Transcript History of Periodic Table

The periodic table is structured so that elements with the same
type of valence electron configuration are arranged in columns.
Flame Test
Flame Test
• a QUALITATIVE test for the presence of metals
in chemical compounds.
• When the compound is excited by heating it in a
flame, the metal ions will begin to emit
light. (WHY??!!)
• Based on the emission spectrum of the element,
the compound will turn the flame a characteristic
color.
• This technique of using certain chemical
compounds to color flames is widely used in
pyrotechnics to produce the range of colors
seen in a firework display.
Viewing spectrum through cobalt
glass
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Certain metal ions will turn the flame very distinctive colors, these colors intern
can help identify the presence of a particular metal in a compound. However,
some colors are produced by several different metals, making it hard to
determine the exact ion or concentration of the ion in the compound. Some
colors are very weak and are easily overpowered by stronger colors. For
instance, the presence of a Potassium Ion in a compound will color a flame
violet / lilac, on the other hand, even trace amounts of Sodium ions in a
compound produce a very strong yellow flame, often times making the
Potassium ion very difficult to detect. To counteract the effects of any Sodium
impurities, one can view the flame through a piece of Cobalt blue glass . The
Cobalt glass absorbs the yellow light given off by Sodium while letting most
other wavelengths of light pass through. More recently, didymium glass has
been substituted for Cobalt glass due to its superior ability to block undesirable
light.
Flame Colorants
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Lithium Chloride
Strontium Chloride
Calcium Chloride (a bleaching powder)
Sodium Chloride (table salt)
or Sodium Carbonate
Yellowish Green Borax
Green Copper Sulfate or Boric Acid
Blue Copper Chloride
Violet 3 parts Potassium Sulfate
1 part Potassium Nitrate (saltpeter)
Purple Potassium Chloride
White Magnesium Sulfate (Epsom salts)
• http://www.khanacademy.org/video/valence
-electrons
September
• Color the Periodic Table
th
27
Which electron
arrangements are stable
and which are not?
• Both atoms and ions considered
• When an atom has its outer level full to
the maximum number of electrons
allowed, the atom is particularly stable
electronically and very unreactive.
• This is the situation with the Noble
Gases: He is [2], neon is [2,8] and
argon is [2,8,8]
• There atoms are the most reluctant to
lose, share or gain electrons in any sort
of chemical interaction because they
are so electronically stable.
• For all elements most of their chemistry is
about what outer electrons do or don't!
• [2], [2,8] and [2,8,8] etc. are known as the
'stable Noble Gas arrangements', and the
atoms of other elements try to attain this
sort of electron structure when reacting to
become more stable.
• The most reactive metals have just
one outer electron.
• These are the Group 1 Alkali Metals,
lithium [2,1], sodium [2,8,1], potassium
[2,8,8,1]
• With one outer shell electron, they have one
more electron than a stable Noble Gas
electron structure.
• So, they readily lose the outer electron
when they chemically react to try to form (if
possible) one of the stable Noble Gas
electron arrangements - which is why atoms
react in the first place!
• When Group 1 Alkali Metal atoms lose an
electron they form a positive ion because
the positive proton number doesn't change,
but with one negative electron lost, there is a
surplus of one + charge e.g.
• sodium atom ==> sodium ion
• Na ==> Na+
• is [2.8.1] ==> [2.8] electronically
• in fundamental particles [11p + 11e] ==> [11p
+ 10e]
• IONS are atoms or group of atoms
which carry an overall electrical
charge i.e. not electrically neutral.
• The most reactive non-metals are just
one electron short of a full outer
shell. These are the Group 7
Halogens, namely fluorine [2,7],
chlorine [2,8,7] etc.
• These atoms are one electron short of a
stable full outer shell and seek an 8th
outer electron to become
electronically stable - yet again, this is
why atoms react!
• They readily gain an outer electron, when
they chemically react, to form one of the
stable Noble Gas electron arrangements
either by sharing electrons (in a covalent
bond) or by electron transfer forming a singly
charged negative ion (ionic bonding) e.g.
• chlorine atom ==> chloride ion
• Cl ==> Cl• is [2.8.7] ==> [2.8.8] electronically
• in fundamental particles [17p + 17e] ==> [17p
+ 18e]
• the positive proton number of Cl doesn't
change but the chloride ion carries one extra
negative electron to give the surplus charge
of a single - on the ion.
Nuclide notation and ions
• sodium-24 isotope
ion, 11 protons, 13
neutrons, 10
electrons (one
electron lost to form
a positive ion)
Nuclide notation and ions
• sodium-23 isotope
ion, 11, protons, 12
neutrons, 10
electrons (one
electron lost to form
a positive ion)
Nuclide notation and ions
• isotope sulfur-32 in the
form of the sulfide ion,
16 protons, 16
neutrons, 18 electrons
(two electrons gained
to form the double
charged negative ion)
• We will discuss more on electron structure
and chemical changes and compound
formation LATER
• The 'chemical structure' of the periodic
Table, that is the chemical similarity of the
vertical groups elements, was known well
before the electronic structure of atoms
was understood. However, it wasn't
understood why they behaved in the same
way chemically nor was it understood at first
why Noble Gases were so unreactive towards
other elements. BUT once the electronic
structure of atoms was understood,
'electronic' theories could then be applied to
explain the chemical similarity of elements in
a vertical Group of the Periodic Table.
• Originally they were laid out in order
of ' relative atomic mass' (the old term
was 'atomic weight'). This is not correct
for some elements now that we know
their detailed atomic structure in terms
of protons, neutrons and electrons, and
of course, their chemical and physical
properties.
• Hydrogen, 1, H, does not readily fit
into any group
Group
• is a vertical column of chemically and
physically similar elements e.g. Group 1
The Alkali Metals (Li, Na, K etc.), Group 7
The Halogens (F, Cl, Br, I etc.) and Group 0
The Noble Gases (He, Ne, Ar etc.). The group
number equals the number of electrons in the
outer shell (e.g. chlorine's electron
arrangement is 2.8.7, the second element
down Group 7 on period 3).
Period
• is a horizontal row of elements with a
variety of properties (left to right goes
from metallic to non-metallic elements.
All the elements use the same number
of electron shells which equals the
period number (e.g. sodium's electron
arrangement 2.8.1, the first element in
Period 3).
• The periodicity of elements i.e. the
repetition of very chemically similar
elements in a group is due to the
repetition of a the same outer electron
structure - check out the last number
from element 3 onwards.
• Knowledge of the electron configuration of
different atoms is useful in understanding
the structure of the periodic table of
elements.
• The electron configuration is based on the
Quantum Model of the atom where
electrons exist in orbitals, energy levels,
follow rules of Pauli, Hund, etc
• The Periodic Table was developed LONG
before this time however.
• How did they construct a periodic table that
graphically exhibits the electron configurations so
well…without knowing anything about the
quantum model of the atom?
• the Periodic Table is based on the
electronic structure of atoms BUT WAS
CREATED BEFORE KNOWING THE electronic
structure of atoms. How can this be true?
• Mendeleev (Russian chemist) first published
his 'Periodic Table' work simultaneously in
1869 with the work of Lothar Meyer (German
chemist) who looked at the physical
properties of all known elements. Lothar
Meyer noted 'periodic' trend patterns e.g.
peaks and troughs when melting or boiling
points, specific heat and atomic volume
values were plotted against 'atomic weight' what we now call relative atomic mass.
Mendeleev's early versions
of the Periodic Table