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Message-ID: <>
Newsgroups: alt.drugs
From: (Dalamar)
Date: Fri,  8 Jul 1994 06:59:04 UTC
Subject: CHEMISTRY: Atomic Structure

The Structure of the Atom

To obtain a model for the atom we must first examine the three basic types of
'building-blocks' from which atoms are constructed. These 'building-blocks' are
known as the proton, neutron and the electron. You will sometimes see these
referred to as 'subatomic particles'. Each of these particles has different
properties and plays a different role in an atom. Protons are positively
charged, each carrying a charge of +1. Neutrons, as the name might suggest, are
electrically neutral particles of about the same mass as a proton. Electrons are
negatively charged, each carries a charge of -1, exactly opposite and equal to
that on a proton. However, electrons are tiny when compared to the proton or
neutron - electrons have around 1/1836 the mass of a proton. This information is
presented in the table below. Mass is measured in atomic mass units, where 1
amu is equivalent to the mass of a proton or neutron.

Particle      Charge        Mass           Symbol

Proton         +1           1 amu            p

Neutron         0           1 amu            n                   

Electron       -1          1/1836 amu        e

It has been determined that in an atom the protons and neutrons bind together
to form a nucleus around which the electrons orbit. It is easy to see why this
model of the atom has been likened to a minature solar system. The nucleus of
the atom is the 'sun' and the electrons are the small orbiting 'planets'.
The number of protons in the nucleus of an atom is known as the _atomic number_.
The atomic number of an atom tells us which element it is from. For example an
atomic number of 3 tells us we are looking at a lithium atom and an atomic
number of 9 tells us we are looking at a fluorine atom. Atoms, when taken as a
whole, are electrically neutral. This means that the number of protons in the
nucleus must be matched by an equal number of orbiting electrons. Any excess or
deficiency in the number of electrons orbiting the nucleus, compared to the
number of protons in the nucleus, gives an overall charge imbalance. This
imbalance will be -1 extra for each surplus electron supplied above the number
of protons. If we add two electrons to a neutral atom it will acquire a net
charge of -2. If electrons are stripped away from a neutral atom we are left
with an excess in the number of protons over the number of electrons. As each
proton carries a +1 charge, each electron deficiency gives a +1 extra charge on
the atom. If we take three electrons away from a neutral atom it
acquires a net charge of +3. These charged atoms are known as _ions_.
Positive ions are known as _cations_ and negative ions are known as _anions_.
Before moving on a few examples will help to illustrate these ideas.

If possible find a copy of the periodic table of the elements. The elements in
the table are listed in order of increasing atomic number from left to right.
The horizontal rows are also known as _periods_. Each element in a period has
one more proton in its nucleus than the element to its immediate left. When
the far right end of a period is reached the addition of the next proton moves
us back to the left and one row down. The vertical columns of the table are
known as _groups_. Elements which make up groups are found to have very
similair properties to each other and this is not just mere coincidence, it
has its reasons rooted in something we shall go on to consider - the way in
which an atoms electrons are positioned around its nucleus.

Find fluorine in the periodic table, symbol F. In the box which details this
element will be its symbol, atomic number and _mass number_. The _mass number_
is the total number of protons plus neutrons in the nucleus, or the total number
of _nucleons_, a term which collectively refers to both protons and neutrons.
Sometimes these two numbers appear as superscript and subscript to the left of
the elements symbol. The superscript is the _mass number_, the total number
of protons plus neutrons. The subscript is the _atomic number_, the total number
of protons alone. For fluorine these values are 9 and 19. Now we have all the
information we need to formulate a picture of a fluorine atom. In the nucleus,
as indicated by the atomic number, are 9 protons. The mass number 19 tells us
that the total number of nucleons is 19, so the number of neutrons must be
(19 - 9) = 10 neutrons. Atoms are electrically neutral, therefore to balance
the +9 charge which the 9 protons introduce, there must be 9 orbiting electrons
giving a cancelling charge of -9. The electrons are held in orbit by the
electrostatic attractive force they feel from the positively charged nucleus.
Remember that charges of the opposite sign _attract_ one another, whilst
charges of the _same_ sign repel. If we now add an electron to the fluorine
atom the total number of electrons becomes 10, one more than the number of
protons in the F nucleus. This extra electron brings with it a -1 charge which
has no cancelling +1 proton in the nucleus. The fluorine 'atom' now carries a
net negative charge of -1. We no longer have a fluorine 'atom', but a
_fluoride ion_, in this case an _anion_ because it is negatively charged.
A diagram will illustrate these points further.

