Download Unit 9: Periodicity

Unit 5: Periodicity
History of the PT
 Dimitri Mendeleev (1836-1907)
– Russian chemistry professor
 In writing a textbook of general chemistry, he compiled vast
quantities of data on the elements. He realized as he examined
the data that the chemical properties of the elements repeated
themselves when the elements were placed in order of atomic
mass. Based on his realization, Mendeleev arranged the
elements in order of increasing atomic mass, and in such a way
that elements with similar chemical properties fell in the same
column. He published a primitive version of today’s periodic
table in 1869. It contained the 62 elements that had been
discovered at that time.
Genius of Mendeleev’s Work
 Left spaces for
elements not yet
discovered.
 He predicted that
some still-unknown
elements must exist to
fit in the holes.
Dmitri Mendeleev (1869)
http://www.chem.msu.su/eng/misc/mendeleev/welcome.html
Inductive vs. Deductive Reasoning
 Inductive reasoning is the use of detailed
facts to form a general principle or model
(going from specific to general). How did
Mendeleev use inductive reasoning?
 Deductive reasoning is the use of a
general principle or model to draw specific
inferences (going from general to specific).
How did Mendeleev use deductive
reasoning?
Mendeleev’s Periodic Table
 Contained 1 inconsistency.
– He placed the elements in order of atomic mass
 Forced to break pattern a couple times to preserve
the patterns he had discovered.
Henry Moseley
 Shortly after Rutherford’s discovery of
the proton in 1911, Henry Moseley
(1887-1915) did experiments with xrays to determine the number of
protons in various elements. When
Moseley arranged the elements of
Mendeleev’s periodic table according to
increasing atomic number and not
atomic mass, the inconsistencies
associated with Mendeleev's table were
eliminated. The modern periodic table is
based on Moseley's arrangement by
atomic number. At age 28, Moseley was
killed in action during World War I. As a
direct result, Britain adopted the policy
of exempting scientists from fighting in
wars.
Periodic Law
 From Mendeleev’s and Moseley’s work comes the
Periodic Law: The properties of the elements are
periodic functions of their atomic numbers.
 What this means is that if we arrange the elements
in order of increasing atomic number, we will
periodically encounter elements that have similar
chemical and physical properties. These elements
appear in the same vertical column (group).
Elements
 Science has come along
way since Aristotle’s theory
of Air, Water, Fire, and
Earth.
 Scientists have identified
90 naturally occurring
elements, and created
about 28 others.
Periodic Table
 The periodic table organizes the elements in a
particular way. A great deal of information about an
element can be gathered from its position in the
period table.
 For example, you can predict with reasonably
good accuracy the physical and chemical
properties of the element. You can also predict
what other elements a particular element will react
with chemically.
 Understanding the organization and plan of the
periodic table will help you obtain basic
information about each of the 118 known
elements.
Key to the Periodic Table
 Elements are organized on
the table according to their
atomic number, usually
found near the top of the
square.
– The atomic number
refers to how many
protons an atom of that
element has.
– For instance, hydrogen
has 1 proton, so it’s
atomic number is 1.
– The atomic number is
unique to that element.
No two elements have
the same atomic
number.
What’s in a square?
 Different periodic tables
can include various bits
of information, but
usually:
– atomic number
– symbol
– atomic mass
– number of valence
electrons
– state of matter at
room temperature.
Periodic Table Expanded View
The way the periodic table usually
seen is a compress view, placing
the Lanthanides and actinides at
the bottom of the stable.
The Periodic Table can be arrange
by sublevels. The s-block is Group I
and & 2, the p-block is Group 13 - 18.
The d-block is the transition metals,
and the f-block are the Lanthanides
and Actinide metals
Valence Electrons
 The number of valence
electrons an atom has
may also appear in a
square.
 Valence electrons are
the electrons in the
outer energy level of an
atom.
 These are the
electrons that are
transferred or shared
when atoms bond
together.
Nonmetals
Metals
Properties of Metals
 Metals are good
conductors of heat and
electricity.
 Metals are shiny.
 Metals are ductile (can
be stretched into thin
wires).
 Metals are malleable
(can be pounded into
thin sheets).
 A chemical property of
metal is its reaction with
water which results in
corrosion.
Properties of Non-Metals
 Non-metals are poor
conductors of heat and
electricity.
