Transcript Chapter 5: The Periodic Law

Chapter 5:
The Periodic Law
5.1: History of the Periodic Table
5.2: Electron Configuration & the Periodic Table
5.3: Electron Configuration & Periodic Properties
5.1: History of the Periodic
Table
• September 1860- International Congress of
Chemists in Karlsruhe, Germany.
• Problem: how to determine the atomic
mass… so that everyone is getting the
same #.
• Italian Chemist , Stanislao Cannizzaro’s
method of determining atomic masses was
accepted- created a standard value.
Döbereiner
• In the early 1800’s J.W. Döbereiner began
classifying elements into groups of three
(TRIADS). For example, lithium, sodium and
potassium; calcium, strontium and barium; AND
chlorine, bromine, and iodine are three sets of
triads. The elements within the triads
share similar chemical properties.
Newlands
• Then in 1865, J.A.R Newlands
noticed that when the known
elements were arranged in
increasing atomic mass, regular
patterns would occur, AND they
would repeat every 8 elements (law
of octaves).
Mendeleev & Chemical
Periodicity
• In 1869, Dmitri Mendeleev (and Lothar Meyer
independently) published practically identical
schemes for classifying elements.
Mendeleev also noticed regular patterns in
the elements when they were arranged by
increasing atomic mass. He produced the
first periodic table of elements and actually
left spaced to be filled in by undiscovered
elements. He was able to predict there
would be more and about where they
would go!
Mendeleev's Periodic Table of the
Elements
Moseley & the Periodic Law
• The periodic table Mendeleev came up with
was a good start but it wasn’t perfect. It
wasn’t until 1913 when H.G.J. Moseley
determined that the correct way to organize
the periodic table is by ATOMIC NUMBER.
• The PERIODIC LAW- when elements are
arranged in order of increasing atomic
number, their physical and chemical
properties show a periodic pattern.
Increases in
atomic #
The Modern Periodic Table
• Periodic Table- is an arrangement of the elements
in order of their atomic numbers so that elements
with similar properties fall in the same column, or
group.
5.2
• Organizing the Squares:
– When you look at the modern Periodic Table of Elements
(PTE) you will notice it is arranged in to columns (vertically)
and rows (horizontally). The columns are called GROUPS or
FAMILIES (18). The rows are called PERIODS (7).
Periods
G
r
o
u
p
s
Labeling & Naming Groups
• There are 3 methods for numbering the columns (the American,
the European and the IUPAC) The only 1 your book uses is the
IUPAC.
– IUPAC: 1-18
– AMERICAN: 1A, 2A, 3B-8B, 11B, 12B, 3A-8A
• The families are given names:
–
–
–
–
1- ALKALI METALS (except hydrogen)
2-ALKALINE EARTH METALS
17- HALOGENS
18- NOBLE GASES
(on the PTE I gave you)
Electronic Configuration &
the PTE:
• The electrons in the outermost energy level are
called the VALENCE electrons.
– These are the electrons which will allow one atom to
interact with another atom and are responsible for the
elements chemical properties.
• The members in each group/family have the same
# of valence electron and that is why they share
similar chemical properties.
The s-, p-, d- and f-block
Elements:
•
•
•
•
S-BLOCK- groups 1&2
P-BLOCK- groups 13-18
D-BLOCK- groups 3-12
F-BLOCK- the lanthanide & actinide series (the two rows at the very
bottom of the PTE)
The s-Block: groups 1&2
• The metals that are the most
reactive are found in the two-column
s-block of the PTE. These are called
the alkali metals and the alkaline
earth metals.
ALKALI METALS
• These are the metals in Group 1 (1A): Li, Na, K, Rb,
Cs, and Fr. The name for the family is derived from
the Arabic word meaning “ashes” because we find
Na and K in the ashes of burnt plants.
• *Properties* These metals are: malleable, ductile
and good conductors with low densities and low
melting points. They are “soft” and are able to be
cut with a knife and when exposed to the air, tarnish
easily.
• THESE METALS ARE EXTREMELY REACTIVE!!!
– The react readily with water and air and must be stored
in evacuated containers… or placed in oil.
ALKALINE EARTH METALS
• This comprises the elements in Group 2 (2A). Be, Mg,
Ca, Sr, Ba, and Ra.
• *Properties* These metals have higher densities and
higher melting points as well as higher ionization
energies. Ergo they are not as reactive as the alkali
metals. They all have 2 valence electrons which are
fairly easy to lose, forming a 2+ ion. In order for the
Alkaline earth metals to combine with halogens you
will need 2 halogens for every 1 alkaline earth metal
(MX2).
Hydrogen: One of a Kind
• It has 1 valence electron with a high ionization energy but
will form a + ion. It is a nonmetal and a gas at room
temperature which is odorless, colorless and also diatomic
(H2).“Most of the Earth’s hydrogen is combined with oxygen
as water”. Remember… it also combines readily with C to
make hydrocarbons (fuels). It is the 9th most abundant
element on earth but the absolute most abundant element
in the universe!
– We get H from hydrocarbons, commercially. And we use H in
ammonia, which can be used as a fertilizer, and in other organic
compounds like alcohols, gases (methane), etc.
The d-Block: groups 3-12
• Transition metals play an important role in
living organisms, and are also extremely
valuable as strong, structurally useful
materials. From iron captured hemoglobin
and cobalt in vitamin B12 to gold and silver
jewelry, there are many different and
important uses for the transition metals.
• *Properties* They vary from family to family
but most have high densities and melting
points, are malleable, ductile and good
conductors.
The p- Block: groups 13-18
• P-block & s-block elements constitute the maingroup elements.
• Properties- very greatly… includes all of the
nonmetals and 6 metalloids.