            Mass number    =  19  FFFFFFFF               
            Atomic number  =   9  F

                x x                      x x             In these diagrams the F
            x                        x         x         represents the nucleus
              x  F  x x                x  F  x           with its 9 protons and
            x                        x         x         10 neutrons. Each x 
                x x                      x x             represents an orbiting
          A fluorine atom           A fluoride ion

        Electrically neutral       Net charge of -1


The protons and neutrons (nucleons) of an atom are held tightly bound together
by a force known as the _strong nuclear force_. This force is extremely strong
and is required to overcome the repulsive forces that the protons exert on one
another due to their close proximity. Remember that the closer you try and bring
charges of opposite sign together, the greater is the replusive force they exert
on each other - much like trying to put the north pole ends of two magnets
together. To alleviate some of this repulsion is the function of the neutrons.
The neutrons act by 'diluting' the concentration of positive charge in the
nucleus by forcing the protons to be on average further apart. As the atomic
number rises, so do the repulsive forces present in the nucleus, with the result
that more neutrons are needed to 'dilute' the charge concentration.
Some elements display varying numbers of neutrons in the nuclei of their atoms.
For example, an atom of hydrogen has one proton in its nucleus and no neutrons.
But what if we introduce a neutron to the nucleus ? Remember, it is the number
of protons which determines which element we have, not the number of neutrons.
So what is this new atom we have created which has one proton, one neutron and
one orbiting electron ? The new atom is known as an _isotope_ of hydrogen.
Isotopes are elements with identical numbers of protons but differing numbers
of neutrons in their nuclei. In the case of hydrogen the isotope with the
1 neutron is known as _deuterium_. There also exists a hydrogen atom with
1 proton and 2 neutrons, known as _tritium_. However, in the case of hydrogen,
the fraction of deuterium atoms in any given sample is miniscule compared with
the number of 'normal' hydrogen atoms. We say that the natural abundance of
deuterium is small compared with the natural abundance of hydrogen.
If you look at the mass numbers for the elements you will see that alot of them
are _not_ whole number values. This is due to the presence of isotopes. The
number indicated as the mass number is an average of the isotopic masses
weighted for natural abundance. For example, chlorine exists as a mixture of
Cl-35 and Cl-37. When these mass numbers are averaged, taking into account the
percentage of each isotope present in a sample, the mass number comes out as
35.45. Because they are the same element, isotopes are identical in terms of
chemical reactivity, hence we never notice that chlorine is a mixture of 2

Electron Energy Levels  
So far you have seen that the atom consists of the proton, the neutron and
the electron. The protons and neutrons together form the nucleus of the atom,
around which orbit the electrons. The number of electrons must exactly match
the number of protons in order for overall electrical neutrality to be achieved.
The function of the neutrons is to stabilise the nucleus by diluting the
repulsive forces of the protons and that elements whose atoms can have differing
numbers of neutrons are known as isotopes.
When we come to examine the arrangement of the electrons around the nucleus a
distinct pattern emerges. It is found that the electrons occupy 'shells' which
are of well defined energy and distance from the nucleus. Electrons occupying
different shells are of different energies and distances from the nucleus.
The number of electrons a shell can hold is fixed and this number cannot be
exceeded. The first shell filled is the K shell, which can hold a maximum of
two electrons. The K shell is also the closest to the nucleus, which means
that electrons in it will be the most tightly held. When the K shell has been
filled by 2 electrons the next shell to fill is the L shell. The L shell is
capable of holding _eight_ electrons before it becomes full. The electrons in
the L shell are further away from the nucleus than those in the K shell, so are
not held so tightly by the attractive force from the nucleus. To build up a
picture of the occupancy of these shells in an atom whose atomic number we
know we use the following rules.