 Non-metals are not
ductile or malleable.
 Solid non-metals are
brittle and break easily.
 They are dull.
 Many non-metals are
gases.
Sulfur
Properties of Metalloids
 Metalloids (metal-like)
have properties of both
metals and non-metals.
 They are solids that can
be shiny or dull.
 They conduct heat and
electricity better than nonmetals but not as well as
metals.
 They are ductile and
malleable.
Silicon
 Columns of elements are
called groups or families.
 Elements in each family
have similar but not
identical properties.
 For example, lithium (Li),
sodium (Na), potassium
(K), and other members of
family 1 are all soft, white,
shiny metals.
 All elements in a family
have the same number of
valence electrons.
 Each horizontal row of elements is called a period.
 The elements in a period are not alike in properties.
 In fact, the properties change greatly across even given
row.
 The first element in a period is always an extremely
active solid. The last element in a period, is always an
inactive gas.
Hydrogen
 The hydrogen square sits atop Family 1, but
it is not a member of that family. Hydrogen is
in a class of its own.
 It’s a gas at room temperature.
 It has one proton and one electron in its one
and only energy level.
 Hydrogen only needs 2 electrons to fill up its
valence shell.
Alkali Metals
 The alkali family is found
in the first column of the
periodic table.
 Atoms of the alkali
metals have a single
electron in their
outermost level, in other
words, 1 valence
electron.
 They are shiny, have the
consistency of clay, and
are easily cut with a
knife.
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 bonded with
another element.
Alkaline Earth Metals
 They are never found uncombined in nature.
 They have two valence electrons.
 Alkaline earth metals include magnesium and
calcium, among others.
Transition Metals
 Transition Elements
include those elements in
the B families.
 These are the metals you
are probably most familiar:
copper, tin, zinc, iron,
nickel, gold, and silver.
 They are good conductors
of heat and electricity.
Transition Metals
 The compounds of
transition metals are
usually brightly
colored and are
often used to color
paints.
 Transition elements
have 1 or 2 valence
electrons, which
they lose when they
form bonds with
other atoms. Some
transition elements
can lose electrons in
their next-tooutermost level.
Transition Elements
 Transition elements have properties similar
to one another and to other metals, but their
properties do not fit in with those of any
other family.
 Many transition metals combine chemically
with oxygen to form compounds called
oxides.
Boron Family
 The Boron Family is named
after the first element in the
family.
 Atoms in this family have 3
valence electrons.
 This family includes a
metalloid (boron), and the
rest are metals.
 This family includes the
most abundant metal in the
earth’s crust (aluminum).
Carbon Family
 Atoms of this family have 4
valence electrons.
 This family includes a nonmetal (carbon), metalloids,
and metals.
 The element carbon is
called the “basis of life.”
There is an entire branch
of chemistry devoted to
carbon compounds called
organic chemistry.
Nitrogen Family
 The nitrogen family is named
after the element that makes
up 78% of our atmosphere.
 This family includes nonmetals, metalloids, and
metals.
 Atoms in the nitrogen family
have 5 valence electrons.
They tend to share electrons
when they bond.
 Other elements in this family
are phosphorus, arsenic,
antimony, and bismuth.
Oxygen Family
 Atoms of this family have
6 valence electrons.
 Most elements in this
family share electrons
when forming compounds.
 Oxygen is the most
abundant element in the
earth’s crust. It is
extremely active and
combines with almost all
elements.
Halogen Family
 The elements in this
family are fluorine,
chlorine, bromine, iodine,
and astatine.
 Halogens have 7 valence
electrons, which explains
why they are the most
active non-metals. They
are never found free in
nature.
 Halogen atoms only need
to gain 1 electron to fill
their outermost energy
level.
 They react with alkali
metals to form salts.
Noble Gases
 Noble Gases are colorless
gases that are extremely unreactive.
 One important property of
the noble gases is their
inactivity. They are inactive
because their outermost
energy level is full.
 Because they do not readily
combine with other elements
to form compounds, the
noble gases are called inert.
 The family of noble gases
includes helium, neon, argon,
krypton, xenon, and radon.
 All the noble gases are found
in small amounts in the
earth's atmosphere.
Rare Earth Elements
 The thirty rare earth
elements are composed
of the lanthanide and
actinide series.
