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Periodic Table of Elements
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.
Elements

The elements,
alone or in
combinations,
make up our
bodies, our world,
our sun, and in
fact, the entire
universe.
Mendeleev – Father of the
Periodic Table

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In 1869, Dmitri Mendeléev created the
first accepted version of the periodic
table.
He grouped elements according to
their atomic mass, and as he did, he
found that the families had similar
chemical properties.
Blank spaces were left open to add
the new elements he predicted would
occur – scandium, germanium.
Glenn T. Seaborg (1912 – 1999)
Seaborg, McMillan, Kennedy, & Wahl, isolated element 94,
plutonium (Pu) in 1940.
Isolated Uranium-233 in 1941 & established Th's nuclear fuel
potential.
1944, Seaborg formulated the 'actinide (Ac) concept' of heavy
element electronic structure which predicted the actinides – specifically, the first eleven trans-uranium elements. This resulted
in a rearrangement of the periodic table.
The Ac series formed a transition series analogous to the rare
earth series of the lanthanide (La) elements.
Considered most significant change to the periodic table since
Mendeleev designed the table in the mid-1800’s. The actinide
concept showed how the trans-uranium elements fit into the
periodic table.
From 1944 -1958, Seaborg identified eight additional elements –
Americium (95), Curium (96), Berkelium (97), Californium (98),
Einsteinium (99), Fermium (100), Mendelevium (101), and
Nobelium (102). Element 106, Seaborgium, bears his name.
In 1951, Seaborg shared the Nobel Prize in Chemistry with E.M.
McMillan for "for their discoveries in the chemistry of the transuranium elements"
Reference:
U.S.
DOE. (2012). Glenn T.
Seaborg
Contributions to
Advancing Science.
Research and
Development from the
U.S. DOE.
Downloaded from:
“http://www.osti.gov/acc
omplishments/seaborg.
html”
Names of Elements

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Property – Sodium
from suda (Arabic
for headache)
Mineral in which
found – Lithium,
Boron, Argentum

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Place of discovery –
Germanium,
Californium, Yttrium,
Ytterbium, Terbium,
Erbium
Person (scientist) –
Curium, Einsteinium,
Seaborgium
Chemical Symbols - Berzelius

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Shorthand for element
Use first letter of name
(capital letter)
Second letter is
important in name - if
needed (lower case)

IUPAC has final
say: International
Union of Pure
and Applied
Chemistry
(Seaborgium
story)
Key to the Periodic Table

Elements are organized on our
modern periodic 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.
Atomic Number


Bohr Model of Hydrogen Atom
This refers to how many
protons an atom of that
element has.
No two elements, have
the same number of
protons.
Quantum Mechanical
Wave Model
Valence Electrons
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The number of valence
electrons an atom has can
be found from the last digit
of the group number for the
s and p blocks. (Arsenic is in
group 15 – has 5 valence
electrons)
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.
Properties of Metals

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Metals are good conductors
of heat and electricity.
Metals are shiny “luster”.
Metals are ductile (can be
stretched into thin wires).
Metals are malleable (can be
pounded into thin sheets).
Metals lose electrons to
become stable. They are the
“losers” and become
positively charged cations.
They lose their outer energy
level, becoming smaller as a
cation than they were as an
atom.
Properties of Non-Metals
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Sulfur
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.
They gain electrons to
form negative anions.
They become larger
when they gain elctrons.
Properties of Metalloids
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Silicon
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Metalloids are located
along the stairstep line.
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. Si and Ge are
used in computer chips.
They are ductile and
malleable.
Families

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Columns of elements are
called groups or families.
Elements in each family
have similar properties.
For example, lithium (Li),
sodium (Na), potassium
(K), and other members of
family IA are all soft,
white, shiny metals.
All elements in a family
have the same number of
valence electrons and
outer electron
configuration ending.
Periods
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Each horizontal row of
elements is called a period.
Elements in the period have
their valence electrons in
the same energy level.
The elements in a period
are not alike in properties.
In fact, the properties
change greatly across any
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
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Hydrogen sits on top of Group 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 one energy level.
Hydrogen only needs 2 electrons to fill
up its valence shell, but is actually more
stable with no electrons at all.
Alkali Metals
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The alkali metal 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 and end in s1 in the
same energy level as the
period.
They are shiny, have the
consistency of clay, and are
easily cut with a knife.
Alkali Metals
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Alkali
video
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They are the most reactive
metals.
They react violently with
water. More violent as
you go down group.
Alkali metals are never
found as free elements in
nature. They are always
bonded with another
element.
They lose 1 electron and
form a +1 cation.
What does it mean to be
reactive?
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We will be describing elements according to their
reactivity.
Elements that are reactive bond easily with other
elements to make compounds.
Some elements are only found in nature bonded
with other elements.
What makes an element reactive?