1. The shells are filled in order from lowest energy (closest to nucleus) to
   higher energy (further from nucleus).

2. The current shell _must_ be completely filled before moving on to fill the 
   next one of higher energy.

When this is done the atom is said to be in its _ground state_, the atom is
at a minimum of energy, all electrons occupy the lowest energy levels available.
The number of electrons in each shell can be indicated by listing the shells in
order of increasing energy, together with the number of electrons in that shell.
Hydrogen has one proton in its nucleus, so it must also have only one electron.
This single electron must occupy the shell of lowest energy - the one nearest
the nucleus - and this is the K shell. This may be written as K1, indicating
the lone occupancy of the K shell. The next element, helium, has an atomic
number of 2 indicating 2 protons in its nucleus. This is matched by 2 orbiting
electrons. Following our rules we must place _both_ of these electrons in the
K shell, which is then full. The electronic configuration of helium is therefore
K2. With the third element, lithium, we begin the filling of the L shell which
is capable of holding 8 electrons. The start of the new shell can be noticed in
the periodic table, where we jump from helium on the far right, to lithium on
the far left. If you count all the elements in the Li row, including Li, you
will see that there are 8, the same as the number of electrons the L shell may
hold before becoming full. The electronic configuration of Li, atomic number
three, is therefore K2 L1. The L shell will continue to fill as we traverse the
row, until we reach the element with the configuration K2 L8 (neon). Neon, like
helium, has a _full_ outer shell of electrons. It is the electrons in the
outermost shell of an atom which is responsible for the elements chemical
reactivity. The next shell to fill is the M shell which is capable of holding
18 electrons before becoming full. The element after neon, sodium, with atomic
number 11, therefore has the electronic configuration K2 L8 M1. Sodium, like
lithium, has only one electron in its outermost shell. Also, both sodium and
lithium are, like the rest of the group, soft metals with similar reactivity.
If you were to sit down and work out the electronic configurations of all the
group I metals (Li, Na, K etc) you would see that they all have one electron
in the outermost shell of their neutral atoms. It is this similarity in
electronic structure which causes the similarity in properties in the group I
metals and for other groups in the periodic table as well.
If you were to work out the electronic configurations for the atoms of the noble
gases (He, Ne, Ar etc), you would see that they all have their outermost shells
completely full. The noble gases are also extremely unreactive. This can be
attributed to the full outer shell of electrons, which provides stability and
unreactivity. This idea of a full outer shell of electrons providing stability
can be used as a powerful rationalising tool when discussing bonding between
atoms, where an atom will strive to acquire a full outer shell, either by the
gaining of electrons, loss of electrons or the sharing of electrons. I shall
cover bonding theory in another file, but first we need to look at a few more
of the properties of atoms which will aid us in predicting reactivity.

             *                                 *   *
                                          *             *
      *    * C *    *                          * Ne *    
                                          *             *

             *                                 *   *

       Carbon K2 L4                          Neon K2 L8

Ionisation Energy

The first ionisation energy of an atom is the amount of energy required to
remove one electron from the outermost shell to an infinite distance.
This may be represented by the equation :

                  E  ========>  E(+)    +    e(-)

Note that the total charge on either side of any equation is always equal, in
this particular case both sides are neutral (the positive charge on the cation
balances the negative charge of the electron).

The second ionisation energy is the energy required to remove a second electron
from the now unipositive ion. This process may be represented by the equation :

                  E(+)  ========>  E(2+)    +    e(-)   

Again, the charges on each side of the equation balance, in this case there is
a plus one charge on each side (the -1 charge on the single electron cancels one
of the two positive charges on E(2+) leaving a net +1.
Removing electrons from an atom requires us to do work, that is we must supply
sufficient energy in order to overcome the attractive force between nucleus and
electron. As we have already seen, the electrons occupy shells which are of
varying distance from the nucleus. Consequently electrons in different shells
experience different attractive forces from the nucleus and they will therefore
differ in the amount of energy needed to remove them. Remember, the closer the
electrons are to the nucleus, the harder it will be to remove them.
If we examine the first ionisation energy as a function of atomic number a
regular pattern emerges.

1. Across a period there is a steady _increase_ in first ionisation energy,
   which peaks at each noble gas. 

2. Down a group the first ionisation energy markedly _decreases_ from element to

The increase in I.E. across a period is due to the increasing nuclear charge
exerting a greater force on the orbiting electrons. Across a period the
electrons are being fed into the same shell, so they are all no further away
from the nucleus. However, the nuclear charge is _increasing_ and this naturally
has the effect of binding those electrons more tightly. This then leads to the
increase in I.E. which is observed in crossing a period.
Based on the argument of increasing nuclear charge you may have expected the
I.E. to increase down a group too, as each group member has more protons in its
nucleus than the one above it. This, you would reason, would cause an increase
in the attractive forces those outer electrons are going to feel and hence a
rise in I.E. However, we are forgetting that for each successive group member
the outermost electrons are in shells which are progressively further from the
nucleus. This increase in electron to nucleus distance produces a drop in the
attractive force which outweighs the increase in atomic number. The result is
a decrease in I.E. on descending any group.

From the above discussion it should now be clear that the elements with the
highest ionisation energies are those to the top and right of the periodic
table (eg O, F, Ne, Cl). These elements have ionisation energies in excess of
15 eV. The elements with the lowest I.E.s are those to the left and bottom of
the periodic table (eg Cs, Fr,). These elements have I.E.s around or below
below 5 eV. Knowing the exact figures isn't important as long as you have an
idea of the trends. Knowledge of an elements I.E. can allow us to predict, for
example, whether that element will be an oxidising or reducing agent. As an
example of how I.E.s differ down a group here are the first and second I.E.s of
the group I metals.

Metal       First I.E.       Second I.E.  

