Cr
Use PT for determination of Pd
Use diagonal rule for determination of Pd
7s
7p
7d
7f
7s
7p
7d
7f
6s
6p
6d
6f
6s
6p
6d
6f
5s
5p
5d
5f
5s
5p
5d
5f
4s
4p
4d
4f
4s
4p
4d
4f
3s
3p
3d
3s
3p
3d
2s
2p
2s
2p
1s
1s
Summary
• Elements are arranged by increasing atomic #
across the periodic table.
• Periods are the horizontal rows. The period # =
PEL.
• Elements are grouped vertically by similar
chemical and physical properties.
• Groups (or Families) are the vertical columns.
The group # = # of valence electrons.
Textbook pg 186-191
Practice problems 5.1-5.3
Atomic Radius
• Half the distance between the nuclei of two like atoms
in a diatomic molecule.
• Atom size vs ion size
Combine to form O2
Pg 187-188
As you go down a group,
Atomic Radius increases
As you go across a period,
Atomic Radius decreases
Why?
• As you go down a group the number of PELs
increases, more electrons are present to fill
these energy levels, so atomic radius increases.
• As you go across a period, atomic # increases
which means that the # of protons in the nucleus
increases, so nuclear charge is increasing and
attracting electrons with a greater force.
Opposite attract.
Ionization Energy
• The amount of energy required to remove an outer electron
• The more difficult it is to remove an electron, the greater
the ionization energy
• smaller atoms have greater ionization energy since the
valence electrons are closer to the nucleus and more
strongly attracted and, therefore, more difficult to remove
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• X + energy → X+ + e• First Ionization Energy
• X + energy → X+ + e-
• Second Ionization Energy (greater than 1st)
• X+ + energy → X+2 + e-
• Third Ionization Energy (greater than 2nd)
• X+2 + energy → X+3 + e-
Pg 189
As you go down a group,
Ionization Energy decreases
Group 2 and 18
2100
ionization energy
If the ionization energy is high,
that means it takes a lot of
energy to remove the outermost
electron. If the ionization energy
is low, that means it takes only a
small amount of energy to
remove the outermost electron
1600
1100
600
100
-400
0
10
20
30
40
atomic number
Top trend is group 18
Bottom trend is group 2
50
60
As you go across a period,
Ionization energy increases
Why?
• As you go down a group the number of PELs
increases. Attraction is less for the electrons
furthest from the nucleus so it takes less energy
for electrons to be pulled away.
• As you go across a period, nuclear charge is
increasing and attracting electrons with a greater
force. Since that force is increasing, it takes
more energy for the electrons to be pulled away.
Going towards inert gases which have a full
valence shell and are extremely resistant to give
up any electrons.
Electronegativity
• An atom’s affinity for
electrons
• Arbitrary scale from 04
– 0 is least
electronegative
– 4 is most
electronegative
Pg 189-191
Fluorine is the
most electronegative
4
Francium is the
least electronegative
0.7
Neon and the other
noble gases
have an
Electronegativity of
0
As you go down a group,
Electronegativity decreases
Electronegativity is a
measure of the
tendency of an atom to
attract electrons.
The arbitrary scale of 04 is the most commonly
used. Fluorine (the most
electronegative
element) is assigned a
value of 4.0.
As you go across a period,
Electronegativity increases
Why?
• As you go down a group the number of
PELs increases. Electron attraction to the
nucleus is less when they are farther from
the nucleus.
• As you go across a period, nuclear charge
is increasing and thus attracting electrons
to a greater extent.
• Why do the inert gases have an
electronegativity of 0?
Chemical Reactivity
• Francium is the most reactive metal.
• Fluorine is the most reactive non-metal.
Chemical
Reactivity
Electronegativity
Ionization Energy
Elements on opposite Sides of the
Periodic Table are attracted to
each other.
• Sodium likes to combine with Chlorine – Why?
• Atoms become cations due to less ionization
energy.
– METALS
– Fr
• Atoms become anions due to high
electronegativity.
– NONMETALS
–F