 One element of the
lanthanide series and
most of the elements in
the actinide series are
called trans-uranium,
which means synthetic
or man-made.
Octet Rule
 Is 8 valence electrons.
 Is associated with the stability of the noble gases.
 Helium (He) is stable with 2 valence electrons
 Valence Electrons
–
–
–
–
He 1s2
Ne 1s2 2s2 2p6
Ar 1s2 2s2 2p6 3s2 3p6
Kr 1s2 2s2 2p63s2 3p6 4s2 3d10 4p6
2
8
8
8
Electron Shielding
Shielding electrons:
electrons in the
energy levels
between the nucleus
and the valence
electrons. They are
called "shielding"
electrons because
they "shield" the
valence electrons
from the force of
attraction exerted by
the positive charge in
the nucleus.
Video
Periodic Trends
 Trends in properties of the elements that follow a
pattern down a group and across a period in the
periodic table.
1
2
3
4
5
6
7
Trends in Atomic Radius (size)
 atomic radius: distance from center
of nucleus to edge of electron cloud
 group trend: increases going down
a group.
– As you move down a group, energy
levels are added, thus increasing the
size of the electron cloud, so the atoms
get larger.

 periodic trend: decreases going left
to right across a period.
– [NOTE: from now on, “across a
period” will refer to the left-to-right
direction]
Atomic Radius
 Atomic Radius
– Increases to the LEFT and DOWN
1
2
3
4
5
6
7
Atomic Radius
Trends in Ionization Energy
 ionization energy: the measure of the energy required to
remove an electron from the outermost energy level of an
atom.
 group trend: decreases going down a group
– This is due to the shielding effect - an electron in the outer energy
level of a large atom is easier to remove because it is well-shielded
from the pull of the nucleus by the inner electrons.
 periodic trend: increases going across a period
– This is due to nuclear charge - across a period, nuclear charge
increases, so it becomes more difficult to remove an electron (held
tighter).
– Note that this periodic trend supports the idea that metals have a
much greater tendency to lose electrons than nonmetals do.
Ionization Energy
 First Ionization Energy
– Increases UP and to the RIGHT
1
2
3
4
5
6
7
Trends in Electron Affinity
 electron affinity (electron-liking): the energy change that
accompanies the addition of an electron to an atom.
– *If energy is released in the process of adding an electron, EA is negative. A
negative EA means that the atom wants to gain the electron. The larger the negative
number, the more it wants to gain the electron. It is a favorable process. Nonmetal
atoms have large negative EAs.
– *If energy is absorbed in the process of adding an electron, EA is positive. A
positive EA means that the atom does not want to gain the electron. The larger the
positive number, the more it does not want to gain the electron. It is an unfavorable
process. Metal atoms have small negative or positive EAs.
 group trend: EA decreases (less favorable/value is more +) going
down a group.
– It is harder to add the electron when it is farther from the nucleus. The nucleus can’t
“grab onto it” as well
 periodic trend: EA increases (more favorable/value is
more - ) going across a period.
– This is due to nuclear charge - across a period, nuclear charge increases, so it
becomes easier to add an electron.
– Note that this periodic trend supports the idea that nonmetals have a much greater
tendency to gain electrons than metals do.
Electron Affinity
 Electron Affinity
– Increases UP and to the RIGHT
1
2
3
4
5
6
7
Trends in Electronegativity
 electronegativity: the tendency of an atom to attract
electrons to itself when it is chemically bonded with another
element.
– numerical scale which combines both ionization energy and electron
affinity.
– can be used to predict whether atoms will form ionic or covalent
bonds in molecules.
 Diagram of water molecule:
– In H2O, oxygen is more electronegative than hydrogen, so it pulls
the electrons closer, and thus obtains a partially negative charge.
 group trend: decreases down a group.
– Larger atoms have more energy levels, so it is harder for them to
attract electrons to the nucleus (shielding effect).
 periodic trend: increases across a period.
– Nonmetallic character increases across a period, and nonmetals
attract electrons more than metals do.
Electronegativity
 Electronegativity
– Increases UP and to the RIGHT
1
2
3
4
5
6
7
Summary of Periodic Trends
 Ionization Energy, Electron Affinity, and
Electronegativity
– Increases UP and to the RIGHT
 Atomic Radius
– Increases to the LEFT and DOWN
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