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INTRODUCTION
The Periodic Table of Elements has developed into a very important tool for the chemist.
It can supply much needed information about the atoms, such as; atomic mass, atomic number
and electron configuration. It is set up in such a way as to also allow us an idea of activity. The
arrangement of the elements on the Periodic Table allow us to better understand how a substance
will react with another substance.
HISTORY
The periodic table is a tool that has not change much. At one time it was believed that
the way to describe the makeup of all things was to use the relationship of fire, air, water and
earth. Not much was known about the elements at this time. As more and more elements were
discovered there had to be a way of arranging them into a usable pattern.
Johann Dobereiner did some work in the late 1820's. He noted a relationship between the
properties of certain elements and their atomic masses. Dobereiner knew that chlorine, iodine
and bromine all had similar properties. He arranged these elements according to their atomic
masses. The lightest (chlorine) was placed first and the heaviest iodine was place last. Bromine
was then placed in the middle. Dobereiner noticed that the difference in the mass between
chlorine and bromine was about the same as the difference in mass between bromine and iodine.
This same relationship was observed in other groups as well (sulfur, selenium and tellurium;
calcium, strontium, and barium; lithium, sodium and potassium). These groups of three similar
elements are known as Dobereiner's triads.
In 1864 John Newlands saw a connection between the properties of elements and their
atomic masses. He indicated that if the elements were arranged according to increasing mass
that every eight element would have similar properties. (1-8, 2-9, 3-10...). It must be understood
that at this time the inert gasses were not yet discovered.
In 1869 Dimitri Mendeleev developed a table that classified the elements according to
their atomic mass. This table is the basis for the one we use today. He placed elements with
similar properties in the same row. He left blank space for elements that were not yet
discovered. Mendeleev was the first to publish a classification of elements.
In 1914 Henry Mosley suggested that the elements on Mendeleev's table be arranged by
increasing atomic number rather than their mass. This made some minor changes in Mendeleev's
table, which placed the elements with similar properties in the same group (Look at Te and I).
This table became the table that we are now using. We can now state the PERIODIC LAW:
The chemical and physical properties of the elements are periodic functions of
their atomic number.
STRUCTURE
The Periodic Table is made up of three types of elements; metals, nonmetals, and
metalloids (semimetals). Metals make up the largest number of elements and are found on the
left side of the table. These metals will form positive ions. They form positive ions because as
atoms they have one, two or three electrons in their outer shell. When they form ions they will
lose these electrons to gain electron configurations that make them stable. The number of
protons do not change therefore when a metallic atom loses an electron (negatively charge
particle) it becomes positively charged because it now has less negative charges than it originally
had. When they lose these electrons they will end up with a smaller radius than when they were
atoms.
Group I or all those elements in the first column, except hydrogen, will all have similar
properties. This group is called the Alkali metals. They will all form ions with a +1 charge.
They all have one s electron in their outer shell. These are very active metals and are seldom
found free in nature. As a matter of fact these elements are not found free in nature. In order to
keep some pure material, such as sodium, it must be stored under kerosene because it would
react with the moisture in the air.
The second column is made up of metals called the alkaline earth metals. These are
also very active elements, but not as active as those in group 1. These group 2 elements have
properties similar to each other. They all have two s electrons in their outer shell and will
therefore produce a +2 ion.
Another group of metals is the transition elements. These metals are located in the
center of the periodic table. They make up groups 3 through 12. When they form ions it
becomes more complicated because they will have multiple energy levels involved. They
therefore produce variable charges (oxidation states) when ions are formed. This area represents
the filling of the d sublevel. Not all will have variable oxidation states. One important
characteristic is that we can have colored solutions formed when transition elements are present.
All transition elements don't form colored solutions, but if you have a colored solution you must
have a transition element present.
The rare earth elements are the two rows at the bottom of the periodic table. These
elements are rare. Most are unstable and will have only the most stable isotope listed. Because
they are unstable many of these are associated with radioactivity. In this area represents the
filling of the f sublevel.
The nonmetals are found on the right side of the periodic table. They are only a small
number of elements - mostly gases. These elements form ions that are larger than their atoms
because they have more electrons in their outer shell so it is easier for them to gain electrons to
complete their outer shell than it would be for them to loses electrons. When they gain electrons
they will have more negative particles around the nucleus therefore the will form ions that have a
negative charge. Remember the number of protons can't change. The two most important
groups of nonmetals are the halogens and the inert gases.
The halogens are those elements that are located in group 17. It is the only group on the
periodic table that has a solid, liquid, and gas. Each member of this group forms an ion, when it
is the nonmetal by itself, of -1. This is because they each have 7 electrons in their outer shell and
each would like to have eight electrons in their outer shell. Fluorine is a member of this group
and is considered to be the most active nonmetal. Fluorine and chlorine are gases at room
temperature, while iodine is a liquid and bromine and astatine are solids.