An incomplete valence electron level.
All atoms (except hydrogen) want to have 8 electrons in
their very outermost energy level (This is called the octet
rule.)
Atoms bond until this level is complete. Atoms with few
valence electrons (4 or less) lose them during bonding.
Atoms with 6, 7, or 8 valence electrons gain electrons
during bonding. They can also share electrons.
Elements, Compounds,
Mixtures

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Sodium is an element.
Chlorine is an
element.
When sodium and
chlorine bond they
make the compound
sodium chloride,
commonly known as
table salt.
Compounds have different properties
than the elements that make them up.
Table salt has different properties than
sodium, an explosive metal, and chlorine,
a poisonous gas.
5
Alkaline Earth Metals
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Have 2 valence electrons, lose them, become +2
cation.
Harder, more dense, stronger, and have higher
melting points than Group 1.
Not as reactive as Group 1.
They are never found uncombined in nature.
Halogen Family
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The elements in this family
are fluorine, chlorine,
bromine, iodine, and
astatine.
They react with alkali metals
to form salts.
Halogens have 7 valence
electrons, which explains
why they are the most
reactive non-metals.
Found only as compounds in
nature.
Halogen atoms only need
to gain 1 electron to fill their
outermost energy level, form
-1 anion.
All are diatomic elements –
F2, Cl2, Br2, I2
Noble Gases
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Noble Gases are colorless gases that are extremely un-reactive.
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 (none in nature), the noble gases are called inert.
Only xenon and krypton can be forced to form compounds in the
lab.
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.
Transition Metals
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Transition Elements
include those elements in
the middle d block.
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
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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-to-outermost d sublevel.
Most form +2 cations and others
Transition Elements

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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.
Inner Transition Metals - “Rare
Earth Elements”
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The thirty inner transition
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.
Not as rare as once
thought to be.
Chalcogen or Oxygen Family
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Atoms of this family have 6
valence electrons, gain 2 to be
-2 anion.
Most elements in this family also
share electrons when forming
compounds.
Chalcogen means “chalkformer”
Oxygen is the most abundant
element in the earth’s crust. It is
extremely active and combines
with almost all elements.
Tom
Lehrer
“Elements” song
Boron Family
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The Boron Family is
named after the first
element in the family.
Atoms in this family have 3
valence electrons, lose
them to form +3 cations.
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
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Atoms of this family have 4
valence electrons, form +2, +4
cations. Carbon can form
carbide, a -4 anion.
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
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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, but can gain 3
electrons to form -3 anions.
Other elements in this family
are phosphorus, arsenic,
antimony, and bismuth.
ELECTRON CONFIGURATION
can easily be determined from the
layout of the periodic table
We can find:
 Sublevel blocks
 Configuration ending
 Main (highest) energy level for
elements
 Number of valence electrons
SUBLEVEL BLOCKS

s sublevel
groups 1 and 2: first two groups and helium

p sublevel
groups 13-18: last six groups

d sublevel
groups 3-12: middle elements (transition)

f sublevel
Bottom two rows: lanthanide and actinide
series (inner transition)
VALENCE ELECTRONS
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The electrons in the highest energy level (s and
p sublevels) are called the valence electrons
Generally, these are the electrons used for
bonding
The number of valence electrons can be
determined from the last digit of the group
number for s and p blocks (middle all have 2)
For example: arsenic is in group 15 so it has 5
valence electrons
Stability depends on valence
electrons
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1. Octet Rule: having 8 valence
electrons makes an atom unreactive.
How does helium fit this rule? (full outer
level)
2. Filled or half-filled sublevel is more
stable (less reactive)
Cr: [Ar]4s23d4 → [Ar]4s13d5
Cu: [Ar]4s23d9 → [Ar]4s13d10
PERIODIC PATTERNS
Unit 4 – Periodic Table
WXHS, Honors Chemistry
ATOMIC RADIUS
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Position in the table and properties of the elements
are due to its electron configuration.
Radius is the distance from the center of the nucleus
to the “edge” of the electron cloud. The “edge” is
considered the 90% probability calculation of finding
the electron.
Atomic radii are usually measured in picometers
(pm) or angstroms (Å). An angstrom is1x10-10 m,
and a pm = 1x10-12 m.
ATOMIC RADIUS
Covalent Bond Determination
Since a cloud’s edge is difficult to define,
scientists calculate the atomic radii in one of
two manners. For covalently bound atoms, the
radius is simply ½ the distance from the center
of one atom to the other (in the diatomic
molecule).