– Metalloids- aka semiconductors, fall on both sides of the
line that separate nonmetals and metals.
– Other metals
– Other non-metals
– Halogens
– Nobel gases
THE HALOGENS
• The name comes from the Greek meaning “salt
former” because most will form salts. They each
have 7 valence electrons and will accept ONE
electron to form a – ion. ALL of the halogens are
diatomic and react with most metals and some
nonmetals. They are HIGHLY REACTIVE and are
found combined with other elements in nature.
– WHY? Because all they need in 1 more electron and
will get it from any element that is willing to give or share
it!
THE NOBLE GASES
• The most abundant is Argon and it was the first be
discovered (1894). However, it does not form any
compounds (argon means “lazy one” in Greek). Now… a
whole new, previously unknown and not predicted group of
element has been discovered, too. Soon after the other
noble gases were found and added to the PTE.
– It was originally though that the noble gases never formed
compounds until 1962 when XeF2, XeF4, and XeF6 were formed.
Then came KrF2, but to date there are no know compounds of
He, Ne, or Ar.
• He is the most commercially important noble gas. It is used
as a coolant for performing experiments at low temps.
• Ar is used in light bulbs to conduct heat away from the
filament.
• Ne makes the red light in “neon” signs.
The f- Block: Lanthanides &
Actinides
• The elements of the 4f series are generally
called the lanthanides, after the element
lanthanum (La), which is the first member of
the series. Similarly, the elements of the 5f
series are called actinides, after the first
member of that series, actinium (Ac).
THE LANTHANIDES
• *Properties* they are very similar to each other and
easily lose 3 electrons forming 3+ ions. All are
silvery and quite soft but heavier and less reactive
than the alkali metals. They tarnish easily but react
slowly in water. They are found together in nature
and are difficult to separate.
• *USES* some alloys are used to make special
steel. Some compounds are found in the color in
your television.
THE ACTINIDES
• These elements are all radioactiveEVERY LAST ONE OF THE ISOTOPES
OF EVERY ELEMENT! Only Th and U
exist in nature… all others are primarily
man made.
5.3 Electron Configuration
and Periodic Properties
• As you move across the periodic table
you begin to see trends/predictable
patterns in the chemical and physical
properties of the elements. This
phenomenon is known as PERIODIC
TRENDS.
Atomic Radius
• Atoms have a radius called the ATOMIC RADIUS- the
distance from the center of the atom to the farthest most/ highest
energy level. This distance is not exact. We estimate.
• Contrary to popular opinion… atoms with more electrons
do not always have a larger atomic radius. The trend is:
– As you move from TOP to BOTTOM down the PTE the radius
grows larger
– As you move from RIGHT to LEFT across the PTE the radius
grows larger
ATOMIC RADII
Largest
• The first trend has a simple explanation: as you
move down through a group/family you are actually
adding additional principle quantum numbers (ex
group 1: 1s, 2s, 3s, 4s, 5s, 6s, and 7s) with corresponding
increases in the number of electrons from 1, 3, 11,
19, 37, 55, and 87. Ergo the radius must increase to
accommodate these increases in electrons.
• The second trend is more complex. As you move
across a period you are increasing both the # of
electrons as well as the number of protons. This
allows the protons to pull more strongly on the
electrons in the same energy level thus shrinking
the radius.
Ionic Size
• We know atoms can gain or lose electron to form Ions.
Depending whether they gain (becoming negatively charged) or lose
(becoming positively charged) will depend on whether ionic size will
shrink or grow.
• When you gain electrons- the ionic size grows
• As you lose electrons- the ionic size shrinks (more protons pulling
on fewer electrons= shrinking).
– The trend is complex- Figure 5-19 p.149
Ionization Energy
• The energy needed to strip an atom of an electron is called
IONIZATION ENERGY. Typically measured in J/atom you
can think of it as a measure of how strongly the atom holds
on to its valence electrons.
• Its trend is exactly opposite the atomic radius trend:
– As you move from BOTTOM to TOP the ionization energy
increases
– As you move from LEFT to RIGHT the ionization energy
increases
Ionization Energy
Hardest
Easiest
Successive Ionization
Energies
What if you want to pull more than one electron from an
atom? What happens to the ionization energy for the 2nd,
3rd, or even 4th electron?
• Successive ionization energies refer to the energy required
to remove the 2nd, 3rd and more electrons. The energy
required increases, partially because of electron-electron
repulsion, however it is not a linear increase.
– *note* there is always a dramatic increase in ionization energy
when you start extracting electrons from the inner noble gas core.
Electron Affinity
• If you can pull an electron from an atom… you
can ADD one to an atom. When an atom
gains an electron energy is REQUIRED! They
are reported in J/mol (NOT ATOMS). Not all
energy used in the addition of an electron is
gained by the atom that the electron is being
added to. It is possible to have + and –
electron affinities. If the value is + then the
atom gaining the electron is also gaining
energy. If the value is – then the atom gaining
the electron is releasing energy.
• General rule for the electron affinity trend– Nonmetals have more NEGATIVE electron
affinities than the metals.
– The exception is the noble gases.
• ** OCTET RULE** atoms tend to gain, lose or even share
electrons in order to have a full set of valence electrons (8)- to
have a full s and p orbital.
More
Negative
Electronegativity
• This is the ability of an atom to attract an electron in
a chemical bond. This is a unit less value.
• The least electronegative elements are in the
LOWER LEFT CORNER of the PTE. The trend is
the same as Ionization energy.
• When bonds occur… the atom with the higher
electronegativity would attract the electron more
strongly. Depending on the discrepancy of the
attraction there will form different types of bonds.
More electronegative