Li            520              7296               The measurements here are in
                                                  kilo-joules per mole. The mole
Na            496              4563               is a unit of measurement of
K             419              3069               
Rb            403              2650         

Cs            375              2420                            

For sodium the first I.E. is 496 kJ/mol, this represents the amount of energy
required to remove the single M electron to leave the Na+ cation (K2 L8).
The amount of energy required to remove the second electron is huge compared
with the first - 4563 kJ/mol. There are two reasons for this.

1. The second electron is being removed from a _full_ orbital shell which
   contains electrons closer to the nucleus than the original single M
   electron already removed ie the process is K2 L8 ====> K2 L7 in which
   we are breaking into a _full shell_ in which the electrons are closer
   to the nucleus.

2. The second electron is being removed from an already positively charged
   cation, with the result that we need to do more work in order to overcome
   this extra attractive force.

   Na ========> Na(+)    +      e(-)      requires _less_ energy than :

   Na(+) ========> Na(2+)    +     e(-)

Size of Atoms and Ions

Across a period in the periodic table, electrons are being fed into the
same shell, so you may have expected no change in atomic size as we cross
the period. However, in traversing the period we introduce more and more
positive nuclear charge, with the result that the electrons being fed into the
current shell feel the pull of the nucleus more strongly, thus there is a
contraction in atomic size. Down groups there is an _increase_ in atomic size
as, going from one element to the next in the group, the outermost electrons
are in shells progressively further from the nucleus.

Anions (negative ions) are always larger than their parent atoms. The reason
being that the addition of an electron to the atom will cause an increase in the
replusive field that the orbiting electrons mutually feel. This increase causes
the electrons to spread out more in space thus increasing the size of the ion in
comparison to the size of the atom.

Cations (positive ions) are always smaller than their parent atoms. A loss of
one or more electrons causes a reduction in the repulsive forces between the
electrons and thus an overall contraction in radius. Also, the electrons which
are lost may totally empty the outer shell, which will naturally lead to a
reduction in radius as the next inner shell is closer to the nucleus. A sodium
atom, for example, has the electronic configuration K2 L8 M1. Loss of a single
electron gives a sodium ion, Na+, which has the stable noble gas electronic
configuration of neon, K2 L8. The loss of the single electron from the M shell
gives a natural reduction to the radius of the cation vs atom, as the outermost
electrons are now in the L shell and not the M shell. In addition, the ratio of
positive charges on the nucleus to the number of orbital electrons is increased.
Thus the effective nuclear charge is increased and the electrons are pulled in.
The greater the charge on the cation, the smaller it becomes. 

For Sodium:

Atomic Radius Na (K2 L8 M1) = 1.57 Angstroms             1 angstrom =
Ionic Radius Na+ (K2 L8) = 0.98 Angstroms               0.0000000001 metres


The electronegativity of an atom is a measure of its ability to attract
electrons to itself when the atom is bonded to others as part of a compound.
An atoms ability to attract electrons to itself depends greatly upon its size.
Generally, the smaller the atom, the greater is its electronegativity ie the
better is its ability to attract electrons. We have already seen that across
a period there is a decrease in atomic size which corresponds to increasing
nuclear charge. Down groups in the table there is a marked increase in size
as the outermost electrons are in orbital shells progressively further from the
nucleus. These trends indicate that across periods there is an _increase_ in
electronegativity and down groups there is a _decrease_ in electronegativity.
Therefore the most electronegative elements are to be found at the top right
of the periodic table (N, O, F, Cl) and the least electronegative are at the
bottom left (Rb, Cs).
The electronegativities of the elements can be placed on a scale of 0-4, with
fluorine, the most electronegative element, assigned the value of 4. The
following partial periodic table lists some electronegativity values.

H  (2.1)
Li (1.0)    Be (1.5)   B (2.0)    C (2.5)    N (3.0)   O (3.5)    F (4.0)
Na (0.9)                                                          Cl (3.0)
K  (0.8)                                                          Br (2.8)
Rb (0.8)                                                          I (2.5)
Cs (0.7)

This particular scale is known as the Pauling scale after its inventor.

Atoms whose electronegativity falls below the 2.1 mark compete poorly for
electrons, in fact these elements are sometimes referred to as electropositive
because they have very little pulling power. They also happen to be the elements
with low ionisation energy. The lower the value of EN below 2.1 the more
electropositive the element will be, so that Cs, with an EN value of around
0.7 is very electropositive indeed (has very low ionisation energy and competes
poorly for electrons when it is part of a compound).

Electronegativity is a useful concept for chemists. For example, the difference
in electronegativity between two bonding atoms can be used to predict whether
that bond will be predominantly _ionic_ or _covalent_. If you do not understand
what is meant by these two terms then don't worry - I shall cover them in the
next file : Bonding and Structure.


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