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BROMINE = Br2
228
114
pm
pm
114
pm
ATOMIC RADIUS
Trends in the Periodic Table
ATOMIC RADII TRENDS

DOWN A FAMILY OR
GROUP
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I
n
c
r
e
a
s
e
s


WHY?
As you move down a
family, or a group, a new
energy level is added.
In other words, a new
floor is added to the
building with each new
period.
Shielding also increases
(more electrons in
between )
ATOMIC RADII TRENDS

ACROSS A PERIOD
DECREASE


S

WHY?
As you go across a period the
number of protons increases.
Thus, the nuclear charge
increases which pulls the
electrons in closer to the
nucleus.
This decreases the radius.
Radius size is affected by:

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Distance from nucleus – more energy
levels = bigger atom
Size of nuclear charge – more protons
in same period = smaller atom
Shielding effect of inner electrons –
lowers the attraction of the nucleus for
the outer electrons
SHIELDING EFFECT
As more electrons are added to
atoms, the inner layers of electrons
shield the outer electrons from the
positive charge of the nucleus.
The effective nuclear charge on
those outer electrons is less, and so
the outer electrons are less tightly
held.
Example of Shielding Effect
IONS

Metals
Lose electrons becoming
positive. Lose entire outer
level so they become
smaller than the original
atom.
i.e. Calcium – Ca
[Ar]4s2
Loses 4s2 e-s becoming Ca+2,
and the cation now has the
electronic configuration of
the Noble gas [Ar].
The cation obeys the Octet
Rule

Gain
NonMetals
electrons becoming
negative. Nuclear charge is
spread among more electrons,
so they are not held as tightly.
This makes them larger than the
original atom.
i.e. Chlorine – Cl
[Ne]3s23p5
Gains one e- becoming Cl-1 and
now has the [Ar] Noble gas
configuration.
The anion also obeys the Octet
Rule.
IONIC RADII TRENDS
Radii values given in (Angstroms) & pm
Scientific Reasoning
DOWN A FAMILY OR GROUP
INCREASES

When we proceed
down a group, the
energy level
increases with each
new period.

The outer energy level
is further away from the
nucleus, and the ion’s
radius increases.
IONIC RADII TRENDS
Radii values given in (Angstroms) & pm
SCIENTIFIC REASONING
ACROSS A PERIOD

DECREASES
then INCREASES

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There is a big jump in size when
cross from metal to nonmetal,
but otherwise follows same trend
as the atomic radii.
Metals are smaller than the atom
because they lose outer level.
When we look at the nonmetals,
the radius increases in size
because the atom, or anion,
gained electrons; thus, the outer
electron, or electrons, is (are)
not held as tightly by the
nucleus.
METALLIC ATOM AND ION COMPARISON
*Radii given in units of Angstroms
NONMETALLIC ATOM AND
ION COMPARISON
* Radii given in units of nm
Why
do the
Noble Gases not
have an ionic
Radius?
ATOM AND ION COMPARISON
*Radii given in units of nanometers (nm)
Why
does
Hydrogen not
have an ionic
Radius?
IONIZATION ENERGY
The energy required to remove an
electron from an atom is ionization
energy.
Successive ionization energies are the
respective energies required to remove
subsequent electrons from an ion.
Ionization energies are generally
measured in kilojoules per mole
(kJ/mol), and sometimes in electron
volts (eV).
PERIODIC TABLE TRENDS with
respect to IONIZATION ENERGY
The
larger the atom is, the easier its
electrons are to remove.
Ionization energy and atomic radius are
inversely proportional.
Increases across period – nuclear charge
increases so harder to remove
Decreases down group – shielding and
energy levels increase with each period
Metals have low IE, nonmetals have high IE
Measured in gas phase.
IONIZATION
ENERGY (cont)
Ionization Energy (eV) v Atomic Number
Dips
show where atom is
more stable with half or
filled sublevel. Between s
and p blocks (full s) and
between p3 and p4 (halffilled p sublevel)
PERIODIC TABLE TRENDS
in IONIZATION ENERGIES
INCREASES
INCREASES
ELECTRON AFFINITY
Electron Affinity is a measure of the energy
change when an electron is added to a neutral
atom, in the gaseous state, to form a negative ion.
This measures the ability to attract electrons as a
neutral atom.
Decreases down group-More energy levels and
shielding decrease attraction felt from nucleus
Increases across period – increased nuclear
charge
Energy is released as atom becomes more stable.
ELECTRON AFFINITY
Why
do the Alkaline Earth
Metals and Noble Gases not
have measurable Electron
Affinities?
ELECTRONEGATIVITY
Electronegativity is a measure of the
tendency of an atom, in a molecule, to
attract electrons in a bond when they
are being shared).
Trend is same as electron affinity and
for the same reasons.
Highest electronegativity – Fluorine
(most reactive nonmetal)
ELECTRONEGATIVITY of the
Elements in the Periodic Table
Reaction Tendencies are
Determined by:
Number of energy levels
 Shielding effect of inner electrons
 Size of nuclear charge
 Stability of filled or half-filled sublevels
ISOELECTRONIC SPECIES
 When a noble gas atom and ions all have
the same configuration
 Examples: N3-, O2-, F-, Ne, Na+, Mg2+, Al3+

INCREASES
METALLIC TREND
Summary of the major trends in the
Periodic Table Trends
Electronegativity
Electronegativity