The other important group of nonmetals is the inert gases. Sometimes these are called
the Noble gases or the inert gas structures. They have completed outer shells and are therefore
stable elements. This is called a stable octet (except He which only has 2 electrons in it's outer
shell). The neither want to gain or lose electrons. They appear to have the right number of
electrons in their outer shell. These elements were some of the last elements discovered because
they were unreactive. These elements will not react with other elements and this is why they are
called inert gases (unreactive). We find now that some of the larger molecular structures can be
made to react with Fluorine under certain conditions.
Metalloids are the other type of elements on the periodic table. They are found along the
line drawn like steps between boron and aluminum, aluminum and silicon and so on. This
appears as a dark line on your periodic table. Aluminum is the exception here. Those elements
found along this line are metalloids, except aluminum. Aluminum is a metal. They separate the
metals from the nonmetals. They have characteristics that resemble both metals and nonmetals.
TRENDS
There are trends that we can follow on the periodic table. We will look at the group or
family (vertical columns) and the period or series (horizontal rows).
As you go from top to bottom (vertical columns) in a group or family you are adding an
additional energy level for each successive atom. Because chemical activity will occur in the
outer shell, it is important to understand that the outer shell gets farther and farther from the
nucleus. The farther an electron gets from the nucleus the less attractive force on those electrons.
We can look at some properties and use this relationship to explain what happens. Because the
elements are arranged by increasing atomic number we will look at these trends in the same was
- increasing atomic number:
Groups or Families
1.
Metallic properties increase. Why? Metals want to lose electrons the farther the
electron is from the nucleus the easier it will be to lose the electron.
2.
Radius increases. There is an additional energy level added to
each successive atom.
3.
Ionization energy decreases (Table K). It takes less energy to form
an ion.
4.
Electronegativity decreases (Table K). Electronegativity is the attraction an atom has for
an electron. If the outer shell (valence
shell) is farther from the nucleus there will not be as
great attraction
if it were closer.
5.
Oxidation states (charges) remains the same because they all have the same number of
electrons in the outer shell and will therefore behave the same.
6. They all have similar properties.
Periods or Series
1. Metallic properties decrease. As you go from one element to the
next the number of
electrons in the outer shell (valence shell) will
increase. A metal wants to lose electrons
and a nonmetal wants to gain electrons. The more electrons present in this outer shell the harder
it is to lose them all.
2.
Radius decreases. All the elements that are on the same period
will have the same
number of energy levels. The number of protons in the nucleus increases making the nuclear
charge stonger.
It therefore pulls the electrons toward it making the radius smaller as you
go from one element to the next.
3.
Ionization energy increases. There are more electrons in the outer shell and it is
therefore harder to remove them. The electrons are also closer to the nucleus so there is a greater
force holding them there.
4.
Electronegativity increases. These elements want to gain electrons to complete their
outer shell, so they have a greater force of attraction.
5. Oxidation states change (+1, +2, +3, + - 4, -3 , -2, -1)). They all have a different number of
electrons in the outer shell.
6.
Properties are all different
IONIZATION ENERGY
Ionization energy is the energy required to remove an electron. First ionization energy,
as indicated on Table K, is the energy needed to remove the most loosely bound electron. Metals
have low ionization energies because they want to lose electrons to form ions. The electrons that
are farthest from the nucleus will have the least attraction by the nucleus and will therefore take
the least amount of energy to be removed from the atom to form an ion. In the example below
the first ionization energy for sodium (Na) will be higher than potassium (K) because the
electron that it is removing is closer to the nucleus. The closer an electron is to the nucleus the
stronger the force of attraction and therefore the more energy needed to overcome that force.
Na + energy (1st ionization energy) 
Na+ + e-
K + energy (1st ionization energy)  K+ + eBoth potassium (K) and sodium (Na) are in the same family. They both have one
electron in their outer shell. To complete the outer shell they must either lose 1 electron or gain
7 electrons. If you associate a certain amount of energy with each electron it would be easier to
lose just one electron that to try and gain 7. Potassium needs less energy to remove its electron
because the outer shell is on the fourth energy level (period 4). It takes 100 kcal/mole of atoms.
Sodium only needs to lose one electron also, but its outer shell is closer to the nucleus so it take
more energy to remove it. It takes 119 kcal/mole of atoms.
ELECTRONEGATIVITY
Electronegativity is the attraction an atom has for an electron. It is an arbitrary scale
with Fluorine having the highest value of 4. Nonmetal have higher electronegativities than
metals because they want to attract electrons to complete their outer shells. The farther the
electron is from the nucleus the weaker the attractive force. That is why bromine has a lower
electronegativity than fluorine because the attraction that bromine has for an electron would
attract it on the fourth energy level as apposed to the second energy level like fluorine. Table K
will give us the electronegativity of many elements. These values will be very important to us in
the next chapter on bonding.
We should make sure that we understand the relationship of electron configurations on
the periodic table. As we move from one element to the next on the periodic table (increasing
atomic number) we add one additional electron. This was covered in the last chapter on Atomic
Structure. There is a certain order in which the energy levels outside the nucleus are arranged.